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AP Subject Review

3 Stoichiometry

Percentage Composition

Calculate the percent of each element in the total mass of the compound

(#atoms of the element)(atomic mass of element) x 100 (molar mass of the compound

Determining the empirical formula

Determine the percentage of each element in your compound

Treat % as grams, and convert grams of each element to moles of each element

Find the smallest whole number ratio of atoms

If the ratio is not all whole number, multiply each by an integer so that all elements are in whole number ratio

Determining the molecular formula

Find the empirical formula mass

Divide the known molecular mass by the empirical formula mass, deriving a whole number, n

Multiply the empirical formula by n to derive the molecular formula

Examples:

60 grams propane gas is burned in excess oxygen: C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(l)

How much water is produced?

7.321 mg of an organic compound containing carbon, hydrogen, and oxygen was analyzed by combustion. The amount of

carbon dioxide produced was 17.873 mg and the amount of water produced was 7.316 mg. Determine the empirical

formula of the compound.

Sodium metal reacts vigorously with water to produce a solution of sodium hydroxide and hydrogen gas:

2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g)

What mass of hydrogen gas can be produced when 10 grams of sodium is added to 15 grams of water?

0.1101 gram of an organic compound containing carbon, hydrogen, and oxygen was analyzed by combustion. The

amount of carbon dioxide produced was 0.2503 gram and the amount of water produced was 0.1025 gram. A

determination of the molar mass of the compound indicated a value of approximately 115 grams/mol. Determine the

empirical formula and the molecular formula of the compound.

4 Chemical Equations and Stoichiometry

Solving Stoichiometry Problems for Reactions in Solution

Identify the species present in the combined solution, and determine what reaction occurs.

Write the balanced net ionic equation for the reaction.

Calculate the moles of reactants.

Determine which reactant is limiting.

Calculate the moles of product(s), as required.

Convert to grams or other units, as required.

Example:

10.0 mL of a 0.30 M sodium phosphate solution reacts with 20.0 mL of a 0.20 M lead(II) nitrate solution (assume no

volume change).

Types of Reactions

Precipitation Reactions

A double displacement reaction in which a solid forms and separates from the solution.

Simple Rules for Solubility

Most nitrate (NO3-) salts are soluble.

Most alkali metal (group 1A) salts and NH4+ are soluble.

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Most Cl-, Br

-, and I

-salts are soluble (except Ag

+, Pb

2+, Hg2

2+).

Most sulfate salts are soluble (except BaSO4, PbSO4, Hg2SO4, CaSO4).

Most OH- are only slightly soluble (NaOH, KOH are soluble, Ba(OH)2, Ca(OH)2 are marginally soluble).

Most S2-

, CO32-

, CrO42-

, PO43-

salts are only slightly soluble, except for those containing the cations in Rule 2.

Complete Ionic Equation

All substances that are strong electrolytes are represented as ions.

Ag+(aq) + NO3

-(aq) + Na

+(aq) + Cl

-(aq)

AgCl(s) + Na+(aq) + NO3

-(aq)

Net Ionic Equation

Includes only those solution components undergoing a change.

Show only components that actually react.

Acid–Base Reactions (Brønsted–Lowry)

Acid—proton donor

Base—proton acceptor

For a strong acid and base reaction:

H+(aq) + OH

–(aq) H2O(l)

Redox Reactions

Reactions in which one or more electrons are transferred.

Rules for Assigning Oxidation States

Oxidation state of an atom in an element = 0

Oxidation state of monatomic ion = charge of the ion

Oxygen = -2 in covalent compounds (except in peroxides where it = -1)

Hydrogen = +1 in covalent compounds

Fluorine = -1 in compounds

Sum of oxidation states = 0 in compounds

Sum of oxidation states = charge of the ion in ions

Balancing Oxidation–Reduction Reactions by Oxidation States

Write the unbalanced equation.

Determine the oxidation states of all atoms in the reactants and products.

Show electrons gained and lost using “tie lines.”

Use coefficients to equalize the electrons gained and lost.

Balance the rest of the equation by inspection.

Add appropriate states.

Example:

What coefficients are needed to balance the remaining elements?

