2015 chemhandouts1

Partido State UniversitySE Enhancement Class 2015

Chemistry-handouts

SI or Metric system powers/multiples of 10 2 types of SI units

fundamental derived

7 base units/fundamental units meter – length. width or height/ distance Kelvin – temperature second – time kilogram – mass Ampere – electric current mole – amount of substance candela – luminous of intensity

precision vs. accuracy

precision values of set or series of measurements are closer to each other

accuracy measure of how close your measured value is to the actual value

Percent error (% error)

% error=measured valueactual value

x 100

Mass vs. Weight Mass - measure of the amount of matter in an object (constant). Weight - measure of the force of gravitational attraction between the object and the Earth

(changing)

Significant figures1. Nonzero digits are ALWAYS significant2. Leading zeros are NOT significant3. Captive zeros are ALWAYS significant4. Trailing zeros at the right end of a number are NOT significant; but are ALWAYS significant after

a decimal point

Significant figures in Measurements Multiplication/Division

number of significant figures in the answer—is the same as the number with LEAST significant figures

Addition/Subtraction answer will have the same number of places as the number with the LEAST places to the

right of the decimal point

Scientific Notation When multiplying exponential terms, add exponents. When dividing exponential terms, subtract exponents. When raising exponential terms to a power, multiply exponents.

Molecules: The Law of Definite Proportions Law of definite proportions - a given pure compound always contains: the same elements in

exactly the same proportions by mass e.g.: Water will always consist of 2 atoms of hydrogen and one atom of oxygen.

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Molecular formulas contain a collection of elemental symbols which as a group represents one molecule.

Formula Name Proportions

CO Carbon monoxide 1 atom of carbon1 atom of oxygen

H2O water 2 atoms of hydrogen1 atom of oxygen

Properties of Pure Substancesa. Physical properties

can be observed without changing the composition of the substance. include: color, odor, taste, solubility , density, melting point and boiling point

b. Chemical properties can be observed when a substance undergoes a change in composition. include: iron rusting, gasoline burning in air, water undergoing electrolysis and chlorine

reacts with sodium

Changes of Pure Substances

a. Physical changes occur without a change in the composition of the substance (conversion from one state

of matter to another)

b. Chemical changes observed when a change in the composition of a substance occurs. New substances are

formed with different physical and chemical properties.Cl2 + Na ---> 2 NaCl

chlorine gas (poisonous) plus sodium (reactive metal) produces a new substance--sodium chloride (table salt) which has totally different physical and chemical properties

Energy Energy is defined as the ability to do work or to transfer heat

principal types of energy are: mechanical, heat, electrical, chemical and light Energy can either be

potential energy - energy possessed by its position in space or its chemical composition

kinetic energy - due to motion

Heat energy - energy is transferred from one substance to another when there is a temperature difference between the substances.

Measurement of Energy (units) Calorie (cal) Joule (j) 1 cal = 4.184 J

A calorie is the quantity of heat required to raise the temperature of 1 gram of water from 14.5 oC to 15.5 oC

Specific Heat

Specific heat (SpH) physical property amount of energy (calories) required to produce a given change in temperature ( oC) in

relation to the mass of a substance

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Chemistry-handouts

in equation:

specific heat= caloriesmass(∆T )

Example:Calculate the SpH of a metal with a mass of 25 grams and it takes 250 calories of heat energy to raise the temperature of the metal from 20 oC to 25 oC.

specific heat= 250 calories25g (25℃−20℃)

=2.0cal /g℃

Law of Conservation of Mass and Conservation of Energy Law of Conservation of Mass:

Mass can neither be created nor destroyed. The total mass of any system always remains constant.

Law of Conservation of Energy: Energy can neither be created nor destroyed. Energy may be transformed from one type to another.

Division of the Elements metals nonmetals metalloids

Metals have certain physical properties: high luster; conduct electricity; malleable; ductile; most have high densities; many have high

m.p.; most are hardMetals have certain chemical properties:

do not readily combine with other metals do combine with nonmetals few are found in the free state (example Au, Ag, Cu, and Pt)

Nonmetals have certain physical properties: Not lusterous; poor conductors; not malleable or ductile; brittle; most have low densities; many

have low melting points (mp); most are softNonmetals have certain chemical properties:

combine with metals or other nonmetals few exist in nature in the free state

Dalton's Atomic Theory elements are composed of atoms atoms are indestructible atoms of the same element are identical atoms combine in whole number ratios atoms of different elements can unite in different ratio

The Structure of the Atom The atom consists of three principle subatomic particles:

Particle Symbol Charge Relative mass

Electron E Positive (+) Approx. zero (0)

Proton P Negative (-) 1

Neutron N Neutral (o) 1

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Mass Number

sum of the number of protons and neutrons.

