general chemistry chem. 101 for - hamohammed - pbworkshamohammed.pbworks.com/f/10112.pdf ·...

Prof.H.A.Mohammed General Chemistry – Chem 101

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Taibah University

Faculty of Science and Arts

Chemistry Department

For

Chemistry , Physics and Biology Students

General Chemistry Chem. 101


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General Chemistry (101 Chem.)

Course Description

Syllabus : The course will cover the following topics

Uيحتوي المقررعلى الموضوعات التاليةU :

Chemical Foundations-1 الكيميائية االسس •

Atoms, Molecules and Ions-2الذرات والجزيئات وااليونات •

Stoichiometry-3 الحسابات الكيميائية •

Types of Chemical Reactions-4 انواع التفاعالت الكيميائية

Atomic Structure and-5 الدوري البناء الذري والتصنيف • Periodicity

Chemical Bonding-6 الترابط الكيميائي •


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Chapter 1: Chemical Foundations

Chemistry: An overview

1.1 The Scientific Method

1.2 Units of Measurement

1.3 Uncertainty in Measurement

1.4 Significant Figures and Calculations

1.5 Dimensional Analysis 1.6 Temperature

1.7 Classification of Matter


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Chemical Foundations

Chemistry: An overview مقدمة:الكيمياء

The Scientific Method 1.1 الطريقة العلمية

Units of Measurement 1.2 وحدات القياس

Uncertainty in Measurement 1.3 الشك في القياس

UChemistry: An overview:

• An Observational Science

• An Experimental Science

• A Laboratory Science

• An Interesting Science

• An Important Science

• A “Hard” Science

UWhat Is Chemistry?

The Study of Matter and its Properties, the Changes that Matter Undergoes, and the Energy Associated with those Changes.

For example: Oxygen atom, hydrogen atom, water molecule


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Water turning to oxygen and hydrogen

Everything is made of tiny particles called atoms and molecules.

Chemists study these particles, looking at the kinds, numbers, structure, size which produce varying chemical and physical properties


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Humans are by nature curious.

Science is just exploring nature.

A scientists is just a person exploring.

You begin to organize your thoughts into Observation, you group those observations into Hypotheses, using Experimentation, and formulate Laws or Theories


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1.1 Steps in the Scientific Method

1--Observations

• quantitative : water boils at 100C • qualitative: the sky is blue, water is liquid

2.Formulating hypotheses

• possible explanation for the observation

3.Performing experiments : gathering new information to decide whether the hypothesis is valid Theory (Model)

• A set of tested hypotheses that give an overall explanation of some natural phenomenon.

Natural Law

The same observation applies to many different systems Example - Law of Mass Conservation “In a chemical reaction matter is either created nor destroyed.”

Speed of Light, E = mc2

Law vs. Theory

A law summarizes what happens;

A theory (model) is an attempt to explain why it happens.

Laws : Very specific, “What will happen” often expressed in mathematical equations.


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Theories: Very general, “Why it will happen,” often includes many “Laws

1.2 Units of Measurement

The Fundamental SI Units

Physical Quantity

Name of Unit Abbreviation

Mass kilogram kg

Length meter m

Time second s

Temperature kelvin K

Electric current ampere A

Amount of substance mole mol

Luminous intensity candela cd

Prefixes Used to Modify Standard Unit

1. kilo = 1000 times base unit = 103

k=1000 or k = 103 1 kg = 103g

• nano = 10-9 times the base unit


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n=.000000001 or n = 10-9 1 nL = 10-9L

Tro's "Introductory Chemistry", Chapter 2

۲۰

•Apply the solution map:

InformationGiven:7.8 km

Find:? miConversion Factor:1 km = 0.6214 mile

Solution Map:km → mi

Example:Convert 7.8 km to miles.

km 1mi .62140

mi km 1

mile 0.6214km 7.8 =×

= 4.84692 mi

= 4.8 mi •Significant figures and round:

2 significant figures

2 significant figures

A digit that must be estimated is called uncertain. A measurement always has some degree of uncertainty Precision and Accuracy

Accuracy refers to the agreement of

a particular value with the true value.