Zn(s) + 2HCl(aq) Zn2+

(aq) + 2Cl–(aq) + H2(g)

dilute nitric acid is added to crystals of pure calcium oxide.

hydrogen sulfide is bubbled through a solution of silver nitrate.

sodium metal is added to water.

a solution of tin(II) chloride is added to a solution of iron(III) sulfate.

phosphorus(V) oxytrichloride is added to water.

solid sodium oxide is added to water.

excess concentrated potassium hydroxide solution is added to a precipitate of zinc hydroxide.

excess concentrated sodium hydroxide solution is added to solid aluminum hydroxide.

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5 Gases

Standard conditions

STP = 1 atm (760mmHg or 101kPa) and 273K ALL temperatures must be in Kelvin! (C + 273)

Gas Laws:

Boyles Law: P1V1 = P2V2

Charles Law: V1/T1 = V2/T2

Guy Lussac’s Law: P1/T1 = P2/T2

Combined Gas Law: V1P1/n1T1 = V2P2/n2T2

Ideal Gas Law: PV = nRT volume in liters R = 0.0821 atm L/molK

Daltons Law: Pt = P1 + P2 …

Mole Fraction: mol A/total moles

Root Mean Square: FW

RTuu rms

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Grahams Law: FW

FW =

2 gas of effusion of Rate

1 gas of effusion of Rate

1

2

Examples:

A sample of hydrogen gas (H2) has a volume of 8.56 L at a temperature of 0ºC and a pressure of 1.5 atm. Calculate the

moles of H2 molecules present in this gas sample.

A sample of gas at 15ºC and 1 atm has a volume of 2.58 L. What volume will this gas occupy at 38ºC and 1 atm ?

Sulfur dioxide (SO2), a gas that plays a central role in the formation of acid rain, is found in the exhaust of automobiles

and power plants. Consider a 1.53- L sample of gaseous SO2 at a pressure of 5.6 x 103 Pa. If the pressure is changed to 1.5

x 104 Pa at a constant temperature, what will be the new volume of the gas ?

Suppose we have a sample of ammonia gas with a volume of 3.5 L at a pressure of 1.68 atm. The gas is compressed to a

volume of 1.35 L at a constant temperature. Use the ideal gas law to calculate the final pressure

Mixtures of helium and oxygen are used in scuba diving tanks to help prevent “the bends.” For a particular dive, 46 L He

at 25ºC and 1.0 atm and 12 L O2 at 25ºC and 1.0 atm were pumped into a tank with a volume of 5.0 L. Calculate the

partial pressure of each gas and the total pressure in the tank at 25ºC.

Calculate the ratio of the effusion rates of hydrogen gas (H2) and uranium hexafluoride (UF6), a gas used in the

enrichment process to produce fuel for nuclear reactors.

7 • Atomic Theory and Bonding

Know the names and locations of the groups on the periodic table: alkali metals, alkaline earth metals, inner and outer

transition metals, halogens and noble gases and their common charges as ions

Electrons

Be able to write electron configurations for any element

Aufbau Principle Electrons fill their orbitals from lower to higher energy: 1s2s293s3p4s3d4p5s4d5p6s4f5d6p7s

Hund’s Rule An electron will half fill an orbital until all orbitals of that sublevel are occupied

Periodic Trends

Ionization Energy

o Energy required to remove an electron from a gaseous atom or ion. X(g) → X+(g) + e

–

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o In general, as we go across a period from left to right, the first ionization energy increases.

Why? Electrons added in the same principal quantum level do not completely shield the

increasing nuclear charge caused by the added protons.

o In general, as we go down a group from top to bottom, the first ionization energy decreases.

Why? The electrons being removed are, on average, farther from the nucleus.

Electron Affinity

o Energy change associated with the addition of an electron to a gaseous atom. X(g) + e– → X

–(g)

o In general as we go across a period from left to right, the electron affinities become more negative.

o In general electron affinity becomes more positive in going down a group.

Atomic Radius

o In general as we go across a period from left to right, the atomic radius decreases.

o Effective nuclear charge increases, therefore the valence electrons are drawn closer to the nucleus,

decreasing the size of the atom.

o In general atomic radius increases in going down a group.

o Cations have lost electrons and are generally smaller than the neutral atom (more nuclear attraction)

o Anions have gained electrons and are larger than the neutral atom (less nuclear attraction)

Examples

Write the electron configuration for Strontium and Sodium using the shorthand notation for the noble gas cores.

How many unpaired electrons are there in a nitrogen atom?

Arrange the elements S, Ge, P, and Si in order of increasing atomic size.

Arrange the ions Na+, K

+, Cl , and Br in order of increasing size.

Arrange the elements Be, Ca, N, and P in order of increasing ionization energy.

