chpt1 11

MatterAnd

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© 2009, Prentice-Hall, Inc.

Chemistry

In this science we study matter and the changes it undergoes.

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Matter

We define matter as anything that has mass and takes up space.

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Matter

• Atoms are the building blocks of matter.

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Matter

• Atoms are the building blocks of matter.• Each element is made of the same kind of atom.

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Matter

• Atoms are the building blocks of matter.• Each element is made of the same kind of atom.• A compound is made of two or more different kinds of

elements.

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States of Matter

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Classification of Matter

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Properties and Changes of

Matter

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Types of Properties

• Physical Properties…– Can be observed without changing a

substance into another substance.• Boiling point, density, mass, volume, etc.

• Chemical Properties…– Can only be observed when a substance is

changed into another substance.• Flammability, corrosiveness, reactivity with

acid, etc.

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Types of Properties

• Intensive Properties…– Are independent of the amount of the

substance that is present.• Density, boiling point, color, etc.

• Extensive Properties…– Depend upon the amount of the substance

present.• Mass, volume, energy, etc.

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Types of Changes

• Physical Changes– These are changes in matter that do not

change the composition of a substance.• Changes of state, temperature, volume, etc.

• Chemical Changes– Chemical changes result in new substances.

• Combustion, oxidation, decomposition, etc.

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Chemical Reactions

In the course of a chemical reaction, the reacting substances are converted to new substances.

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Compounds

Compounds can be broken down into more elemental particles.

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Separation of Mixtures

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Filtration

In filtration solid substances are separated from liquids and solutions.

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Distillation

Distillation uses differences in the boiling points of substances to separate a homogeneous mixture into its components.

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Chromatography

This technique separates substances on the basis of differences in solubility in a solvent.

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What do these countries have in common?

US, Liberia and Burma


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What do these countries have in common?

US, Liberia and Burma

• They use the imperial system


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View of Countries using Metric


LiberiaBerma

USA

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Units of Measurement

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SI Units

• Système International d’Unités• A different base unit is used for each quantity.

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Metric System

Prefixes convert the base units into units that are appropriate for the item being measured.

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Volume

• The most commonly used metric units for volume are the liter (L) and the milliliter (mL).– A liter is a cube 1 dm

long on each side.– A milliliter is a cube 1 cm

long on each side.

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Uncertainty in Measurement

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Uncertainty in Measurements

Different measuring devices have different uses and different degrees of accuracy.

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1 ml

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0.1 ml

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Accuracy versus Precision

• Accuracy refers to the proximity of a measurement to the true value of a quantity.

• Precision refers to the proximity of several measurements to each other.

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Significant Figures

• The term significant figures refers to digits that were measured.

• When rounding calculated numbers, we pay attention to significant figures so we do not overstate the accuracy of our answers.

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Significant Figures

1. All nonzero digits are significant.

2. Zeroes between two significant figures are themselves significant.

3. Zeroes at the beginning of a number are never significant.

4. Zeroes at the end of a number are significant if a decimal point is written in the number.

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Significant Figures

• When addition or subtraction is performed, answers are rounded to the least significant decimal place.

• When multiplication or division is performed, answers are rounded to the number of digits that corresponds to the least number of significant figures in any of the numbers used in the calculation.

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Temperature

By definition temperature is a measure of the average kinetic energy of the particles in a sample.

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Temperature• In scientific

measurements, the Celsius and Kelvin scales are most often used.

• The Celsius scale is based on the properties of water.– 0C is the freezing point

of water.– 100C is the boiling point

of water.

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Temperature

• The Kelvin is the SI unit of temperature.

• It is based on the properties of gases.

• There are no negative Kelvin temperatures.

• K = C + 273.15

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Temperature

• The Fahrenheit scale is not used in scientific measurements.

F = 9/5(C) + 32 C = 5/9(F − 32)

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Density

Density is a physical property of a substance.

d =mV

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Dimensional Analysis

• We use dimensional analysis to convert one quantity to another.

• Most commonly dimensional analysis utilizes conversion factors (e.g., 1 in. = 2.54 cm)

1 in.

2.54 cm

2.54 cm

1 in.or

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Use the form of the conversion factor that puts the sought-for unit in the numerator.

Given unit desired unitdesired unit

given unit

Conversion factor

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• For example, to convert 8.00 m to inches,– convert m to cm– convert cm to in.

8.00 m100 cm

1 m

1 in.

2.54 cm 315 in.

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Atomic Theory of Matter

The theory that atoms are the fundamental building blocks of matter reemerged in the early 19th century, championed by John Dalton.

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Dalton's Postulates

Each element is composed of extremely small particles called atoms.

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Dalton's Postulates

All atoms of a given element are identical to one another in mass (?) and other properties, but the atoms of one element are different from the atoms of all other elements.

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Dalton's Postulates

Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.

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Dalton’s Postulates

Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms.

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Law of Constant CompositionJoseph Proust (1754–1826)

• This is also known as the law of definite proportions.

• It states that the elemental composition of a pure substance never varies.

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Law of Conservation of Mass

The total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place.

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The Electron

• Streams of negatively charged particles were found to emanate from cathode tubes.

• J. J. Thompson is credited with their discovery (1897).

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The Electron

Thompson measured the charge/mass ratio of the electron to be 1.76 108 coulombs/g.

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Millikan Oil Drop Experiment

Once the charge/mass ratio of the electron was known, determination of either the charge or the mass of an electron would yield the other.

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Millikan Oil Drop Experiment

Robert Millikan (University of Chicago) determined the charge on the electron in 1909.

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Radioactivity

• Radioactivity is the spontaneous emission of radiation by an atom.

• It was first observed by Henri Becquerel.

• Marie and Pierre Curie also studied it.

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Radioactivity• Three types of radiation were discovered by

Ernest Rutherford: particles particles rays

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The Atom, circa 1900

• The prevailing theory was that of the “plum pudding” model, put forward by Thompson.

• It featured a positive sphere of matter with negative electrons imbedded in it.

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Discovery of the Nucleus

Ernest Rutherford shot particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.

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The Nuclear Atom

Since some particles were deflected at large angles, Thompson’s model could not be correct.

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The Nuclear Atom• Rutherford postulated a very small,

dense nucleus with the electrons around the outside of the atom.

• Most of the volume of the atom is empty space.

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Other Subatomic Particles

• Protons were discovered by Rutherford in 1919.

• Neutrons were discovered by James Chadwick in 1932.

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Subatomic Particles

• Protons and electrons are the only particles that have a charge.

• Protons and neutrons have essentially the same mass.

• The mass of an electron is so small we ignore it.

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Symbols of Elements

Elements are symbolized by one or two letters.

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Atomic Number

All atoms of the same element have the same number of protons:

The atomic number (Z)

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Atomic Mass

The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom.

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Isotopes

• Isotopes are atoms of the same element with different masses.

• Isotopes have different numbers of neutrons.

116C

126C

136C

146C

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Atomic Mass

Atomic and molecular masses can be measured with great accuracy with a mass spectrometer.

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Average Mass

• Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations.

• Average mass is calculated from the isotopes of an element weighted by their relative abundances.

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Periodic Table

• It is a systematic catalog of the elements.

• Elements are arranged in order of atomic number.

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Periodicity

When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities.

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Periodic Table

• The rows on the periodic chart are periods.

• Columns are groups.• Elements in the same

group have similar chemical properties.

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Groups

These five groups are known by their names.

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Periodic Table

Nonmetals are on the right side of the periodic table (with the exception of H).

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Periodic Table

Metalloids border the stair-step line (with the exception of Al, Po, and At).

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Periodic Table

Metals are on the left side of the chart.

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Chemical FormulasThe subscript to the right of the symbol of an element tells the number of atoms of that element in one molecule of the compound.

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Chemical FormulasMolecular compounds are composed of molecules and almost always contain only nonmetals.

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Diatomic Molecules

These seven elements occur naturally as molecules containing two atoms.

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Types of Formulas

• Empirical formulas give the lowest whole-number ratio of atoms of each element in a compound.

• Molecular formulas give the exact number of atoms of each element in a compound.


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Types of Formulas

• Structural formulas show the order in which atoms are bonded.

• Perspective drawings also show the three-dimensional array of atoms in a compound.

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Ions

• When atoms lose or gain electrons, they become ions.– Cations are positive and are formed by elements

on the left side of the periodic chart.– Anions are negative and are formed by elements

on the right side of the periodic chart.

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Ionic Bonds

Ionic compounds (such as NaCl) are generally formed between metals and nonmetals.

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Writing Formulas

• Because compounds are electrically neutral, one can determine the formula of a compound this way:– The charge on the cation becomes the subscript

on the anion.– The charge on the anion becomes the subscript

on the cation.– If these subscripts are not in the lowest whole-

number ratio, divide them by the greatest common factor.

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Common Cations

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Common Anions

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Inorganic Nomenclature

• Write the name of the cation.

• If the anion is an element, change its ending to -ide; if the anion is a polyatomic ion, simply write the name of the polyatomic ion.

• If the cation can have more than one possible charge, write the charge as a Roman numeral in parentheses.

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Patterns in Oxyanion Nomenclature

• When there are two oxyanions involving the same element:– The one with fewer oxygens ends in -ite.