القراءة الصحيحةمدى تطابق القراءة المقاسة مع Precision refers to the degree of agreement among several elements of the same quantity.


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مدى تطابق القراءات المقاسة مع بعضها البعض

بعض االدوات المعملية المستخدمة لقياس الحجم

Prepared by Dr. Amjad Shraim and Dr. Fethi Kooli۲۲


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۲۳

The difference between accuracyوالصحةالفرق بين الدقة and precision

الدقة عدم الصحة و عدم الدقة الصحة و الدقة

1.4 Significant Figures and Calculations

Rules for Counting significant figures

1-Non-Zero integers

2-Zeros:

(a)Leading zeros (b) Capitive zero (c) trailing zeros

3. Exact numbers

1-Nonzero integers always count as significant figures


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2-Zeros

(a)Leading zeros do not count as significant figures.

(b)Captive zeros always count as significant figures

16.07 has 4 sig.figs

(c) Trailing zeros are significant only if t the number contains

a decimal point.

9.300 has 4 sig figs

3-Exact numbers have an infinite number of significant figures

1 inch = 2.54 cm, exactly

Multiplication and Division

# sig figs in the result equals the number in the least precise measurement used in the calculation.

6.38 x 2.0 =12.76

13 (2 sig figs)

Addition and Subtraction:


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# sig figs in the result equals the number of decimal places in

the least precise measurement.

6.8 + 11.934 = 22.4896 ~ =22.5 (3 sig. figs)

Numbers

1-Exact

2-Measured: Significant Figures

• Exact: Sometimes you can determine an exact value for a quality

of an object.

• Often by counting.


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• Books on a table.

• Sometimes by definition

• 1 cm is exactly 1/100th of 1 meter.

• Measured: Whenever you use an instrument to compare a quality

of an object to a standard, there is uncertainty in the comparison.

Writing a Number in Scientific Notation, Continued

0.00012340

1. Locate the decimal point.

0.00012340

2. Move the decimal point to obtain a number between 1 and 10.

1.2340

3. Multiply the new number by 10n .

• Where n is the number of places you moved

the decimal point.

1.2340 x 104

4. If you moved the decimal point to the left, then n is +;

if you moved it to the right, then n is − .

1.2340 x 10-4


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Writing a Number in Standard Form: ( 1.234 x 10-6 )

• Since exponent is -6, make the number smaller by moving the

decimal point to the left 6 places.

• When you run out of digits to move around, add zeros.

• Add a zero in front of the decimal point for decimal numbers.

000 001.234

0.000 001 234

How many sig figs?

45.87360.000239 0.00023900 48000. 48000 3.982×106

1.00040

6

3

5

5

2

4

6

•All digits count

•Leading 0’s don’t

•Trailing 0’s do

•0’s count in decimal form

•0’s don’t count w/o decimal

•All digits count

•0’s between digits count as well as trailing in decimal form

Rounding


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When rounding to the correct number of significant figures, if the number after the place of the last significant figure is:

0 to 4, round down.

Drop all digits after the last significant figure and leave the last significant figure alone.

5 to 9, round up.

Drop all digits after the last significant figure and increase the last

significant figure by one

1.5 Dimensional Analysis

Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Proper use of “unit factors” leads to proper units in your answer.

OK mile: .1 0 621371

kilometer0.62137 mile kilometer

=

NOT OK: 1 kilometer0.62137 mile

1 mile0.62137 kilometer

=

التحويل بين الوحدات المختلفة


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Temperature Scales and Interconversions

Kelvin ( K ) - The “Absolute temperature scale” begins at absolute zero and only has positive values.

Celsius ( oC ) - The temperature scale used by science, formally called centigrade and most commonly used scale around the world.Water freezes at 0oC, and boils at 100oC.