Mass Number = number of protons + number of neutrons

Atomic number

refer to the number of protons or electrons of an atom

Ions

charged atoms could be:

cation (+) anion (- )

The Mole

amount of substance 1 doz = 12 1 mole = 6.02 x 10 23 atoms molecules, particles

equation:

n= MMW

where:n =amount of substance, molesM = mass of substance, kgMW = molar mass/molecular mass

Converting moles to grams

How many grams of lithium are in 3.50 moles of lithium?

M=n x MW

M=3.50moles x 7g Li1mole Li

=24.5 g Li

Converting grams to moles

How many moles of lithium are in 18.2 grams of lithium?

n= MMW

n=18.2 gLi7g Li

=2.6mole Li

Using Avogadro’s Number

How many atoms of lithium are in 3.50 moles of lithium?

N A=n x N

N A=3.50moles Li x 6.02 x1023 atoms

N A=2.07 x1024atoms

Formulas Empirical formula: the lowest whole number ratio of atoms in a compound. Molecular formula: the true number of atoms of each element in the formula of a compound.

molecular formula = (empirical formula)n [n = integer] molecular formula = C6H6 = (CH)6

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empirical formula = CH

STOICHIOMETRY

Chemical Equations

C2H5OH + 3O2 2CO2 + 3H2O

reactants products

Quantitative significance:

1 mole of ethanol reacts with 3 moles of oxygen to produce 2 moles of carbon dioxide and 3 moles of water

Stoichiometry Problem 1

6.50 grams of aluminum reacts with an excess of oxygen. How many grams of aluminum oxide are formed?

4 Al + 3 O2 2Al2O3

6.50 g Al x1mole Al27 g Al

x2mol Al2O34mol Al

x101.96 gmol Al2O31molmol Al2O3

=12.3gmol Al2O3

Stoichiometry Problem 2

How many grams of hydrogen will be produced if 10.0 grams of calcium is added to an excess of hydrochloric acid?

2 HCl + Ca CaCl2 + H2

10 gCa x1moleCa40 gCa

x1mole H 2

1moleCax2 g H 2

1mole H 2

=0.5 g H 2

GASES

Characteristics of Gases Expansion - gases expand indefinitely and uniformly to fill all the space in which they are placed.

Indefinite Shape and Volume - gas has no definite shape or volume, but will fit the vessel in which it is placed.

Compressibility - gases can be highly compressed.

Low density - density of gases is very low and, therefore, measured in grams/liter (g/l)

Diffusion - two or more different gases will normally mix completely and uniformly when in contact with each other.

Pressure of Gases Pressure is defined as force per unit area (force/area). STP (standard temperature and pressure) - the conditions are 0 oC and 1 atm. Units of pressure

14.7 psi = 1 atm 76.0 cm Hg = 1 atm 760 mm Hg = 1 atm 760 torr = 1 atm 1.013 x 105 Pa = 1 atm

Gas LawsA. Boyle's Law

Boyles Law - at constant temperature, the volume is inversely proportional to the pressure.

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As the pressure increases, the volume decreases; as the pressure decreases, the volume increases

Example:A gas has a volume of 2 liters at a pressure of 6 atm. Calculate the new volume if the pressure was decreased to 3 atm and the temperature remained the same.

B. Charles's Law Charles Law - at constant pressures the volume is directly proportional to the temperature.

As the temperature increases, the volume increases, as the temperature decreases, the volume decreases

Example:A gas has a volume of 2 liters at a temperature of 600 K. Calculate the new volume if the temperature was decreased to 300 K and the pressure remains the same.

C. Gay-Lussac's Law Gay-Lussac Law - at constant volume, the pressure is directly proportional to the temperature.

As the temperature increases, the pressure increases; as the temperature decreases, the pressure decreases

Example:A gas has a pressure of 700 torr at a temperature of 300 K. Calculate the new pressure if the temperature was increased to 600 K and the volume remains the same.

In Gay-Lussac's Law, the pressure is proportional to the temperature. Therefore, if the temperature is doubles, the pressure would be expected to double.