• NO2− : nitrite; SO3

2− : sulfite

– The one with more oxygens ends in -ate.• NO3

− : nitrate; SO42− : sulfate

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• The one with the second fewest oxygens ends in -ite.– ClO2

− : chlorite

• The one with the second most oxygens ends in -ate.– ClO3

− : chlorate

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• The one with the fewest oxygens has the prefix hypo- and ends in -ite.

– ClO− : hypochlorite

• The one with the most oxygens has the prefix per- and ends in -ate.

– ClO4− : perchlorate

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Acid Nomenclature

• If the anion in the acid ends in -ide, change the ending to -ic acid and add the prefix hydro- .– HCl: hydrochloric acid– HBr: hydrobromic acid– HI: hydroiodic acid

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Acid Nomenclature

• If the anion in the acid ends in -ite, change the ending to -ous acid.– HClO: hypochlorous

acid

– HClO2: chlorous acid

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Acid Nomenclature

• If the anion in the acid ends in -ate, change the ending to -ic acid.– HClO3: chloric acid

– HClO4: perchloric acid

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Nomenclature of Binary Compounds

• The less electronegative atom is usually listed first.

• A prefix is used to denote the number of atoms of each element in the compound (mono- is not used on the first element listed, however) .

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• The ending on the more electronegative element is changed to -ide.

– CO2: carbon dioxide– CCl4: carbon tetrachloride

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• If the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are often elided into one.

N2O5: dinitrogen pentoxide

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Nomenclature of Organic Compounds

• Organic chemistry is the study of carbon.• Organic chemistry has its own system of

nomenclature.

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The simplest hydrocarbons (compounds containing only carbon and hydrogen) are alkanes.

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The first part of the names above correspond to the number of carbons (meth- = 1, eth- = 2, prop- = 3, etc.).

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• When a hydrogen in an alkane is replaced with something else (a functional group, like -OH in the compounds above), the name is derived from the name of the alkane.

• The ending denotes the type of compound.– An alcohol ends in -ol.

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Law of Conservation of Mass“We may lay it down as an

incontestable axiom that, in all the operations of art and nature,

nothing is created; an equal amount of matter exists both

before and after the experiment. Upon this principle, the whole art

of performing chemical experiments depends.”

--Antoine Lavoisier, 1789

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Chemical Equations

Chemical equations are concise representations of chemical reactions.

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Anatomy of a Chemical Equation

CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)

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Reactants appear on the left side of the equation.

CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)

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Products appear on the right side of the equation.

CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)

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The states of the reactants and products are written in parentheses to the right of each compound.

CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)

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Coefficients are inserted to balance the equation.

CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)

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Subscripts and Coefficients Give Different Information

• Subscripts tell the number of atoms of each element in a molecule.

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Subscripts and Coefficients Give Different Information

• Subscripts tell the number of atoms of each element in a molecule

• Coefficients tell the number of molecules.

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Reaction Types

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Combination Reactions

• Examples:– 2 Mg (s) + O2 (g) 2 MgO (s)

– N2 (g) + 3 H2 (g) 2 NH3 (g)

– C3H6 (g) + Br2 (l) C3H6Br2 (l)

• In this type of reaction two or more substances react to form one product.

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• In a decomposition one substance breaks down into two or more substances.

Decomposition Reactions

• Examples:– CaCO3 (s) CaO (s) + CO2 (g)

– 2 KClO3 (s) 2 KCl (s) + O2 (g)

– 2 NaN3 (s) 2 Na (s) + 3 N2 (g)

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Combustion Reactions

• Examples:– CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)

– C3H8 (g) + 5 O2 (g) 3 CO2 (g) + 4 H2O (g)

• These are generally rapid reactions that produce a flame.

• Most often involve hydrocarbons reacting with oxygen in the air.

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Formula Weights

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Formula Weight (FW)• A formula weight is the sum of the

atomic weights for the atoms in a chemical formula.

• So, the formula weight of calcium chloride, CaCl2, would be

Ca: 1(40.1 amu)

+ Cl: 2(35.5 amu)

111.1 amu

• Formula weights are generally reported for ionic compounds.

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Molecular Weight (MW)

• A molecular weight is the sum of the atomic weights of the atoms in a molecule.

• For the molecule ethane, C2H6, the molecular weight would be

C: 2(12.0 amu)

30.0 amu+ H: 6(1.0 amu)

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Percent Composition

One can find the percentage of the mass of a compound that comes from each of the elements in the compound by using this equation:

% element =(number of atoms)(atomic weight)

(FW of the compound)x 100

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Percent Composition

So the percentage of carbon in ethane is…

%C =(2)(12.0 amu)

(30.0 amu)

24.0 amu

30.0 amu= x 100

= 80.0%

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Moles

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Avogadro’s Number

• 6.02 x 1023

• 1 mole of 12C has a mass of 12 g.

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Molar Mass

• By definition, a molar mass is the mass of 1 mol of a substance (i.e., g/mol).– The molar mass of an element is the mass

number for the element that we find on the periodic table.

– The formula weight (in amu’s) will be the same number as the molar mass (in g/mol).

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Using Moles

Moles provide a bridge from the molecular scale to the real-world scale.

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Mole Relationships

• One mole of atoms, ions, or molecules contains Avogadro’s number of those particles.

• One mole of molecules or formula units contains Avogadro’s number times the number of atoms or ions of each element in the compound.

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Finding Empirical Formulas

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Calculating Empirical Formulas

One can calculate the empirical formula from the percent composition.

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The compound para-aminobenzoic acid (you may have seen it listed as PABA on your bottle of sunscreen) is composed of carbon (61.31%), hydrogen (5.14%), nitrogen (10.21%), and oxygen (23.33%). Find the empirical formula of PABA.

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Assuming 100.00 g of para-aminobenzoic acid,

C: 61.31 g x = 5.105 mol C

H: 5.14 g x = 5.09 mol H

N: 10.21 g x = 0.7288 mol N

O: 23.33 g x = 1.456 mol O

1 mol12.01 g

1 mol14.01 g

1 mol1.01 g

1 mol16.00 g

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Calculate the mole ratio by dividing by the smallest number of moles:

C: = 7.005 7

H: = 6.984 7

N: = 1.000

O: = 2.001 2

5.105 mol0.7288 mol

5.09 mol0.7288 mol

0.7288 mol0.7288 mol

1.458 mol0.7288 mol

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These are the subscripts for the empirical formula:

C7H7NO2

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Combustion Analysis

• Compounds containing C, H and O are routinely analyzed through combustion in a chamber like this.– C is determined from the mass of CO2 produced.

– H is determined from the mass of H2O produced.

– O is determined by difference after the C and H have been determined.

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Elemental Analyses

Compounds containing other elements are analyzed using methods analogous to those used for C, H and O.

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Stoichiometric Calculations

The coefficients in the balanced equation give the ratio of moles of reactants and products.

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Starting with the mass of Substance A you can use the ratio of the coefficients of A and B to calculate the mass of Substance B formed (if it’s a product) or used (if it’s a reactant).

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Starting with 1.00 g of C6H12O6… we calculate the moles of C6H12O6…use the coefficients to find the moles of H2O…and then turn the moles of water to grams.

C6H12O6 + 6 O2 6 CO2 + 6 H2O

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Limiting Reactants

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How Many Cookies Can I Make?

• You can make cookies until you run out of one of the ingredients.

• Once this family runs out of sugar, they will stop making cookies (at least any cookies you would want to eat).

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How Many Cookies Can I Make?

• In this example the sugar would be the limiting reactant, because it will limit the amount of cookies you can make.

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Limiting Reactants

• The limiting reactant is the reactant present in the smallest stoichiometric amount.– In other words, it’s the reactant you’ll run out of first (in

this case, the H2).

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Limiting Reactants

In the example below, the O2 would be the excess reagent.

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Theoretical Yield

• The theoretical yield is the maximum amount of product that can be made.– In other words it’s the amount of product

possible as calculated through the stoichiometry problem.

• This is different from the actual yield, which is the amount one actually produces and measures.

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Percent Yield

One finds the percent yield by comparing the amount actually obtained (actual yield) to the amount it was possible to make (theoretical yield).

Actual YieldTheoretical YieldPercent Yield = x 100

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Solutions

• Solutions are defined as homogeneous mixtures of two or more pure substances.

• The solvent is present in greatest abundance.

• All other substances are solutes.

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Dissociation

• When an ionic substance dissolves in water, the solvent pulls the individual ions from the crystal and solvates them.

• This process is called dissociation.

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Dissociation

• An electrolyte is a substances that dissociates into ions when dissolved in water.

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Electrolytes

• An electrolyte is a substances that dissociates into ions when dissolved in water.

• A nonelectrolyte may dissolve in water, but it does not dissociate into ions when it does so.

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Electrolytes and Nonelectrolytes

Soluble ionic compounds tend to be electrolytes.

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Electrolytes and Nonelectrolytes

Molecular compounds tend to be nonelectrolytes, except for acids and bases.

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Electrolytes

• A strong electrolyte dissociates completely when dissolved in water.

• A weak electrolyte only dissociates partially when dissolved in water.

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Strong Electrolytes Are…

• Strong acids• Strong bases

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Strong Electrolytes Are…

• Strong acids• Strong bases• Soluble ionic salts

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Precipitation Reactions

When one mixes ions that form compounds that are insoluble (as could be predicted by the solubility guidelines), a precipitate is formed.