Fahrenheit ( oF ) - Commonly used scale in America for our weather reports. Water freezes at 32oF,and boils at 212oF.

T (in K) = T (in oC) + 273.15

T (in oC) = T (in K) - 273.15

T (in oF) = 9/5 T (in oC) + 32

T (in oC) = [ T (in oF) - 32 ] 5/9

Problem 3-8:Temperature Conversions

(a) The boiling point of Liquid Nitrogen is -195.8 oC, what is the temperature in Kelvin and degrees Fahrenheit?

T (in K) = T (in oC) + 273.15T (in K) =

T (in oF) = 9/5 T (in oC) + 32T (in oF) =

(b)The normal body temperature is 98.6oF, what is it in Kelvinand degrees Celsius?

T (in oC) = [ T (in oF) - 32] 5/9T (in oC) =

T (in K) = T (in oC) + 273.15T (in K) =


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٥٦

1.6 Density• Inverse relationship between mass and volume.• Solids = g/cm3

– 1 cm3 = 1 mL• Liquids = g/mL• Gases = g/L• Volume of a solid can be determined by water

displacement—Archimedes Principle.• Density : solids > liquids > gases

– Except ice is less dense than liquid water!

VolumeMassDensity =

1.7 Clssification of Matter

Any matter can exist in one of 3 States

• (a) Gas (b) Liquid (c) Solid


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Solids• The particles in a solid are packed

close together and are fixed in position.– Although they may vibrate.

• The close packing of the particles results in solids being incompressible.

• The inability of the particles to move around results in solids retaining their shape and volume when placed in a new container and prevents the particles from flowing.

٦۳

Liquids

• The particles in a liquid are closely packed, but they have some ability to move around.

• The close packing results in liquids being incompressible.

• The ability of the particles to move allows liquids to take the shape of their container and to flow. However, they don’t have enough freedom to escape and expand to fill the container.


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Matter: is it pure or impure

• Pure Substance = All samples are made of the same pieces in the same percentages.– Salt

• Mixtures = Different samples may have the same pieces in different percentages.– Salt water

Pure SubstanceConstant Composition

Homogeneous

MixtureVariable Composition

Matter

Heterogeneous


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Mixtures

1. Made of multiple substances, but appears to be one substance.

2. All portions of a sample have the same composition and properties.

1. Made of multiple substances, whose presence can be seen.

2. Portions of a sample have different composition and properties.

Heterogeneous Homogeneous

٦۷

Matter Summary


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• Matter has PropertiesPhysical Properties are the characteristics of matter that can be changed without changing its composition.

A. Characteristics that are directly observable.

• Chemical Properties are the characteristics that determine how the composition of matter changes as a result of contact with other matter or the influence of energy.

A. Characteristics that describe the behavior of matter.

Matter has Properties, Matter can also go through Changes

• Changes that alter the state or appearance of the matter without altering the composition are called physical changes.

• Changes that alter the composition of the matter are called chemical changes.

A. During the chemical change, the atoms that are present rearrange into new molecules, but all of the original atoms are still present.

Is it a Physical or Chemical Change?

– A physical change results in a different form of the same substance.


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– The kinds of molecules don’t change.

– A chemical change results in one or more completely new substances.

– Also called chemical reactions.

– The new substances have different molecules than the original substances.

You will observe different physical properties because the new substances have their own physical properties


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1-Fundamental Chemical Laws 1-Conservation of Mass

Definite Composition-2Law of Multiple Proportions-3

2-Early Experiments To Characterize Atom 3-The Periodic Table4-Naming Simple Compounds

Chapter :2

Atoms, Molecules and Ions


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Three Laws that Led to the Atomic Theory• Law of Mass Conservation: The total mass of substances does not

change during a chemical reaction (Lavoisier).

• Law of Definite ( or Constant ) Composition: No matter what its source, a particular chemical compound is composed of the same elements in the same parts (fractions) by mass (Proust).