8•Bonding

Types of Chemical Bonds

Ionic Bonding – electrons are transferred

Covalent Bonding – electrons are shared equally by nuclei

Polar Covalent Bond Unequal distribution of electrons between atoms in a molecule resulting in a partial positive

and negative region in the molecule

Nonpolar covalent Bond: Equal distribution of the electron cloud

Electronegativity:

increases as you go across the period table and decreases as you go down. Fluorine has an electronegativity of 4

and is the most electronegative element. Noble gases have 0 electronegativity. The greater the e difference, the

more ionic charater the bond has and the stronger it is.

Dipole Moment

Property of a molecule whose charge distribution can be represented by a center of positive charge and a center of

negative charge.

Use an arrow to represent a dipole moment.

Lattice Energy

The change in energy that takes place when separated gaseous ions are packed together to form an ionic solid.

k = proportionality constant

Q1 and Q2 = charges on the ions

r = shortest distance between the centers of the cations and anions

Bond Energy

ΔH = Σn X D(bonds broken) – Σn X D(bonds formed)

Lewis Structure

Shows how valence electrons are arranged among atoms in a molecule.

Octet Rule All atoms should have 8 electrons exept for hydrogen with 2

Steps for Writing Lewis Structures

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Sum the valence electrons from all the atoms.

Use a pair of electrons to form a bond between each pair of bound atoms.

Arrange the remaining electrons to satisfy the octet rule (or duet rule for hydrogen).

Resonance occurs when you have a combination of multiple and single bonds:

Formal Charge

Formal charge = (# valence e– on free neutral atom) – (# valence e

– assigned to the atom in the molecule).

VSEPR Model

Hybridization

explains why bonds in molecules with different atomic orbitals behave as identical bonds ie. CH4

sigma bond: overlap of two s orbitals or an s and a p orbital or head-to-head p orbitals.

Electron density of a sigma bond is greatest along the axis of the bond.

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Pi (π) bonds--come from the sideways overlap of p atomic orbitals; the region above and below the internuclear

axis. NEVER occur without a sigma bond first!

HYBRIDIZATION # OF HYBRID

ORBITALS

GEOMETRY

sp 2 Linear

sp2 3 Trigonal planar

sp3 4 Tetrahedral

dsp3 5 Trigonal bipyramidal

d2sp

3 6 Octahedral

Example

Draw a Lewis structure for NH3; CH4; SF6

What molecular shapes are associated with the following electron pairs around the central atom?

a. 3 bonding pairs and 2 lone pairs

b. 4 bonding pairs and 1 lone pair

Name the shapes and hybridization for the following molecules or polyatomic ions:

a. O3

b. GaH3

Determine whether the following molecules are polar or nonpolar:

a. CCl4

b. XeF4

10 • Liquids and Solids

Intermolecular Forces

London Dispersion forces found in all molecules. It is the only IMF in nonpolar molecules. The larger the

molecule, the stronger the LDFs

Dipole dipole forces found in polar molecules

Hydrogen bonding between H and F,O or N

A polar molecule has polar bonds and is asymmetric

Some Properties of a Liquid

Surface Tension: The resistance to an increase in its surface area (polar molecules). High ST indicates strong

IMF’s.

Capillary Action: Spontaneous rising of a liquid in a narrow tube.

Viscosity: Resistance to flow (molecules with large intermolecular forces). Modeling a liquid is difficult.

Gases have VERY SMALL IMFs and lots of motion. Solids have VERY HIGH IMFs and next to no motion. Liquids

have both strong IMFs and quite a bit of motion.

Types of Solids

Crystalline Solids: highly regular arrangement of their components [often ionic, table salt (NaCl), pyrite (FeS2)].

Amorphous solids: considerable disorder in their structures (glass).

Representation of Components in a Crystalline Solid

Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make

up the substance.

(a) network covalent—carbon in diamond form—here each molecule is covalently bonded to each neighboring C

with a tetrahedral arrangement. Graphite on the other hand, make sheets that slide and is MUCH softer! (pictured

later)

(b) ionic salt crystal lattice

(c) ice—notice the “hole” in the hexagonal structure and all the H-bonds. The “hole” is why ice floats—it makes

it less dense than the liquid!

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Types of Crystalline Solids

Ionic Solid: contains ions at the points of the lattice that describe the structure of the solid (NaCl). VERY high

MP’s. Hard. Ion-Ion Coulombic forces are the strongest of all attractive forces. “IMF” usually implies

covalently bonded substances, but can apply to both types.