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Metathesis (Exchange) Reactions• Metathesis comes from a Greek word that

means “to transpose.”

AgNO3 (aq) + KCl (aq) AgCl (s) + KNO3 (aq)

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Metathesis (Exchange) Reactions• Metathesis comes from a Greek word that

means “to transpose.”• It appears the ions in the reactant

compounds exchange, or transpose, ions.


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Cl- Br- I- NO3- SO4

2- CO32- PO4

3-

Li+ S S S S S S S S

Na+ S S S S S S S S

K+ S S S S S S S S

Mg2+ NS S S S S S NS NS

Ca2+ S S S S S S NS NS

Sr2+ S S S S S NS NS NS

Ba2+ S S S S S NS NS NS

Fe2+ NS S S S S S NS NS

Fe3+ NS S S S S S NS NS

Ni2+ NS S S S S S NS NS

Cu+ NS S S S S S NS NS

Cu2+ NS S S S S S NS NS

Al3+ NS S S S S S NS NS

Zn2+ NS S S S S S NS NS

Ag+ NS NS NS NS S S NS NS

Pb2+ NS NS NS NS S NS NS NS

Solubility of different compounds(NS = non soluble in water, S = soluble in water)

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Solution Chemistry

• It is helpful to pay attention to exactly what species are present in a reaction mixture (i.e., solid, liquid, gas, aqueous solution).

• If we are to understand reactivity, we must be aware of just what is changing during the course of a reaction.

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Molecular Equation

The molecular equation lists the reactants and products in their molecular form.


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Ionic Equation• In the ionic equation all strong electrolytes (strong

acids, strong bases, and soluble ionic salts) are dissociated into their ions.

• This more accurately reflects the species that are found in the reaction mixture.

Ag+ (aq) + NO3- (aq) + K+ (aq) + Cl- (aq)

AgCl (s) + K+ (aq) + NO3- (aq)

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Net Ionic Equation• To form the net ionic equation, cross out anything

that does not change from the left side of the equation to the right.

Ag+(aq) + NO3-(aq) + K+(aq) + Cl-(aq)

AgCl (s) + K+(aq) + NO3-(aq)

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• The only things left in the equation are those things that change (i.e., react) during the course of the reaction.

Ag+(aq) + Cl-(aq) AgCl (s)

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• The only things left in the equation are those things that change (i.e., react) during the course of the reaction.

• Those things that didn’t change (and were deleted from the net ionic equation) are called spectator ions.

Ag+(aq) + NO3-(aq) + K+(aq) + Cl-(aq)

AgCl (s) + K+(aq) + NO3-(aq)

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Writing Net Ionic Equations

1. Write a balanced molecular equation.

2. Dissociate all strong electrolytes.

3. Cross out anything that remains unchanged from the left side to the right side of the equation.

4. Write the net ionic equation with the species that remain.

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Acids

• Arrhenius defined acids as substances that increase the concentration of H+ when dissolved in water.

• Brønsted and Lowry defined them as proton donors.

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Acids

There are only seven strong acids:• Hydrochloric (HCl)• Hydrobromic (HBr)• Hydroiodic (HI)

• Nitric (HNO3)

• Sulfuric (H2SO4)

• Chloric (HClO3)

• Perchloric (HClO4)

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Bases• Arrhenius defined bases

as substances that increase the concentration of OH− when dissolved in water.

• Brønsted and Lowry defined them as proton acceptors.

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Bases

The strong bases are the soluble metal salts of hydroxide ion:• Alkali metals• Calcium• Strontium• Barium

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Acid-Base Reactions

In an acid-base reaction, the acid donates a proton (H+) to the base.

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Neutralization ReactionsGenerally, when solutions of an acid and a base are combined, the products are a salt and water.

CH3COOH (aq) + NaOH (aq) CH3COONa (aq) + H2O (l)

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Neutralization ReactionsWhen a strong acid reacts with a strong base, the net ionic equation is…

HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)

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H+ (aq) + Cl- (aq) + Na+ (aq) + OH-(aq)

Na+ (aq) + Cl- (aq) + H2O (l)

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H+ (aq) + Cl- (aq) + Na+ (aq) + OH-(aq)

Na+ (aq) + Cl- (aq) + H2O (l)

H+ (aq) + OH- (aq) H2O (l)

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Gas-Forming Reactions

• Some metathesis reactions do not give the product expected.

• In this reaction, the expected product (H2CO3) decomposes to give a gaseous product (CO2).

CaCO3 (s) + HCl (aq) CaCl2 (aq) + CO2 (g) + H2O (l)

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When a carbonate or bicarbonate reacts with an acid, the products are a salt, carbon dioxide, and water.

CaCO3 (s) + HCl (aq) CaCl2 (aq) + CO2 (g) + H2O (l)

NaHCO3 (aq) + HBr (aq) NaBr (aq) + CO2 (g) + H2O (l)

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Similarly, when a sulfite reacts with an acid, the products are a salt, sulfur dioxide, and water.

SrSO3 (s) + 2 HI (aq) SrI2 (aq) + SO2 (g) + H2O (l)

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• This reaction gives the predicted product, but you had better carry it out in the hood, or you will be very unpopular!

• But just as in the previous examples, a gas is formed as a product of this reaction.

Na2S (aq) + H2SO4 (aq) Na2SO4 (aq) + H2S (g)

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Oxidation-Reduction Reactions

• An oxidation occurs when an atom or ion loses electrons.

• A reduction occurs when an atom or ion gains electrons.

• One cannot occur without the other.

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Oxidation Numbers

To determine if an oxidation-reduction reaction has occurred, we assign an oxidation number to each element in a neutral compound or charged entity.

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Oxidation Numbers

• Elements in their elemental form have an oxidation number of 0.

• The oxidation number of a monatomic ion is the same as its charge.

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Oxidation Numbers

• Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions.Oxygen has an oxidation number of −2,

except in the peroxide ion in which it has an oxidation number of −1.

Hydrogen is −1 when bonded to a metal, +1 when bonded to a nonmetal.

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Oxidation Numbers

• Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions.Fluorine always has an oxidation number

of −1.The other halogens have an oxidation

number of −1 when they are negative; they can have positive oxidation numbers, however, most notably in oxyanions.

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Oxidation Numbers

• The sum of the oxidation numbers in a neutral compound is 0.

• The sum of the oxidation numbers in a polyatomic ion is the charge on the ion.

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Displacement Reactions

• In displacement reactions, ions oxidize an element.

• The ions, then, are reduced.

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In this reaction,

silver ions oxidize

copper metal.

Cu (s) + 2 Ag+ (aq) Cu2+ (aq) + 2 Ag (s)

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The reverse reaction,

however, does not

occur.

Cu2+ (aq) + 2 Ag (s) Cu (s) + 2 Ag+ (aq) x

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Activity Series

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Molarity• Two solutions can contain the same

compounds but be quite different because the proportions of those compounds are different.

• Molarity is one way to measure the concentration of a solution.

moles of solute

volume of solution in litersMolarity (M) =

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Mixing a Solution

• To create a solution of a known molarity, one weighs out a known mass (and, therefore, number of moles) of the solute.

• The solute is added to a volumetric flask, and solvent is added to the line on the neck of the flask.

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Dilution• One can also dilute a more concentrated

solution by– Using a pipet to deliver a volume of the solution to

a new volumetric flask, and– Adding solvent to the line on the neck of the new

flask.

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DilutionThe molarity of the new solution can be determined from the equation

Mc Vc = Md Vd,

where Mc and Md are the molarity of the concentrated and dilute solutions, respectively, and Vc and Vd are the volumes of the two solutions.

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Using Molarities inStoichiometric Calculations

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Titration

Titration is an analytical technique in which one can calculate the concentration of a solute in a solution.

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Chapter 5Thermochemistry

John D. BookstaverSt. Charles Community College

Cottleville, MO

Chemistry, The Central Science, 11th editionTheodore L. Brown; H. Eugene LeMay, Jr.;

and Bruce E. Bursten

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Energy

• Energy is the ability to do work or transfer heat.– Energy used to cause an object that has

mass to move is called work.– Energy used to cause the temperature of

an object to rise is called heat.

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Potential Energy

Potential energy is energy an object possesses by virtue of its position or chemical composition.

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Kinetic Energy

Kinetic energy is energy an object possesses by virtue of its motion.

12

KE = mv2

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Units of Energy

• The SI unit of energy is the joule (J).

• An older, non-SI unit is still in widespread use: the calorie (cal).

1 cal = 4.184 J

1 J = 1 kg m2

s2

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Definitions:System and Surroundings

• The system includes the molecules we want to study (here, the hydrogen and oxygen molecules).

• The surroundings are everything else (here, the cylinder and piston).

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Definitions: Work

• Energy used to move an object over some distance is work.

• w = F dwhere w is work, F is the force, and d is the distance over which the force is exerted.

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Heat

• Energy can also be transferred as heat.

• Heat flows from warmer objects to cooler objects.

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Conversion of Energy

• Energy can be converted from one type to another.

• For example, the cyclist above has potential energy as she sits on top of the hill.

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Conversion of Energy

• As she coasts down the hill, her potential energy is converted to kinetic energy.

• At the bottom, all the potential energy she had at the top of the hill is now kinetic energy.