• The Law of Multiple Proportions: When two elements form a series of compounds, the masses of one element that combine with a fixed mass of the other element are in the ratio of small integers to each other (Dalton).

1-Fundamental Chemical Laws : Lecture :1

1-Conservation of Mass

• Matter is neither created nor destroyed in a chemical reaction.

• In every chemical operation an equal quantity of matter exists before and after the operation.

• Mass is conserved in a chemical reaction.

• Moreover, in chemical change, the mass of the elements is conserved, element by element.

• Development of this law was made possible by the analytical balance.


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Problem

• Potassium chlorate (KClO3) decomposes to potassium chloride (KCl) and oxygen (O2) when heated. In one experiment 100.0 g of KClO3 generated 36.9g of O2 and 57.3 g of KCl. What mass of KClO3 remained unreacted?

• Mass of KClO3 before reaction = mass of KCl + mass of O2 + mass of unreacted KClO3

• 100.0 g of KClO3 = 57.3 g KCl + 36.9g O2 + g unreacted KClO3

• g unreacted KClO3 = 100.0 g - 57.3 g - 36.9 g = 5.8 g

2-Definite Composition

Chemical analysis of a 9.07 g sample of calcium phosphate shows that it contains 3.52 g of Ca.

How much Ca could be obtained from a 1.000 kg

sample?

Mass fraction Ca = 3.52 g Ca / (9.07 g total) = 0.388

(i.e., 38.8% Ca by mass in any sample of compound)

Mass Ca in 1.000 kg =

(1.000 kg total)x(0.388 g Ca/ g total) = 0.388 kg C = 388 g Ca

2-Law of Definite Proportions


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In a given chemical compound, the proportions by mass of the elements that compose it are fixed, regardless of the source of the compound.

The ratio of elements in a compound is fixed regardless of the source of the compound.

Water is made up of 11.1% by mass of hydrogen and 88.9% oxygen.

3-Law of Multiple Proportions

If elements A and B react to form two compounds,the different masses of B that combine with a fixed mass of A can be expressed as a ratio of small whole numbers.

Example: Nitrogen Oxides I & II

Cmpd Mass % N

Mass % O

Mass Ratio O/N

Ratio of Ratios

I 46.68 53.32 1.142 1

II 30.45 69.55 2.284 2


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Law of Multiple Proportions (Dalton)• If elements A & B react to form more than one compound, the different

masses of “B” that combine with a fixed mass of “A” can be expressed as a ratio of SMALL WHOLE NUMBERS.

Ex. Assume two compounds containing just Carbon (C) and Oxygen (O) with the following relative compositions

Carbon Oxide (I): 57.1 % Oxygen and 42.9 % CarbonCarbon Oxide (II): 72.7 % Oxygen and 27.3 % Carbon

Mass Ratios: Oxide (I) = 57.1 / 42.9 C = 1.33 g O / g C Oxide (II) = 72.7 / 27.3 C = 2.66 g O / g C

Ratio Oxide (I) / Oxide (II) = 1.33 / 2.66 = 1/2Expressing the relative mass fractions as “small whole numbers”, the ratio of oxygen atoms to carbon atoms in Oxide I is 1:1 (CO)The ratio of oxygen atoms to carbon atoms in Oxide II is 2:1 (CO2)

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Dalton’s Atomic Theory

• Matter is composed of (indivisible) atoms.

• All atoms of a given chemical element are identical in mass and in all other properties.

• Different chemical elements are composed of different atoms of different masses.

• Atoms are indestructible. They retain their identities in a chemical reaction.

• A compound forms from its elements through the combination of atoms of unlike elements in small whole number ratios.



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Early Experiments To Characterize The Atom

Experiencing Atoms

1. There are about 91 elements found in nature.

– Over 20 have been made in laboratories.