D. General Gas Law Boyles and Charles Law can be combined into a single equation.

Example: solve for pressure2 liters of a gas has an initial pressure of 800 torr at a temperature of 300 K. What would be the new pressure if the temperature was increased to 600 K and the volume remained the same?

Solution:Using the general gas law formula, solve for P2

Example: solve for temperature

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A gas has an initial pressure of 4 atm at a temperature of 300 K with an initial volume of 2 liters. What would be the new temperature if the pressure was increased to 8 atm and the volume doubled?

Solution:Using the general gas law formula, solve for T2

Example: solve for volumeA gas has an initial pressure of 30 psi at a temperature of 600 K with an initial volume of 2 liters. What would be the new volume if the pressure was decreased to 15 psi and the temperature remained the same?

Solution:Using the general gas law formula, solve for V2

E. Ideal Gas Law

Example: solve for molesA gas has a pressure of 4 atm. at a temperature of 300 K with an initial volume of 2000 ml. Calculate the moles of the gas.

Solution:Using the ideal gas law formula, solve for moles.

2. Convert all units to either atm., liters, or kelvin.

Example: solve for pressure2 moles of a gas has a volume of 0.75 liters at a temperature of 27 oC. Calculate the pressure of the gas.

Solution:Using the ideal gas law formula, solve for pressure.

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Example: solve for the gas constant1 mole of a gas at STP (273 K and 1 atm) has a volume of 22.4 liters. Calculate the gas constant R.

Solution:Using the ideal gas law formula, solve for the gas constant R.

Example: solve for temperatureNitrogen gas (N2 - 28 amu) has a volume of 5 liters, a pressure of 15 atm and a mass of 56 grams. Calculate the temperature of the gas.

Solution:Using the ideal gas law formula, solve for temperature.

Example: solve for volume

A gas has a temperature of 300 K, a pressure of 22 psi and 2 moles. Calculate the volume of the gas.

Solution:Using the ideal gas law formula, solve for volume.

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Example: solve for mass

Chlorine gas (Cl2 - 71 amu) has a volume of 2 liters, a pressure of 4 atm and a temperature of 300 kelvin. Calculate the mass of chlorine.

Solution:Using the ideal gas law formula, solve for moles.

grams = moles x mol wt

grams = 0.32 moles x 71 amu = 22.7 grams

SOLUTIONS

Solution homogeneous mixture of two or more substances. generally composed of two substances

solute ---- substance being dissolved and is in smaller quantity. solvent --- substance that dissolves the solute and is in larger quantity

Types of solutions liquid solution

e.g. salt solution (salt & water) gaseous solution

e.g. air solution (various gases like oxygen, nitrogen, carbon dioxide, water and other gases)

solid solutions e.g. metal alloys like brass (copper and zinc)

Factors that Affect Solubility

nature of solute and solvent “LIKE DISSOLVES LIKE". Water dissolves salts, but not oil or gasoline.

Temperature the higher the temperature, more solute will dissolve in a given solvent.

Pressure

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pressure has little or no effect on the solubility of a solid or liquid, but the solubility of gases is greatly affected by pressure (Henry's Law).

Henry's Law Henry's Law - the solubility of a gas in a liquid is directly proportional to the applied

pressure. ---the higher the pressure, the more gas dissolves in a liquid. e.g. Carbonated soft drinks (carbon dioxide gas dissolved in a liquid (water) at

high pressure)

Factors that affect the Rate of Dissolution Particle size

the smaller the particle size, the faster the rate of dissolution. Powders dissolve faster than large lumps.

Temperature the higher the temperature, the faster is the rate of dissolution. It is easier to

dissolve sugar in a glass of hot tea than it is in a glass of iced tea.

Rate of stirring stirring or agitation of the solute in a solvent causes it to dissolve faster.

Concentration if a solution has some solute dissolved in it, the rate of dissolving additional

solute will be slower.

Relative Terms for Expressing Solute Concentration

Solubility the amount of solute that can be dissolved in a given amount of solvent.

Non-quantitative terms: Concentrated solution: the solution contains more solute than a dilute solution.

Dilute solution: the solution contains less solute than concentrated solution.

Unsaturated solution: the solution can dissolve more solute.

Saturated solution: the solution contains the maximum amount of solute the solvent can dissolve.

Supersaturated solution: the solution contains more solute than a saturated solution. This is very unstable condition and slight disturbance causes the excess solute to settle out.

Note: All of these are temperature dependent. A solution that is saturated at 25oC may be unsaturated at 45oC.