Molecular Solid: discrete covalently bonded molecules at each of its lattice points (sucrose, ice).

Atomic Solid: atoms of the substance are located at the lattice points. Carbon—diamond, graphite and the

fullerenes. Boron, and silicon as well.

Know this chart well:

Structure and Bonding in Metals

Metals are characterized by high thermal and electrical conductivity, malleability, and ductility. These properties are

explained by the nondirectional covalent bonding found in metallic crystals.

Electron Sea Model: A regular array of metals in a “sea” of electrons. I A & II A metals pictured at left.

Band (Molecular Orbital) Model: Electrons assumed to travel around metal crystal in MOs formed from valence

atomic orbitals of metal atoms.

Metal alloys: a substance that has a mixture of elements and has metallic properties

substitution alloys—in brass 1/3 of the atoms in the host copper metal have been replaced by zinc atoms. Sterling

silver—93% silver and 7% copper. Pewter—85% tin, 7% copper, 6% bismuth and 2% antimony. Plumber’s

solder—95% tin and 5% antimony.

interstitial alloy—formed when some of the interstices [holes] in the closest packed metal structure are occupied

by small atoms. Steel—carbon is in the holes of an iron crystal. There are many different types of steels, all

depend on the percentage of carbon in the iron crystal.

Network Atomic Solids—a.k.a. Network Covalent

Composed of strong directional covalent bonds that are best viewed as a “giant molecule”. Both diamond and

graphite are network solids. The difference is that diamond bonds with neighbors in a tetrahedral 3-D fashion,

while graphite only has weak bonding in the 3rd

dimension. Network solids are often:

brittle—diamond is the hardest substance on the planet, but when a diamond is “cut” it is actually fractured to

make the facets

do not conduct heat or electricity

carbon, silicon-based

Diamond is hard, colorless and an insulator. It consists of carbon atoms ALL bonded tetrahedrally, therefore sp3

hybridization and 109.5 bond angles.

11 • Solutions

Definitions:

Solution- a homogeneous mixture of two or more substances in a single phase.

solute--component in lesser concentration; dissolvee

solvent--component in greater concentration; dissolver

solubility--maximum amount of material that will dissolve in a given amount of solvent at a given temp. to

produce a stable solution.

Saturated solution- a solution containing the maximum amount of solute that will dissolve under a given set of

conditions.

Unsaturated solution- a solution containing less than the maximum amount of solute that will dissolve under a

given set of conditions. (more solute can dissolve)

Supersaturated solution- a solution that has been prepared at an elevated temperature and then slowly cooled.

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miscible—When two or more liquids mix (ex. Water and food coloring)

immiscible—When two or more liquids DON’T mix.--they usually layer if allowed to set for a while. (ex.

Water and oil)

Units of solution concentration

Molarity (M) = # of moles of solute per liter of solution

Mole fraction () = ratio of the number of moles of a given component to the total number of moles of solution.

Molality (m) = # of moles of solute per kilogram of solvent

Heat of solution (Hsoln) = the enthalpy change associated with

Factors Affecting Solubility

Pressure Effects:

o The solubility of a gas is higher with increased pressure.

o Henry’s Law- the amount of a gas dissolved in a solution is directly proportional to the pressure of the gas

above the solution. P = kC

o P = partial pressure of the gaseous solute k = constant C = concentration of the gas

Temperature Effects:

o The amount of solute that will dissolve usually increases with increasing temperature. Solubility

generally increases with temperature if the solution process is endothermic (Hsoln > 0). Solubility

generally decreases with temperature if the solution process is exothermic (Hsoln < 0). Potassium

hydroxide, sodium hydroxide and sodium sulfate are three compounds that become less soluble as the

temperature rises. This can be explained by LeChatelier’s Principle.

o The solubility of a gas in water always decreases with increasing temperature.

Colligative Properties- properties that depend on the number of dissolved particles

Vapor Pressure Lowering- The presence of a nonvolatile solute lowers the vapor pressure of a solvent. Raoult’s

Law: Psolution = (solvent) (Posolvent)

o Psolution = observed vapor pressure of the solvent in the solution

o solvent = mole fraction of solvent

o Po

solvent = vapor pressure of the pure solvent

An ideal solution is a solution that obeys Raoult’s Law. There is no such thing. In very dilute solutions

o We can find the molecular weight of a solute by using the vapor pressure of a solution.