(a) Each kind of atom is unique


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a. Carbon is not Hydrogen

b. They have different properties

i. Structure, magnetic meaning they can attract and repel other atoms, melting, boiling, electrical, stability, reactivity (attract and repel), etc…

The Atom Is Divisible

• Work done by J. J. Thomson and others proved that the atom had pieces(subatomic particles) called electrons.

• Thomson found that electrons are much smaller than atoms and carry a negative charge.

– The mass of the electron is 1/1836th the mass of a hydrogen atom.

– The charge on the electron is the fundamental unit of charge that we call –1 charge unit.

The Electron

1- A Glass tube from which most of the air has been evacuated.

2-When two metal plates are connected to a high –voltage source, the negatively charged plate, called the cathode, emits an invisible ray.

3- The cathode ray is drawn to the positively charged plate, called the anode, where it passes through a hole and continues travelling to the other of the tube

4-When the ray strikes the specially coated surface, it produces a strong fluorescence, or bright light


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5-The cathode ray is attracted by the plate bearing +ve charges and repelled by the plate bearing –ve charges so it must consists of negatively charged particles called electrons

A cathode-ray tube: the fast-moving electrons excite the gas in the tube, causing a glow between the electrodes(see fig.)

Millikan’s Experiment :Mass of Electron 1. Measured rate of droplet’s fall without voltage: obtained droplet’s mass. 2. Voltage across plates influenced speed, due to charge on droplet. 3. Quantitative effect of voltage w/ laws of physics -> amt. of charge on droplet. 4. RESULT: Different droplets had different charge, but always a multiple of same number -> elementary charge on electron: e = 1.602x10-19 coulombs (negative).


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5- e/m (charge/mass)= - 1.76x108 c/g (Thomson experiments) 5. (Mass/charge) x e- = mass of e-

Rutherford's Experiment On α -Particle Bombardment

of Metal foil :


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Rutherford’s Interpretation— The Nuclear Model

The atom contains a tiny dense center called the nucleus.

– The amount of space taken by the nucleus is only about 1/10 trillionth the volume of the atom.


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2. The nucleus has essentially the entire mass of the atom.

– The electrons weigh so little they contribute practically no mass to the atom.

3. The nucleus is positively charged.

• The amount of positive charge balances the negative charge of the electrons.

4.The electrons are dispersed in the empty space of

the atom surrounding the nucleus.

• Like water droplets in a cloud.

Some Problems:

• How could beryllium have 4 protons stuck together in the nucleus?

• Shouldn’t they repel each other?

• If a beryllium atom has 4 protons, then it should weigh 4 amu, but it actually weighs 9.01 amu! Where is the extra mass coming from?

• Each proton weighs 1 amu.

• Remember: The electron’s mass is only about 0.00055 amu and Be has only 4 electrons—it can’t account for the extra 5 amu of mass.

• There Must Be Something Else

ThereTo answer these questions, Rutherford proposed that there was another particle in the nucleus—it is called a neutron.


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• Neutrons (are discovered by Schadwick upon bombarding a thin sheet of Be with α –particles : high energy radiation was emitted: neutrons)

have no charge and a mass of 1 amu.

– The masses of the proton and neutron are both

– approximately 1 amu.

Mass and Charge of subatomic particles

• The Modern Atom we know atoms are composed of three main pieces—protons, neutrons, and electrons.

• The nucleus contains protons and neutrons.

• The nucleus is only about 10 P

-13P cm in diameter.

• The electrons move outside the nucleus with an average distance of about 10P

-8P cm.

– Therefore, the radius of the atom is about 10P

5P times larger

than the radius of the nucleus.

Some Notes on Charges

Subatomic particle

Mass g

Mass amu

Location in atom

Charge Symbol

Proton

1.67262 x 10-24

1.0073 nucleus 1+ p, p+, H+

Electron 0.00091 x 10-24

0.00055 empty space 1− e, e-

Neutron 1.67493 x 10-24

1.0087 nucleus 0 n, n0


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– There are two kinds of charges, called positive and negative.

– Opposite charges attract : + attracted to –.