Quantitative Terms for Expressing Solute Concentration

A. Percent Concentration of Solute expressed as the percentage of the solute based on the entire solution.

a. The solution is the total of the solute and the solvent. The concentration can be expressed as percent :

a. the percent by mass

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b. percent by volume (for two liquids)

c. combination of the mass of the solute and the volume of the solution.

Example: (% by mass)

20 grams of NaCl is is mixed with 180 ml of water, what would be concentration as %NaCl by mass in this solution?

mass water = 180 g (remember that the density of water = 1.0 g/ml)

Example (% by volume): 200 ml of antifreeze is mixed with 0.8 liters of water, what would be % concentration of antifreeze by volume in this solution?

volume water = 800 ml (remember 0.8 liters = 800 ml)

Example (combination of the mass of the solute and the volume of the solution)150 ml of ethanol is mixed with 1.5 liters of water, what would be % concentration of ethanol by mass/volume in this solution?

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volume water = 1500 ml (remember 1.5 liters = 1500 ml) mass ethanol = 118.5 g (density = 0.79 g/ml and mass = density x vol)

B. The concentration can be expressed as parts by mass of solute per million parts of solution (ppm) (ppm)

Example: A drinking water sample has 15 mg of Pb per 500 ml of solution. Calculate the ppm of Pb in the water.

mass solution = 500 ml = 500 g = 500,000 mg

C. Molarity (M) moles of solute per liter of solution

Example: 120 grams of NaOH is diluted to 750 ml. What would be the molarity of this solution.

First, write the molarity formulas

0.75 liters (750 ml = 0.75 l) determine the formula mass of NaOH calculate the moles of NaOH

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calculate the molarity of NaOH

Example:

500 ml of a 2.5 M solution of Ca(OH)2 needs to be made. What would be the mass of Ca(OH)2 needed to dilute to 800 ml.

write the molarity formulas

convert ml to liters (500 ml / 1000 = liters)

0.5 liters (500 ml = 0.5 l)

determine the formula mass of Ca(OH)2

Ca(OH)2 = 74.1 amu

calculate the moles of Ca(OH)2

calculate the grams of Ca(OH)2

To make the 2.5 M Ca(OH)2 solution, weigh out 92.6 grams of Ca(OH)2 and dilute with water to 800 ml.

D. Normality (N) used when more precise measurements of the concentration of solutions are needed

and the solutions are being use in reactions of acids and bases.

Note that equivalents can be defined in several ways. Here are just a few.

The number of moles of H+ or OH- ions replaced in a chemical reaction.

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The number of replaceable H+ or OH- ions in a compound. The number of moles of electrons transferred in a chemical reaction (red-ox).

Example: 296.4 grams of Ca(OH)2 is diluted to 600 ml. What would be the normality of this solution.

write the normality formulas

convert 600 ml to liters (600ml / 1000 = liters)

0.6 liters (600 ml = 0.6 L)


Ca(OH)2 = 74.1 amu

calculate the number of replaceable OH- in Ca(OH)2

Ca(OH)2 has 2 replaceable hydroxides

calculate the equivalent mass of Ca(OH)2

calculate the equivalents of Ca(OH)2

calculate the normality of Ca(OH)2

Example: 600 ml of a 2.5 M solution of Ca(OH)2 needs to be made. What would be the mass of Ca(OH)2 needed to dilute to 600 ml.

write the normality formulas

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Chemistry-handouts

convert 600 ml to liters (600 ml / 1000 = liters)0.6 liters (600 ml = 0.6 L)


Ca(OH)2 = 74.1 amu

calculate the number of replaceable OH- in Ca(OH)2

Ca(OH)2 has 2 replaceable hydroxides

calculate the equivalent mass of Ca(OH)2

calculate the equivalents of Ca(OH)2

calculate the normality of Ca(OH)2

To make the 2.5 M Ca(OH)2 solution, weigh out 55.6 grams of Ca(OH)2 and dilute with water to 600 ml.

E. Dilution Formula used to dilute a concentrated stock solution (these stock solutions are concentrated so

as to save room during storage and costs in shipping) to a desired concentration.

C1 = the initial concentrationC2 = the final concentrationV1 = the initial volumeV2 = the final volume

Example:

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A stock solution of hydrochloric acid is has a concentration of 12M (C1). 100 ml (V2) of a 6M HCl (C2) solution needs to be made from the stock solution. How many ml (V1) of the stock solution needs to be diluted to 100 ml (V2) to make the 6M (C2) solution?

substitute the actual values for the variables in the formula.