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First Law of Thermodynamics• Energy is neither created nor destroyed.• In other words, the total energy of the universe is

a constant; if the system loses energy, it must be gained by the surroundings, and vice versa.

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Internal EnergyThe internal energy of a system is the sum of all kinetic and potential energies of all components of the system; we call it E.

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Internal EnergyBy definition, the change in internal energy, E, is the final energy of the system minus the initial energy of the system:

E = Efinal − Einitial

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Changes in Internal Energy

• If E > 0, Efinal > Einitial

– Therefore, the system absorbed energy from the surroundings.

– This energy change is called endergonic.

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• If E < 0, Efinal < Einitial

– Therefore, the system released energy to the surroundings.

– This energy change is called exergonic.

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• When energy is exchanged between the system and the surroundings, it is exchanged as either heat (q) or work (w).

• That is, E = q + w.

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E, q, w, and Their Signs

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Exchange of Heat between System and Surroundings

• When heat is absorbed by the system from the surroundings, the process is endothermic.

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Exchange of Heat between System and Surroundings

• When heat is absorbed by the system from the surroundings, the process is endothermic.

• When heat is released by the system into the surroundings, the process is exothermic.

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State Functions

Usually we have no way of knowing the internal energy of a system; finding that value is simply too complex a problem.

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State Functions• However, we do know that the internal energy

of a system is independent of the path by which the system achieved that state.– In the system below, the water could have reached

room temperature from either direction.

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State Functions• Therefore, internal energy is a state function.• It depends only on the present state of the

system, not on the path by which the system arrived at that state.

• And so, E depends only on Einitial and Efinal.

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State Functions

• However, q and w are not state functions.

• Whether the battery is shorted out or is discharged by running the fan, its E is the same.– But q and w are different

in the two cases.

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Work

Usually in an open container the only work done is by a gas pushing on the surroundings (or by the surroundings pushing on the gas).

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WorkWe can measure the work done by the gas if the reaction is done in a vessel that has been fitted with a piston.

w = -PV

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Enthalpy

• If a process takes place at constant pressure (as the majority of processes we study do) and the only work done is this pressure-volume work, we can account for heat flow during the process by measuring the enthalpy of the system.

• Enthalpy is the internal energy plus the product of pressure and volume:

H = E + PV

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Enthalpy

• When the system changes at constant pressure, the change in enthalpy, H, is

H = (E + PV)

• This can be written

H = E + PV

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Enthalpy

• Since E = q + w and w = -PV, we can substitute these into the enthalpy expression:

H = E + PV

H = (q+w) − w

H = q

• So, at constant pressure, the change in enthalpy is the heat gained or lost.

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Endothermicity and Exothermicity

• A process is endothermic when H is positive.

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Endothermicity and Exothermicity

• A process is endothermic when H is positive.

• A process is exothermic when H is negative.

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Enthalpy of Reaction

The change in enthalpy, H, is the enthalpy of the products minus the enthalpy of the reactants:

H = Hproducts − Hreactants

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Enthalpy of Reaction

This quantity, H, is called the enthalpy of reaction, or the heat of reaction.

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The Truth about Enthalpy

1. Enthalpy is an extensive property.

2. H for a reaction in the forward direction is equal in size, but opposite in sign, to H for the reverse reaction.

3. H for a reaction depends on the state of the products and the state of the reactants.

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Calorimetry

Since we cannot know the exact enthalpy of the reactants and products, we measure H through calorimetry, the measurement of heat flow.

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Heat Capacity and Specific Heat

The amount of energy required to raise the temperature of a substance by 1 K (1C) is its heat capacity.

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We define specific heat capacity (or simply specific heat) as the amount of energy required to raise the temperature of 1 g of a substance by 1 K.

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Specific heat, then, is

Specific heat =heat transferred

mass temperature change

s =q

m T

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Constant Pressure Calorimetry

By carrying out a reaction in aqueous solution in a simple calorimeter such as this one, one can indirectly measure the heat change for the system by measuring the heat change for the water in the calorimeter.

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Constant Pressure Calorimetry

Because the specific heat for water is well known (4.184 J/g-K), we can measure H for the reaction with this equation:

q = m s T

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Bomb Calorimetry

• Reactions can be carried out in a sealed “bomb” such as this one.

• The heat absorbed (or released) by the water is a very good approximation of the enthalpy change for the reaction.

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Bomb Calorimetry

• Because the volume in the bomb calorimeter is constant, what is measured is really the change in internal energy, E, not H.

• For most reactions, the difference is very small.

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Enthalpies of Formation

Enthalpy of formation, Hf, is ….

the enthalpy change for the reaction in which a compound is made from its constituent elements in their elemental forms.

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Hess's Law

If a reaction is carried out in a series of steps, H

for the overall reaction = the sum of the enthalpy changes for individual steps.

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Calculation of H

• Imagine this as occurringin three steps:

C3H8 (g) + 5 O2 (g) 3 CO2 (g) + 4 H2O (l)

C3H8 (g) 3 C (graphite) + 4 H2 (g)

3 C (graphite) + 3 O2 (g) 3 CO2 (g)

4 H2 (g) + 2 O2 (g) 4 H2O (l)

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C3H8 (g) + 5 O2 (g) 3 CO2 (g) + 4 H2O (l)

C3H8 (g) 3 C (graphite) + 4 H2 (g)

3 C (graphite) + 3 O2 (g) 3 CO2 (g)

4 H2 (g) + 2 O2 (g) 4 H2O (l)

C3H8 (g) + 5 O2 (g) 3 CO2 (g) + 4 H2O (l)

Calculation of H

• The sum of these equations is:

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Calculation of H

We can use Hess’s law in this way:

H = nHf°products – mHf° reactants

where n and m are the stoichiometric coefficients.

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H = [3(-393.5 kJ) + 4(-285.8 kJ)] – [1(-103.85 kJ) + 5(0 kJ)]

= [(-1180.5 kJ) + (-1143.2 kJ)] – [(-103.85 kJ) + (0 kJ)]= (-2323.7 kJ) – (-103.85 kJ) = -2219.9 kJ

C3H8 (g) + 5 O2 (g) 3 CO2 (g) + 4 H2O (l)

Calculation of H

Hf of the most stable

Form of any element Is 0 ...no formationNeeded.

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Hess’s Law

H is well known for many reactions, and it is inconvenient to measure H for every reaction in which we are interested.

• However, we can estimate H using published H values and the properties of enthalpy.

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Hess’s Law

Hess’s law states that “[i]f a reaction is carried out in a series of steps, H for the overall reaction will be equal to the sum of the enthalpy changes for the individual steps.”

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Hess’s Law

Because H is a state function, the total enthalpy change depends only on the initial state of the reactants and the final state of the products.

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Energy in FoodsMost of the fuel in the food we eat comes from carbohydrates and fats.

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Energy in Fuels

The vast majority of the energy consumed in this country comes from fossil fuels.

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Chapter 6Electronic Structureof Atoms

Chemistry, The Central Science, 11th editionTheodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten

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The Nature of Energy

• Why an object can glow when its temperature increases?

• The wave nature of light does not explain it

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Waves

• To understand the electronic structure of atoms, one must understand the nature of electromagnetic radiation.

• The distance between corresponding points on adjacent waves is the wavelength ().

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Waves

• The number of waves passing a given point per unit of time is the frequency ().

• For waves traveling at the same velocity, the longer the wavelength, the smaller the frequency.

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Electromagnetic Radiation

• All electromagnetic radiation travels at the same velocity: the speed of light (c), 3.00 108 m/s.

• Therefore,c =

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Another mystery in the early 20th century involved the emission spectra observed from energy emitted by atoms and molecules.

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• For atoms and molecules one does not observe a continuous spectrum, as one gets from a white light source.

• Only a line spectrum of discrete wavelengths is observed.

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• Max Planck explained it by assuming that energy comes in packets called quanta.

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• Einstein used this assumption to explain the photoelectric effect.

• He concluded that energy is proportional to frequency:

E = hwhere h is Planck’s constant, 6.626 10−34 J-s.

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• Therefore, if one knows the wavelength of light, one can calculate the energy in one photon, or packet, of that light:

c = E = h

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• Niels Bohr adopted Planck’s assumption and explained these phenomena in this way:1. Electrons in an atom can only

occupy certain orbits (corresponding to certain energies).

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• Niels Bohr adopted Planck’s assumption and explained these phenomena in this way:1. Electrons in permitted orbits

have specific, “allowed” energies; these energies will not be radiated from the atom.

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• Niels Bohr adopted Planck’s assumption and explained these phenomena in this way:1. Energy is only absorbed or

emitted in such a way as to move an electron from one “allowed” energy state to another; the energy is defined by

E = h

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The energy absorbed or emitted from the process of electron promotion or demotion can be calculated by the equation:

E = −hcRH ( )1nf

2

1ni

2-

where RH is the Rydberg constant, and ni and nf are the initial and final energy levels of the electron.

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The Wave Nature of Matter

• Louis de Broglie posited that if light can have material properties, matter should exhibit wave properties.

• He demonstrated that the relationship between mass and wavelength was

=hmv

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The Uncertainty Principle

• Heisenberg showed that the more precisely the momentum of a particle is known, the less precisely is its position known:

• In many cases, our uncertainty of the whereabouts of an electron is greater than the size of the atom itself!