Like charges repel : (+) repels( +) (–) repels( –).

To be neutral, something

must have no charge or equal

amounts of opposite charges

– Elements:

– Each element has a unique number of protons in its nucleus.

– All carbon atoms have 6 protons in their nuclei.

– The number of protons in the nucleus of an atom is called the atomic number.

– Z is the short-hand designation for the atomic number.

– Because each element’s atoms have a unique number of protons, each element can be identified by its atomic number.

– The elements are arranged on the Periodic Table in order of their atomic numbers.

– Each element has a unique name and symbol.


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– The symbol is either one or two letters

One capital letter or one capital letter + one lower case letter

Atomic Definitions

XA

Z

X = Atomic symbol of the element, or element symbol

A = The Mass number; A = Z + N

Z = The Atomic Number, the Number of Protons in the Nucleus

(All atoms of the same element havethe same no. of protons.)

N = The Number of Neutrons in the Nucleus

The Nuclear Symbol of the Atom, or Isotope


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Periodic law : Mendeleev

– Ordered elements by atomic mass.

– Saw a repeating pattern of properties.

– Periodic law —When the elements are arranged in order of increasing relative mass, certain sets of properties recur periodically?

– Used pattern to predict properties of undiscovered elements.

– Where atomic mass order did not fit other properties, he reordered by other properties.

Te & I

Periodicity

٥۰

= Metal

= Metalloid

= Nonmetal


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The Modern Periodic Table

– Elements with similar chemical and physical properties are in the same column.

– Columns are called Groups or Families.

– Designated by a number and letter at top.

– Rows are called Periods.

– Each period shows the pattern of properties repeated in the next period.

– Main group = representative elements = “A” groups.

– Transition elements = “B” groups.

– All metals.

– Bottom rows = inner transition elements = rare earth elements.

– Metals

– Really belong in periods 6 and 7.


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٥۳

= Alkali metals

= Alkali earth metals

= Noble gases

= Halogens

= Lanthanides

= Actinides

= Transition metals

Important Groups—Hydrogen

• Nonmetal.

• Colorless, diatomic gas.

– Very low melting point and density.

• Reacts with nonmetals to form molecular compounds.

– HCl is an acidic gas.

– H2O is a liquid.

• Reacts with metals to form hydrides.

– Metal hydrides react with water to form H2.

hydrogen halides dissolve in water to form acids


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Atom Happiness:

• The number of electrons determines the physical and chemical properties of an atom

• When the protons=electrons atoms are electrically neutral

• But Atom Happiness only comes about when the atom has the same number of Electrons as a noble gas

• Atoms will not gain and loose protons because protons are at the center of atoms, very far away from outside, in an electrical shroud of negative charge

• Atoms will gain and loose electrons, which are on the outside surface of atoms

• When an atom gains or looses an electron the electrical balance is lost

• But, atoms are happier with a charge

Ions:

– Ions with a positive charge are called cations.

– More protons than electrons.

– Form by losing electrons.

– Ions with a negative charge are called anions.

– More electrons than protons.

– Form by gaining electrons.

– Chemically, ions are much different than the neutral atoms.


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– Because they have a different structure.

Atomic Structures of Ions

• Nonmetals form anions.

• For each negative charge, the ion has 1 more electron than the neutral atom.

o F = 9 p+ and 9 e− ; F─ = 9 p+ and 10 e−.

o P = 15 p+ and 15 e− ; P3─ = 15 p+ and 18 e−.

• Anions are named by changing the ending of the name to –ide.

fluorine F + 1e− → F─ fluoride ion

oxygen O + 2e− → O2─ oxide ion

• The charge on an anion can often be determined from the group number on the periodic table.

– Group 7A ⇒ 1−, Group 6A ⇒ 2−.

• Metals form cations.

• For each positive charge the ion has 1 less electron than the neutral atom.

• Na atom = 11 p+ and 11 e−; Na+ ion = 11 p+ and 10 e−.