Electrolytes vs. Nonelectrolytes

electrolytes substance whose aqueous solution conducts an electric current

nonelectrolytes substance whose aqueous solution does not conduct an electric current

ACIDS AND BASES

Acid/Base Definitions Arrhenius Model

Acids produce hydrogen ions in aqueous solutions Bases produce hydroxide ions in aqueous solutions

Bronsted-Lowry Model Acids are proton donors Bases are proton acceptors

Lewis Acid Model Acids are electron pair acceptors Bases are electron pair donors

Properties of Acids Acids are proton (hydrogen ion, H+) donors Acids have a pH lower than 7 Acids taste sour Acids effect indicators

Blue litmus turns red Methyl orange turns red

Acids react with active metals, producing salts and hydrogen gas(H2) Acids react with carbonates Acids neutralize bases

Strong Acids

assumed to be 100% ionized in solution (good H+ donors). include:

hydrochloric acid ----- HCl Sulfuric acid ------ H2SO4

Nitric acid ------ HNO3

Hydriodic acid ----HI Perchloric acid -----HClO4

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Hydrobromic acid ----HBr Hydronium ion -----H3O+

Weak acids include:

Iodic acid ----HIO3

Oxalic acid ---- H2C2O4

sulfurous acid ----H2SO3

phosphoric acid ----H3PO4

nitrous acid ----HNO2

Hydroflouric acid –HF formic acid ---- HCOOH benzoic acid ---C6H5COOH

acetic acid -----CH3COOH

crbonic acid ---H2CO3

organic acids and their sources Citric acid ---H3C6H5O7 – citrus fruit Malic acid – apples Butyric acid – rancid butter Amino acids – protein Nucleic acids – DNA and RNA Ascorbic acid – Vitamin C

Properties of Bases Bases are proton (hydrogen ion, H+) acceptors Bases have a pH greater than 7 Bases taste bitter Bases effect indicators

Red litmus turns blue Phenolphthalein turns purple

Solutions of bases feel slippery Bases neutralize acids

Examples of Bases Sodium hydroxide (lye)---- NaOH Potassium hydroxide ----KOH Magnesium hydroxide ----- Mg(OH)2

Calcium hydroxide (lime) ----- Ca(OH)2

Weak bases ammonia ----NH3

methylamine -----CH3NH2

ethylamine -----C2H5NH2

Diethylamine ---- (C2H5)2NH

Hdroxylamine --- HONH2

Hydrazine ------ H2NNH2

Aniline ---- C6H5NH2

Pyridine ----C5H5N

Self-Ionization of Water

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At 25, [H3O+] = [OH-] = 1 x 10-7

Kw is a constant at 25 C Kw = [H3O+][OH-] Kw = (1 x 10-7)(1 x 10-7) = 1 x 10-14

Calculating pH, pOH

pH = -log10(H3O+) pOH = -log10(OH-)

Relationship between pH and pOH

pH + pOH = 14

Finding [H3O+], [OH-] from pH, pOH

[H3O+] = 10-pH

[OH-] = 10-pOH

Example:

What is the pH of a 0.50 M solution of acetic acid, HC2H3O2, Ka = 1.8 x 10-5? Solve for pH and pOH of the solution

write the equationHC2H3O2 C2H3O2

- + H+

0.50 – x x x

[H+] = 3.0 x 10-3 M

Reaction of Weak Bases with Water

B + H2O BH+ + OH-

Example

What is the pH of a 0.50 M solution of ammonia, NH3, Kb = 1.8 x 10-5 ?

Write the equation for the reactionNH3 + H2O NH4

+ + OH-

Set up the law of mass actionNH3 + H2O NH4

+ + OH-

50-x x x

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1 .8 x 10−5 =( x )( x )

( 0.50 − x )≃ x2

(0 .50)

1 .8 x 10−5 = x2

(0 .50 )

pH =−log(3 .0 x 10−5 )= 4 .52


Chemistry-handouts

[OH-] = 3.0 x 10-3 M

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1 .8 x 10−5 =( x )( x )

( 0.50 − x )≃ x2

(0 .50)

1 .8 x 10−5 = x2

(0 .50 )

pOH =−log (3 .0 x 10−5 )= 4 .52

pH = 14 .00 − pOH = 9. 48

2015 chemhandouts1

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