(x) (mv) h4

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Quantum Mechanics

• Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated.

• It is known as quantum mechanics.

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Schrödinger equation

Time dependent form

Time independent form

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Quantum Mechanics

• The wave equation is designated with a lower case Greek psi ().

• The square of the wave equation, 2, gives a probability density map of where an electron has a certain statistical likelihood of being at any given instant in time.

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Quantum Numbers

• Solving the wave equation gives a set of wave functions, or orbitals, and their corresponding energies.

• Each orbital describes a spatial distribution of electron density.

• An orbital is described by a set of three quantum numbers.

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Principal Quantum Number (n)

• The principal quantum number, n, describes the energy level on which the orbital resides.

• The values of n are integers ≥ 1.

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Angular Momentum Quantum Number (l)

• This quantum number defines the shape of the orbital.

• Allowed values of l are integers ranging from 0 to n − 1.

• We use letter designations to communicate the different values of l and, therefore, the shapes and types of orbitals.

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Angular Momentum Quantum Number (l)

Value of l 0 1 2 3

Type of orbital s p d f

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Magnetic Quantum Number (ml)

• The magnetic quantum number describes the three-dimensional orientation of the orbital.

• Allowed values of ml are integers ranging from -l to l:

−l ≤ ml ≤ l.• Therefore, on any given energy level,

there can be up to 1 s orbital, 3 p orbitals, 5 d orbitals, 7 f orbitals, etc.

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Magnetic Quantum Number (ml)

• Orbitals with the same value of n form a shell.• Different orbital types within a shell are

subshells.

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s Orbitals

• The value of l for s orbitals is 0.

• They are spherical in shape.

• The radius of the sphere increases with the value of n.

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s Orbitals

Observing a graph of probabilities of finding an electron versus distance from the nucleus, we see that s orbitals possess n−1 nodes, or regions where there is 0 probability of finding an electron.

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p Orbitals

• The value of l for p orbitals is 1.• They have two lobes with a node between

them.

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d Orbitals

• The value of l for a d orbital is 2.

• Four of the five d orbitals have 4 lobes; the other resembles a p orbital with a doughnut around the center.

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Energies of Orbitals

• For a one-electron hydrogen atom, orbitals on the same energy level have the same energy.

• That is, they are degenerate.

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Energies of Orbitals

• As the number of electrons increases, though, so does the repulsion between them.

• Therefore, in many-electron atoms, orbitals on the same energy level are no longer degenerate.

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Spin Quantum Number, ms

• In the 1920s, it was discovered that two electrons in the same orbital do not have exactly the same energy.

• The “spin” of an electron describes its magnetic field, which affects its energy.

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Spin Quantum Number, ms

• This led to a fourth quantum number, the spin quantum number, ms.

• The spin quantum number has only 2 allowed values: +1/2 and −1/2.

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Pauli Exclusion Principle

• No two electrons in the same atom can have exactly the same energy.

• Therefore, no two electrons in the same atom can have identical sets of quantum numbers.

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Electron Configurations

• This shows the distribution of all electrons in an atom.

• Each component consists of – A number denoting the

energy level,

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• This shows the distribution of all electrons in an atom


energy level,– A letter denoting the type

of orbital,

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• This shows the distribution of all electrons in an atom.


energy level,– A letter denoting the type

of orbital,– A superscript denoting

the number of electrons in those orbitals.

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First Klechkovsky’s Rule

As atomic number increases, filling of orbitals in the atom goes from orbitals with the smaller sum of n and l (n+l) to orbitals with the larger sum of n and l (n+l).

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Orbital Diagrams

• Each box in the diagram represents one orbital.

• Half-arrows represent the electrons.

• The direction of the arrow represents the relative spin of the electron.

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Hund’s Rule

“For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.”

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Periodic Table

• We fill orbitals in increasing order of energy.

• Different blocks on the periodic table (shaded in different colors in this chart) correspond to different types of orbitals.

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Some Anomalies

For instance, the electron configuration for copper is

[Ar] 4s1 3d5

rather than the expected

[Ar] 4s2 3d4.

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Second Klechkovsky’s Rule

When the sum of n and l (n+l) is identical, the filling of orbitals goes in the direction of rising of the main quantum number (n).

Exceptions: Cr, Cu, Nb, Mo, Ru, Rh, Pd, Ag, Pt, Au

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Some Anomalies

Some irregularities occur when there are enough electrons to half-fill s and d orbitals on a given row.

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Some Anomalies

• This occurs because the 4s and 3d orbitals are very close in energy.

• These anomalies occur in f-block atoms, as well.

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Chapter 7Periodic Properties

of the Elements




Cottleville, MO

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Development of Periodic Table

• Elements in the same group generally have similar chemical properties.

• Physical properties are not necessarily similar, however.

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Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped.

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Mendeleev, for instance, predicted the discovery of germanium (which he called eka-silicon) as an element with an atomic weight between that of zinc and arsenic, but with chemical properties similar to those of silicon.

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Periodic Trends

• In this chapter, we will rationalize observed trends in– Sizes of atoms and ions.– Ionization energy.– Electron affinity.

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Effective Nuclear Charge

• In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons.

• The nuclear charge that an electron experiences depends on both factors.

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Effective Nuclear Charge

The effective nuclear charge, Zeff, is found this way:

Zeff = Z − S

where Z is the atomic number and S is a screening constant, usually close to the number of inner electrons.

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Group

Other electrons in

the same group

Electrons in group(s) with principal quantum number n-1

Electrons in all group(s) with principal quantum number

< n-1

[1s] 0.3 N/A N/A

[ns,np] 0.35 0.85 1

[nd] or [nf] 0.35 1 1

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What Is the Size of an Atom?

The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.

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Sizes of Atoms

Bonding atomic radius tends to… …decrease from left to

right across a row(due to increasing Zeff).

…increase from top to bottom of a column

(due to increasing value of n).

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Sizes of Ions

• Ionic size depends upon:– The nuclear

charge.– The number of

electrons.– The orbitals in

which electrons reside.

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Sizes of Ions

• Cations are smaller than their parent atoms.– The outermost

electron is removed and repulsions between electrons are reduced.

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Sizes of Ions

• Anions are larger than their parent atoms.– Electrons are

added and repulsions between electrons are increased.

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Sizes of Ions

• Ions increase in size as you go down a column.– This is due to

increasing value of n.

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Sizes of Ions

• In an isoelectronic series, ions have the same number of electrons.

• Ionic size decreases with an increasing nuclear charge.

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Ionization Energy

• The ionization energy is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion.– The first ionization energy is that energy

required to remove first electron.– The second ionization energy is that

energy required to remove second electron, etc.

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Ionization Energy• It requires more energy to remove each

successive electron.• When all valence electrons have been removed,

the ionization energy takes a quantum leap.

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Trends in First Ionization Energies

• As one goes down a column, less energy is required to remove the first electron.– For atoms in the same

group, Zeff is essentially the same, but the valence electrons are farther from the nucleus.

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• Generally, as one goes across a row, it gets harder to remove an electron.– As you go from left to

right, Zeff increases.

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However, there are two apparent discontinuities in this trend.

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• The first occurs between Groups IIA and IIIA.

• In this case the electron is removed from a p-orbital rather than an s-orbital.– The electron removed

is farther from nucleus.– There is also a small

amount of repulsion by the s electrons.

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• The second occurs between Groups VA and VIA.– The electron removed

comes from doubly occupied orbital.

– Repulsion from the other electron in the orbital aids in its removal.

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Electron Affinity

Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom:

Cl + e− Cl−

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Trends in Electron Affinity

In general, electron affinity becomes more exothermic as you go from left to right across a row.

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There are again, however, two discontinuities in this trend.

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• The first occurs between Groups IA and IIA.– The added electron

must go in a p-orbital, not an s-orbital.

– The electron is farther from nucleus and feels repulsion from the s-electrons.

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• The second occurs between Groups IVA and VA.– Group VA has no

empty orbitals.– The extra electron

must go into an already occupied orbital, creating repulsion.

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Properties of Metal, Nonmetals,

and Metalloids

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Metals versus Nonmetals

Differences between metals and nonmetals tend to revolve around these properties.

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Metals versus Nonmetals

• Metals tend to form cations.• Nonmetals tend to form anions.

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Metals

They tend to be lustrous, malleable, ductile, and good conductors of heat and electricity.

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Metals

• Compounds formed between metals and nonmetals tend to be ionic.

• Metal oxides tend to be basic.

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Nonmetals

• These are dull, brittle substances that are poor conductors of heat and electricity.

• They tend to gain electrons in reactions with metals to acquire a noble gas configuration.

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Nonmetals

• Substances containing only nonmetals are molecular compounds.

• Most nonmetal oxides are acidic.

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Metalloids

• These have some characteristics of metals and some of nonmetals.

• For instance, silicon looks shiny, but is brittle and fairly poor conductor.

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Group Trends

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Alkali Metals

• Alkali metals are soft, metallic solids.

• The name comes from the Arabic word for ashes.

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Alkali Metals

• They are found only in compounds in nature, not in their elemental forms.

• They have low densities and melting points.• They also have low ionization energies.

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Alkali Metals

Their reactions with water are famously exothermic.