• Ca atom = 20 p+ and 20 e−; Ca2+ ion = 20 p+ and 18 e−.

• Cations are named the same as the metal.

sodium Na → Na+ + 1e− sodium ion


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calcium Ca → Ca2+ + 2e− calcium ion

• The charge on a cation can often be determined from the group number on the periodic table.

Group 1A ⇒ 1+, Group 2A ⇒ 2+, (Al, Ga, In) ⇒ 3+.

Valence Electrons and Ion Charge

• The highest energy electrons in an atom are called the valence electrons.

• Metals form cations by losing their valence electrons to get the same number of electrons as the previous noble gas.

• Main group metals.

• Li+ = 2 e− = He; Al3+ = 10 e− = Ne.

• Nonmetals form anions by gaining electrons to have the same number of electrons as the next noble gas.

• Cl− = 18 e− = Ar; Se2− = 36 e− = Kr.

Ion Charge and the Periodic Table

• The charge on an ion can often be determined from an elements position on the periodic table.

• Metals are always positive ions, nonmetals are negative ions.

• For many main group metals, the cation charge = the group number.

• For nonmetals, the anion charge = the group number – 8.


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٦۲

Li+

Na+

K+

Rb+

Cs+

Be2+

Mg2+

Ca2+

Sr2+

Ba2+

Al3+

Ga3+

In3+

O2−

S2−

Se2−

Te2−

F−

Cl−

Br−

I−

N3−

P3−

As3−

1A

2A 3A 7A6A5A

Structure of the Nucleus

• Soddy discovered that the same element could have atoms with different masses, which he called isotopes.

– There are two isotopes of chlorine found in nature, one that has a mass of about 35 amu and another that weighs about 37 amu.

• The observed mass is a weighted average of the weights of all the naturally occurring atoms.

– The atomic mass of chlorine is 35.45 amu.


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Isotopes

• All isotopes of an element are chemically identical.

– Undergo the exact same chemical reactions.

• All isotopes of an element have the same number of protons.

• Isotopes of an element have different masses.

• Isotopes of an element have different numbers of neutrons.

• Isotopes are identified by their mass numbers.

– Protons + neutrons.

Neon:

Order of Elements in a Formula

• Metals are written first.

• NaCl

• Nonmetals are written in order from Table 5.1.

• CO2


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• There are occasional exceptions for historical or informational reasons.

• H2O, but NaOH .

Ionic Compounds

• Metals + nonmetals.

• No individual molecule units, instead have a 3-dimensional array of cations and anions made of formula units. Ionic compounds are made of ions called cations and anions.

• Cations = + charged ions; anions = − charged ions.

• The sum of the + charges of the cations must equal the sum of the − charges of the anions.

• If Na+ is combined with S2-, you will need 2 Na+ ions for every S2- ion to balance the charges, therefore the formula must be Na2S.

•


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Naming Binary Compounds (Type I; Ionic)

• The cation is always named first and the anion second.

• A monatomic cation takes its name from the name of the element, e.g. Na+ is called sodium in the names of compounds containing this ion.

• A monatomic anion is named by taking the first part of the element and adding –ide, e.g. Cl- is chloride.

Naming Binary Compounds (Type I; Ionic)


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Naming Binary Compounds(Type II; Ionic)

• Applies to cations that can take on alternate charge states

• Using the principle of charge balance determine the cation charge.

• Include in the cation name a Roman numeral indicating the charge.


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Type II

Ionic Compounds with Polyatomic Ions

• Polyatomic ions are assigned special names that must be memorized.

• Special rules apply to anions that contain an atom of a given element and different numbers of oxygen atoms. These anions are called oxyanions.


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Tro's "Introductory Chemistry", Chapter 4 . Start learning these boldface ones


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Rules for Families of OxoanionsFamilies with Two Oxoanions

The ion with more O atoms takes the nonmetal root and thesuffix “-ate”.

The ion with fewer O atoms takes the nonmetal root and thesuffix “-ite”.

Families with Four Oxoanions (usually a Halogen)

The ion with most O atoms has the prefix “per-”, the nonmetalroot and the suffix “-ate”.

The ion with one less O atom has just the suffix “-ate”.

The ion with two less O atoms has the just the suffix “-ite”.

The ion with three less O atoms has the prefix “hypo-” and thesuffix “-ite”.

Binary Compounds (Type III; Covalent – Contain Two Nonmetals:

• The first element in the formula is named first, using the full element name.

• The second element is named as if it were an anion.

• Prefixes are used to denote the numbers of atoms present.

• The prefix mono- is never used for naming the first element, e.g. CO is carbon monoxide.


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NAMING OXOANIONS - EXAMPLES

Prefixes Root Suffixes Chlorine Bromine Iodine

per “ ” ate perchlorate perbromate periodate[ ClO4

-] [ BrO4-] [ IO4

-]

“ ” ate chlorate bromate iodate[ ClO3

-] [BrO3-] [ IO3

-]

“ ” ite chlorite bromite iodite[ ClO2

-] [ BrO2-] [ IO2

-]

hypo “ ” ite hypochlorite hypobromite hypoiodite[ ClO -] [ BrO -] [ IO -]N

o. o

f O a

tom

s

Binary Compounds (Type III; Covalent – Contain Two Nonmetals

• The first element in the formula is named first, using the full element name.

• The second element is named as if it were an anion.

• Prefixes are used to denote the numbers of atoms present.

The prefix mono- is never used for naming the first element, e.g. CO is carbon monoxide


55



Binary Molecular Compounds of Two Nonmetals

1. Name first element in formula first. – Use the full name of the element.

2. Name the second element in the formula with an −ide, as if it were an anion.

– However, remember these compounds do not contain ions!

3. Use a prefix in front of each name to indicate the number of atoms.

– Never use the prefix mono- on the first element.


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Subscript—Prefixes

• 1 = mono-

– Not used on first nonmetal.

• 2 = di-

• 3 = tri-

• 4 = tetra-

• 5 = penta-

• 6 = hexa-

• 7 = hepta-

• 8 = octa-

Drop last “a” if name begins with vowel

Example—Naming Binary Molecular, BFR3 R, Continued

4. Name the first element.

boron.

4. Name the second element with an –ide.

Fluorine ⇒ fluoride.

4. Add a prefix to each name to indicate the subscript.

monoboron, trifluoride.

4. Write the first element with prefix, then the second element with prefix.


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– Drop prefix mono- from first element.

boron trifluoride.

Formula-to-Name Acids

• Acids are molecular compounds that often behave like they are made of ions.

• All names have acid at end.

• Binary Acids = Hydro- prefix + stem of the name of the nonmetal + -ic suffix.

• Oxyacids:

– If polyatomic ion ends in –ate = Name of polyatomic ion with –ic suffix.

– If polyatomic ion ends in –ite = Name of polyatomic ion with –ous suffix.

Naming Oxyacids:

• If polyatomic ion name ends in –ate, then change ending to –ic suffix.

• If polyatomic ion name ends in –ite, then change ending to –ous suffix.

Write word acid at end of all names

Example—Naming Oxyacids, HR2RSOR4R,Continued

4. Identify the anion.


58


SO4 = SO42- = sulfate.

4. If the anion has –ate suffix, change it to –ic. If the anion has –ite suffix, change it to –ous.

SO42- = sulfate ⇒ sulfuric.

4. Write the name of the anion followed by the word acid.

sulfuric acid

(This is kind of an exception, to make it sound nicer!)

Writing Formulas for Acids

• When name ends in acid, formulas starts with H.

• Write formulas as if ionic, even though it is molecular.

• Hydro- prefix means it is binary acid, no prefix means it is an oxyacid.

• For an oxyacid, if ending is –ic, polyatomic ion ends in –ate; if ending is –ous, polyatomic ion ends in –ous.

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