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Alkali Metals• Alkali metals (except Li) react with oxygen to

form peroxides.• K, Rb, and Cs also form superoxides:

K + O2 KO2

• They produce bright colors when placed in a flame.

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Alkaline Earth Metals

• Alkaline earth metals have higher densities and melting points than alkali metals.

• Their ionization energies are low, but not as low as those of alkali metals.

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Alkaline Earth Metals

• Beryllium does not react with water and magnesium reacts only with steam, but the others react readily with water.

• Reactivity tends to increase as you go down the group.

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Group 6A

• Oxygen, sulfur, and selenium are nonmetals.• Tellurium is a metalloid.• The radioactive polonium is a metal.

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Oxygen• There are two allotropes of

oxygen:– O2

– O3, ozone

• There can be three anions:– O2−, oxide– O2

2−, peroxide– O2

1−, superoxide

• It tends to take electrons from other elements (oxidation).

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Sulfur

• Sulfur is a weaker oxidizer than oxygen.

• The most stable allotrope is S8, a ringed molecule.

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Group VIIA: Halogens

• The halogens are prototypical nonmetals.• The name comes from the Greek words halos

and gennao: “salt formers”.

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Group VIIA: Halogens

• They have large, negative electron affinities.– Therefore, they tend to

oxidize other elements easily.

• They react directly with metals to form metal halides.

• Chlorine is added to water supplies to serve as a disinfectant

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Group VIIIA: Noble Gases

• The noble gases have astronomical ionization energies.

• Their electron affinities are positive.– Therefore, they are relatively unreactive.

• They are found as monatomic gases.

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Group VIIIA: Noble Gases

• Xe forms three compounds:– XeF2

– XeF4 (at right)

– XeF6

• Kr forms only one stable compound:– KrF2

• The unstable HArF was synthesized in 2000.

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Chapter 8Concepts of Chemical

Bonding

Chemistry, The Central Science, 11th editionTheodore L. Brown, H. Eugene LeMay, Jr.,



Cottleville, MO

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Chemical Bonds• Three basic types of

bonds– Ionic

• Electrostatic attraction between ions

– Covalent• Sharing of electrons

– Metallic• Metal atoms bonded to

several other atoms

• Lewis Symbols• Octet Rule

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Ionic Bonding

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Energetics of Ionic Bonding

As we saw in the last chapter, it takes 495 kJ/mol to remove electrons from sodium.

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We get 349 kJ/mol back by giving electrons to chlorine.

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But these numbers don’t explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so exothermic!

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• There must be a third piece to the puzzle.

• What is as yet unaccounted for is the electrostatic attraction between the newly-formed sodium cation and chloride anion.

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Lattice Energy

• This third piece of the puzzle is the lattice energy:– The energy required to completely separate a mole

of a solid ionic compound into its gaseous ions.

• The energy associated with electrostatic interactions is governed by Coulomb’s law:

Eel = Q1Q2

d

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Lattice Energy

• Lattice energy, then, increases with the charge on the ions.

• It also increases with decreasing size of ions.

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Energetics of Ionic Bonding:Born-Haber cycle

By accounting for all three energies (ionization energy, electron affinity, and lattice energy), we can get a good idea of the energetics involved in such a process.

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• These phenomena also helps explain the “octet rule.”

• Metals, for instance, tend to stop losing electrons once they attain a noble gas configuration because energy would be expended that cannot be overcome by lattice energies.

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Covalent Bonding

• In covalent bonds atoms share electrons.

• There are several electrostatic interactions in these bonds:– Attractions between electrons

and nuclei– Repulsions between electrons– Repulsions between nuclei

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Polar Covalent Bonds

• Though atoms often form compounds by sharing electrons, the electrons are not always shared equally.

• Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does.

• Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.

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Electronegativity

• Electronegativity is the ability of atoms in a molecule to attract electrons to themselves.

• On the periodic chart, electronegativity increases as you go…– …from left to right across

a row.– …from the bottom to the

top of a column.

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• When two atoms share electrons unequally, a bond dipole results.

• The dipole moment, , produced by two equal but opposite charges separated by a distance, r, is calculated:

= Qr• It is measured in debyes (D).

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The greater the difference in electronegativity, the more polar is the bond.

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Lewis Structures

Lewis structures are representations of molecules showing all electrons, bonding and nonbonding.

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Writing Lewis Structures

1. Find the sum of valence electrons of all atoms in the polyatomic ion or molecule.– If it is an anion, add one

electron for each negative charge.

– If it is a cation, subtract one electron for each positive charge.

PCl3

5 + 3(7) = 26

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2. The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds.

Keep track of the electrons:

26 - 6 = 20

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3. Fill the octets of the outer atoms.


26 - 6 = 20; 20 - 18 = 2

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4. Fill the octet of the central atom.


26 - 6 = 20; 20 - 18 = 2; 2 - 2 = 0

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5. If you run out of electrons before the central atom has an octet…

…form multiple bonds until it does.

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• Then assign formal charges.– For each atom, count the electrons in lone pairs and

half the electrons it shares with other atoms.– Subtract that from the number of valence electrons for

that atom: the difference is its formal charge.

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• The best Lewis structure…– …is the one with the fewest charges.– …puts a negative charge on the most

electronegative atom.

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Resonance

This is the Lewis structure we would draw for ozone, O3. -

+

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Resonance

• But this is at odds with the true, observed structure of ozone, in which…– …both O-O bonds

are the same length.– …both outer

oxygens have a charge of -1/2.

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Resonance

• One Lewis structure cannot accurately depict a molecule like ozone.

• We use multiple structures, resonance structures, to describe the molecule.

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Resonance

Just as green is a synthesis of blue and yellow…

…ozone is a synthesis of these two resonance structures.

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Resonance

• In truth, the electrons that form the second C-O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon.

• They are not localized; they are delocalized.

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Resonance

• The organic compound benzene, C6H6, has two resonance structures.

• It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.

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Exceptions to the Octet Rule

• There are three types of ions or molecules that do not follow the octet rule:– Ions or molecules with an odd number of

electrons– Ions or molecules with less than an octet– Ions or molecules with more than eight

valence electrons (an expanded octet)

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Odd Number of Electrons

Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons.

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Fewer Than Eight Electrons

• Consider BF3:– Giving boron a filled octet places a negative

charge on the boron and a positive charge on fluorine.

– This would not be an accurate picture of the distribution of electrons in BF3.

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Therefore, structures that put a double bond between boron and fluorine are much less important than the one that leaves boron with only 6 valence electrons.

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The lesson is: if filling the octet of the central atom results in a negative charge on the central atom and a positive charge on the more electronegative outer atom, don’t fill the octet of the central atom.

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More Than Eight Electrons

• The only way PCl5 can exist is if phosphorus has 10 electrons around it.

• It is allowed to expand the octet of atoms on the 3rd row or below.– Presumably d orbitals in

these atoms participate in bonding.

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Even though we can draw a Lewis structure for the phosphate ion that has only 8 electrons around the central phosphorus, the better structure puts a double bond between the phosphorus and one of the oxygens.

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• This eliminates the charge on the phosphorus and the charge on one of the oxygens.

• The lesson is: when the central atom in on the 3rd row or below and expanding its octet eliminates some formal charges, do so.

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Covalent Bond Strength

• Most simply, the strength of a bond is measured by determining how much energy is required to break the bond.

• This is the bond enthalpy.• The bond enthalpy for a Cl-Cl bond, D(Cl-Cl),

is measured to be 242 kJ/mol.

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Average Bond Enthalpies

• This table lists the average bond enthalpies for many different types of bonds.

• Average bond enthalpies are positive, because bond breaking is an endothermic process.

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Average Bond Enthalpies

NOTE: These are average bond enthalpies, not absolute bond enthalpies; the C-H bonds in methane, CH4, will be a bit different than the C-H bond in chloroform, CHCl3.

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Enthalpies of Reaction

• Yet another way to estimate H for a reaction is to compare the bond enthalpies of bonds broken to the bond enthalpies of the new bonds formed.

• In other words, Hrxn = (bond enthalpies of bonds broken) -

(bond enthalpies of bonds formed)

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CH4 (g) + Cl2 (g)

CH3Cl (g) + HCl (g)

In this example, one C-H bond and one Cl-Cl bond are broken; one C-Cl and one H-Cl bond are formed.

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So,

H = [D(C-H) + D(Cl-Cl)] - [D(C-Cl) + D(H-Cl)]

= [(413 kJ) + (242 kJ)] - [(328 kJ) + (431 kJ)]

= (655 kJ) - (759 kJ)

= -104 kJ

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Bond Enthalpy and Bond Length

• We can also measure an average bond length for different bond types.

• As the number of bonds between two atoms increases, the bond length decreases.

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Chapter 9Molecular Geometriesand Bonding Theories

Chemistry, The Central Science, 11th editionTheodore L. Brown, H. Eugene LeMay, Jr.,



Cottleville, MO

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Molecular Shapes

• The shape of a molecule plays an important role in its reactivity.

• By noting the number of bonding and nonbonding electron pairs we can easily predict the shape of the molecule.

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What Determines the Shape of a Molecule?

• Simply put, electron pairs, whether they be bonding or nonbonding, repel each other.

• By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.

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Electron Domains

• We can refer to the electron pairs as electron domains.

• In a double or triple bond, all electrons shared between those two atoms are on the same side of the central atom; therefore, they count as one electron domain.

• The central atom in this molecule, A, has four electron domains.

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Valence Shell Electron Pair Repulsion Theory (VSEPR)

“The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.”

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Electron-Domain Geometries

These are the electron-domain geometries for two through six electron domains around a central atom.

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Electron-Domain Geometries

• All one must do is count the number of electron domains in the Lewis structure.

• The geometry will be that which corresponds to the number of electron domains.

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Molecular Geometries

• The electron-domain geometry is often not the shape of the molecule, however.

• The molecular geometry is that defined by the positions of only the atoms in the molecules, not the nonbonding pairs.

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Molecular Geometries

Within each electron domain, then, there might be more than one molecular geometry.

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Linear Electron Domain

• In the linear domain, there is only one molecular geometry: linear.

• NOTE: If there are only two atoms in the molecule, the molecule will be linear no matter what the electron domain is.

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Trigonal Planar Electron Domain

• There are two molecular geometries:– Trigonal planar, if all the electron domains are

bonding,– Bent, if one of the domains is a nonbonding pair.

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Nonbonding Pairs and Bond Angle

• Nonbonding pairs are physically larger than bonding pairs.

• Therefore, their repulsions are greater; this tends to decrease bond angles in a molecule.

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Multiple Bonds and Bond Angles

• Double and triple bonds place greater electron density on one side of the central atom than do single bonds.

• Therefore, they also affect bond angles.

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Tetrahedral Electron Domain

• There are three molecular geometries:– Tetrahedral, if all are bonding pairs,– Trigonal pyramidal if one is a nonbonding pair,– Bent if there are two nonbonding pairs.

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Trigonal Bipyramidal Electron Domain

• There are two distinct positions in this geometry:– Axial– Equatorial

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Lower-energy conformations result from having nonbonding electron pairs in equatorial, rather than axial, positions in this geometry.

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• There are four distinct molecular geometries in this domain:– Trigonal bipyramidal– Seesaw– T-shaped– Linear

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Octahedral Electron Domain

• All positions are equivalent in the octahedral domain.

• There are three molecular geometries:– Octahedral– Square pyramidal– Square planar

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Larger Molecules

In larger molecules, it makes more sense to talk about the geometry about a particular atom rather than the geometry of the molecule as a whole.

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Larger Molecules

This approach makes sense, especially because larger molecules tend to react at a particular site in the molecule.

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Polarity

• In Chapter 8 we discussed bond dipoles.

• But just because a molecule possesses polar bonds does not mean the molecule as a whole will be polar.

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Polarity

By adding the individual bond dipoles, one can determine the overall dipole moment for the molecule.

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Polarity

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Overlap and Bonding

• We think of covalent bonds forming through the sharing of electrons by adjacent atoms.

• In such an approach this can only occur when orbitals on the two atoms overlap.

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Overlap and Bonding

• Increased overlap brings the electrons and nuclei closer together while simultaneously decreasing electron-electron repulsion.

• However, if atoms get too close, the internuclear repulsion greatly raises the energy.

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Hybrid Orbitals

But it’s hard to imagine tetrahedral, trigonal bipyramidal, and other geometries arising from the atomic orbitals we recognize.

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Hybrid Orbitals

• Consider beryllium:– In its ground electronic

state, it would not be able to form bonds because it has no singly-occupied orbitals.

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Hybrid Orbitals

But if it absorbs the small amount of energy needed to promote an electron from the 2s to the 2p orbital, it can form two bonds.

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Hybrid Orbitals

• Mixing the s and p orbitals yields two degenerate orbitals that are hybrids of the two orbitals.– These sp hybrid orbitals have two lobes like a p orbital.– One of the lobes is larger and more rounded as is the s

orbital.

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Hybrid Orbitals

• These two degenerate orbitals would align themselves 180 from each other.

• This is consistent with the observed geometry of beryllium compounds: linear.

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Hybrid Orbitals

• With hybrid orbitals the orbital diagram for beryllium would look like this.

• The sp orbitals are higher in energy than the 1s orbital but lower than the 2p.

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Hybrid Orbitals

Using a similar model for boron leads to…

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Hybrid Orbitals

…three degenerate sp2 orbitals.

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Hybrid Orbitals

With carbon we get…

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Hybrid Orbitals

…four degenerate

sp3 orbitals.

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Hybrid Orbitals

For geometries involving expanded octets on the central atom, we must use d orbitals in our hybrids.

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Hybrid Orbitals

This leads to five degenerate sp3d orbitals…

…or six degenerate sp3d2 orbitals.

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Hybrid Orbitals

Once you know the electron-domain geometry, you know the hybridization state of the atom.

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Valence Bond Theory

• Hybridization is a major player in this approach to bonding.

• There are two ways orbitals can overlap to form bonds between atoms.

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Sigma () Bonds

• Sigma bonds are characterized by– Head-to-head overlap.– Cylindrical symmetry of electron density about the

internuclear axis.

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Pi () Bonds

• Pi bonds are characterized by– Side-to-side overlap.– Electron density

above and below the internuclear axis.

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Single Bonds

Single bonds are always bonds, because overlap is greater, resulting in a stronger bond and more energy lowering.

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Multiple Bonds

In a multiple bond one of the bonds is a bond and the rest are bonds.

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Multiple Bonds

• In a molecule like formaldehyde (shown at left) an sp2 orbital on carbon overlaps in fashion with the corresponding orbital on the oxygen.

• The unhybridized p orbitals overlap in fashion.

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Multiple Bonds

In triple bonds, as in acetylene, two sp orbitals form a bond between the carbons, and two pairs of p orbitals overlap in fashion to form the two bonds.

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Delocalized Electrons: Resonance

When writing Lewis structures for species like the nitrate ion, we draw resonance structures to more accurately reflect the structure of the molecule or ion.

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• In reality, each of the four atoms in the nitrate ion has a p orbital.

• The p orbitals on all three oxygens overlap with the p orbital on the central nitrogen.

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This means the electrons are not localized between the nitrogen and one of the oxygens, but rather are delocalized throughout the ion.

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Resonance

The organic molecule benzene has six bonds and a p orbital on each carbon atom.

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Resonance

• In reality the electrons in benzene are not localized, but delocalized.

• The even distribution of the electrons in benzene makes the molecule unusually stable.

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Molecular Orbital (MO) Theory

Though valence bond theory effectively conveys most observed properties of ions and molecules, there are some concepts better represented by molecular orbitals.

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• In MO theory, we invoke the wave nature of electrons.

• If waves interact constructively, the resulting orbital is lower in energy: a bonding molecular orbital.

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If waves interact destructively, the resulting orbital is higher in energy: an antibonding molecular orbital.

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MO Theory

• In H2 the two electrons go into the bonding molecular orbital.

• The bond order is one half the difference between the number of bonding and antibonding electrons.

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MO Theory

For hydrogen, with two electrons in the bonding MO and none in the antibonding MO, the bond order is

12

(2 - 0) = 1

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MO Theory

• In the case of He2, the bond order would be

12

(2 - 2) = 0

• Therefore, He2 does not exist.

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MO Theory

• For atoms with both s and p orbitals, there are two types of interactions:– The s and the p orbitals

that face each other overlap in fashion.

– The other two sets of p orbitals overlap in fashion.

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MO Theory

• The resulting MO diagram looks like this.

• There are both and bonding molecular orbitals and * and * antibonding molecular orbitals.

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MO Theory

• The smaller p-block elements in the second period have a sizeable interaction between the s and p orbitals.

• This flips the order of the and molecular orbitals in these elements.

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Second-Row MO Diagrams

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Chapter 10Gases



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Characteristics of Gases

• Unlike liquids and solids, gases– expand to fill their containers;– are highly compressible;– have extremely low densities.

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• Pressure is the amount of force applied to an area.

Pressure

• Atmospheric pressure is the weight of air per unit of area.

P =FA

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Units of Pressure

• Pascals– 1 Pa = 1 N/m2

• Bar– 1 bar = 105 Pa = 100 kPa

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Units of Pressure

• mm Hg or torr–These units are literally the difference in the heights measured in mm (h) of two connected columns of mercury.

• Atmosphere1.00 atm = 760 torr

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Standard Pressure

• Normal atmospheric pressure at sea level is referred to as standard pressure.

• It is equal to1.00 atm

760 torr (760 mm Hg)101.325 kPa

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Manometer

This device is used to measure the difference in pressure between atmospheric pressure and that of a gas in a vessel.

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Boyle’s Law

The volume of a fixed quantity of gas at constant temperature is inversely proportional to the pressure.

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As P and V areinversely proportional

A plot of V versus P results in a curve.

Since

V = k (1/P)This means a plot of V versus 1/P will be a straight line.

PV = k

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Charles’s Law

• The volume of a fixed amount of gas at constant pressure is directly proportional to its absolute temperature.

A plot of V versus T will be a straight line.

• i.e.,VT

= k

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Avogadro’s Law

• The volume of a gas at constant temperature and pressure is directly proportional to the number of moles of the gas.

• Mathematically, this means V = kn

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Ideal-Gas Equation

V 1/P (Boyle’s law)V T (Charles’s law)V n (Avogadro’s law)

• So far we’ve seen that

• Combining these, we get

V nTP

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Ideal-Gas Equation

The constant of proportionality is known as R, the gas constant.

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Ideal-Gas Equation

The relationship

then becomes

nTP

V

nTP

V = R

or

PV = nRT

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Densities of Gases

If we divide both sides of the ideal-gas equation by V and by RT, we get

nV

PRT

=

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• We know thatmoles molecular mass = mass

Densities of Gases

• So multiplying both sides by the molecular mass () gives

n = m

PRT

mV

=

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Densities of Gases

• Mass volume = density

• So,

Note: One only needs to know the molecular mass, the pressure, and the temperature to calculate the density of a gas.

PRT

mV

=d =

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Molecular Mass

We can manipulate the density equation to enable us to find the molecular mass of a gas:

Becomes

PRT

d =

dRTP =

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Dalton’s Law ofPartial Pressures

• The total pressure of a mixture of gases equals the sum of the pressures that each would exert if it were present alone.

• In other words,

Ptotal = P1 + P2 + P3 + …

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Partial Pressures

• When one collects a gas over water, there is water vapor mixed in with the gas.

• To find only the pressure of the desired gas, one must subtract the vapor pressure of water from the total pressure.

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Kinetic-Molecular Theory

This is a model that aids in our understanding of what happens to gas particles as environmental conditions change.

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Main Tenets of Kinetic-Molecular Theory

Gases consist of large numbers of molecules that are in continuous, random motion.

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The combined volume of all the molecules of the gas is negligible relative to the total volume in which the gas is contained.

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Attractive and repulsive forces between gas molecules are negligible.

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Energy can be transferred between molecules during collisions, but the average kinetic energy of the molecules does not change with time, as long as the temperature of the gas remains constant.

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The average kinetic energy of the molecules is proportional to the absolute temperature.

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Effusion

Effusion is the escape of gas molecules through a tiny hole into an evacuated space.

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Effusion

The difference in the rates of effusion for helium and nitrogen, for example, explains a helium balloon would deflate faster.

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Diffusion

Diffusion is the spread of one substance throughout a space or throughout a second substance.

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Graham's Law

KE1 KE2=

1/2 m1v12 1/2 m2v2

2=

=m1

m2

v22

v12

m1

m2

v22

v12

= v2

v1

=

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Real Gases

In the real world, the behavior of gases only conforms to the ideal-gas equation at relatively high temperature and low pressure.

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Real Gases

Even the same gas will show wildly different behavior under high pressure at different temperatures.

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Deviations from Ideal Behavior

The assumptions made in the kinetic-molecular model (negligible volume of gas molecules themselves, no attractive forces between gas molecules, etc.) break down at high pressure and/or low temperature.

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Corrections for Nonideal Behavior

• The ideal-gas equation can be adjusted to take these deviations from ideal behavior into account.

• The corrected ideal-gas equation is known as the van der Waals equation.

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The van der Waals Equation

) (V − nb) = nRTn2aV2(P +

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Chapter 11Intermolecular Forces,

Liquids, and Solids

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States of MatterThe fundamental difference between states of matter is the distance between particles.

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States of MatterBecause in the solid and liquid states particles are closer together, we refer to them as condensed phases.

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The States of Matter

• The state a substance is in at a particular temperature and pressure depends on two antagonistic entities:

– the kinetic energy of the particles;

– the strength of the attractions between the particles.

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Intermolecular Forces

The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together.

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They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.

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These intermolecular forces as a group are referred to as van der Waals forces.

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van der Waals Forces

• Dipole-dipole interactions

• Hydrogen bonding

• London dispersion forces

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Ion-Dipole Interactions

• Ion-dipole interactions (a fourth type of force), are important in solutions of ions.

• The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.

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Dipole-Dipole Interactions

• Molecules that have permanent dipoles are attracted to each other.– The positive end of one is

attracted to the negative end of the other and vice-versa.

– These forces are only important when the molecules are close to each other.

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Dipole-Dipole Interactions

The more polar the molecule, the higher is its boiling point.

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London Dispersion Forces

While the electrons in the 1s orbital of helium would repel each other (and, therefore, tend to stay far away from each other), it does happen that they occasionally wind up on the same side of the atom.

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At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side.

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Another helium nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1.

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London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole.

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• These forces are present in all molecules, whether they are polar or nonpolar.

• The tendency of an electron cloud to distort in this way is called polarizability.

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Factors Affecting London Forces

• The shape of the molecule affects the strength of dispersion forces: long, skinny molecules (like n-pentane tend to have stronger dispersion forces than short, fat ones (like neopentane).

• This is due to the increased surface area in n-pentane.

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Factors Affecting London Forces

• The strength of dispersion forces tends to increase with increased molecular weight.

• Larger atoms have larger electron clouds which are easier to polarize.

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Which Have a Greater Effect?Dipole-Dipole Interactions or Dispersion Forces

• If two molecules are of comparable size and shape, dipole-dipole interactions will likely the dominating force.

• If one molecule is much larger than another, dispersion forces will likely determine its physical properties.

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How Do We Explain This?

• The nonpolar series (SnH4 to CH4) follow the expected trend.

• The polar series follows the trend from H2Te through H2S, but water is quite an anomaly.

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Hydrogen Bonding

• The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong.

• We call these interactions hydrogen bonds.

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Hydrogen Bonding

• Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine.

Also, when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.

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Summarizing Intermolecular Forces

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Intermolecular Forces Affect Many Physical Properties

The strength of the attractions between particles can greatly affect the properties of a substance or solution.

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Viscosity• Resistance of a liquid

to flow is called viscosity.

• It is related to the ease with which molecules can move past each other.

• Viscosity increases with stronger intermolecular forces and decreases with higher temperature.

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Surface Tension

Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.

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Solids

• We can think of solids as falling into two groups:– crystalline, in which

particles are in highly ordered arrangement.

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Solids

• We can think of solids as falling into two groups:– amorphous, in which

there is no particular order in the arrangement of particles.

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Crystalline SolidsBecause of the ordered in a crystal, we can focus on the repeating pattern of arrangement called the unit cell.

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Crystalline Solids

There are several types of basic arrangements in crystals, like the ones depicted above.

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Crystalline Solids

We can determine the empirical formula of an ionic solid by determining how many ions of each element fall within the unit cell.

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Ionic Solids

• What are the empirical formulas for these compounds?– (a) Green: chlorine; Gray: cesium– (b) Yellow: sulfur; Gray: zinc– (c) Gray: calcium; Blue: fluorine

CsCl ZnS CaF2

(a) (b) (c)

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Attractions in Ionic CrystalsIn ionic crystals, ions pack themselves so as to maximize the attractions and minimize repulsions between the ions.

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Types of Bonding in Crystalline Solids

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Covalent-Network andMolecular Solids

• Diamonds are an example of a covalent-network solid, in which atoms are covalently bonded to each other.– They tend to be hard and have high melting

points.

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Covalent-Network andMolecular Solids

• Graphite is an example of a molecular solid, in which atoms are held together with van der Waals forces.– They tend to be softer and have lower melting

points.

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Metallic Solids

• Metals are not covalently bonded, but the attractions between atoms are too strong to be van der Waals forces.

• In metals valence electrons are delocalized throughout the solid.

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Phase Changes

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Energy Changes Associated with Changes of State

The heat of fusion is the energy required to change a solid at its melting point to a liquid.

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The heat of vaporization is defined as the energy required to change a liquid at its boiling point to a gas.

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• The heat added to the system at the melting and boiling points goes into pulling the molecules farther apart from each other.

• The temperature of the substance does not rise during a phase change.

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Vapor Pressure

• At any temperature some molecules in a liquid have enough energy to escape.

• As the temperature rises, the fraction of molecules that have enough energy to escape increases.

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Vapor Pressure

As more molecules escape the liquid, the pressure they exert increases.

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Vapor Pressure

The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condense at the same rate.

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Vapor Pressure

• The boiling point of a liquid is the temperature at which it’s vapor pressure equals atmospheric pressure.

• The normal boiling point is the temperature at which its vapor pressure is 760 torr.

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Phase Diagrams

Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.

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Phase Diagrams

• The circled line is the liquid-vapor interface.• It starts at the triple point (T), the point at

which all three states are in equilibrium.

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Phase Diagrams

It ends at the critical point (C); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.

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Phase Diagrams

Each point along this line is the boiling point of the substance at that pressure.

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Phase Diagrams

• The circled line in the diagram below is the interface between liquid and solid.

• The melting point at each pressure can be found along this line.

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Phase Diagrams• Below the triple point the substance cannot

exist in the liquid state.• Along the circled line the solid and gas

phases are in equilibrium; the sublimation point at each pressure is along this line.

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Phase Diagram of Water

• Note the high critical temperature and critical pressure.– These are due to the

strong van der Waals forces between water molecules.

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Phase Diagram of Water

• The slope of the solid-liquid line is negative.– This means that as the

pressure is increased at a temperature just below the melting point, water goes from a solid to a liquid.

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Phase Diagram of Carbon Dioxide

Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO2 sublimes at normal pressures.

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Phase Diagram of Carbon Dioxide

The low critical temperature and critical pressure for CO2 make supercritical CO2 a good solvent for extracting nonpolar substances (like caffeine)

chpt1 11

Education

matter atoms

building blocks of matter

proximity ofa measurement

science westudy matter

burmamatter andmeasurement

mlmatter andmeasurement

chemical properties

different degrees of