Boiling Point Elevation

BPE = mki m = molality k = BP constant i= number of ions

Freezing point depression

FPD = mki m = molality k = BP constant i= number of ions

12 • Chemical Kinetics: Rates of Reaction

FACTORS THAT AFFECT REACTION RATES

Nature of the reactants--Some reactant molecules react in a hurry, others react very slowly.

Concentration of reactants--more molecules, more collisions.

Temperature--heat >em up & speed >em up;

Catalysts--accelerate chemical reactions but are not themselves transformed.

Surface area of reactants--exposed surfaces affect speed

Rate = change in concentration of a species per time interval

When writing rate expressions, they can be written in terms of reactants disappearance or products appearance.

* Rate is not constant, it changes with time. Graphing the data of an experiment will show an average rate of reaction.

You can find the instantaneous rate by computing the slope of a straight line tangent to the curve at that time.

reaction rate--expressed as the Δ in concentration of a reagent per unit time or Δ[A]/Δt

Initial rxn rate = k[A]m

[B]n

[C]p

Exponents can be zero, whole numbers or fractions and are determined by experiment

ORDER OF A REACTION

order with respect to a certain reactant is the exponent on its concentration term in the rate expression

order of the reaction is the sum of all the exponents on all the concentration terms in the expression

Zero order: The change in concentration of reactant has no effect on the rate. These are not very common. General

form of rate equation: Rate = k

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First order: Rate is directly proportional to the reactants concentration; doubling [rxt], doubles rate. These are

very common! Nuclear decay reactions usually fit into this category. General form of rate equation: Rate = k [A]

Second order: Rate is quadrupled when [rxt] is doubled and increases by a factor of 9 when [rxt] is tripled etc.

These are common, particularly in gas-phase reactions.

General form of rate equation: Rate = k [A]2

TWO TYPES OF RATE LAWS

o differential rate law--expresses how the rate depends on concentration

o integrated rate law--expresses how the concentrations depend on time

INTEGRATED RATE LAW: CONCENTRATION/TIME RELATIONSHIPS

first order: second order:

ln[A] = -kt + ln[A]o 1/[A] = kt + 1/[A]o

HALF-LIFE AND REACTION RATE FOR FIRST ORDER REACTIONS, t1/2

the time required for one half of one of the reactants to disappear.

T1/2 = 0.693

k

Half life is INDEPENDENT OF ORIGINAL CONCENTRATION for 1st

order!!!

HALF-LIFE AND REACTION RATE FOR SECOND ORDER REACTIONS, t1/2

the time required for one half of one of the reactants to disappear.

k t2 = 1

[A]o

HALF-LIFE AND REACTION RATE FOR ZERO ORDER REACTIONS t2 = [A]o rxn

RATE EXPRESSIONS FOR ELEMENTARY STEPS--the rate expression cannot be predicted from overall

stoichiometry. The rate expression of an elementary step is given by the product of the rate constant and the

concentrations of the reactants in the step.

ELEMENTARY STEP MOLECULARITY RATE EXPRESSION

Aproducts unimolecular rate = k[A]

A + B products bimolecular rate = k[A][B]

A + A products bimolecular rate = k[A]

2

2 A + B products* termolecular* rate = k[A]

2[B]

THE EFFECT OF TEMPERATURE OF REACTION RATE: ARRHENIUS EQUATION

k = reaction rate constant = Ae-E*/RT

o R is the ―energy‖ R or 8.31 x 10-3

kJ/K mol

o A is the frequency factor units of L/(mol s) & depends on the frequency of collisions and the fraction

of these that have the correct geometry--# of effective collisions

o e-E*/RT

is always less than 1 and is the fraction of molecules having the minimum energy required for reaction

CATALYSIS

catalysts are not altered during the reaction--they serve to lower the activation energy and speed up the reaction by offering

a different pathway for the reaction

HETEROGENEOUS CATALYST--

different phase than reactants, usually involves gaseous reactants adsorbed on the surface of a solid catalyst

HOMOGENEOUS CATALYST—exists in the same phase as the reacting molecules.

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13 Chemical Equilibria

THE EQUILIBRIUM EXPRESSION:

A general description of the equilibrium condition proposed by Gudberg and Waage in 1864 is known as the Law of Mass

Action. Equilibrium is temperature dependent, however, it does not change with concentration or pressure.

equilibrium constant expression--for the general reaction

aA + bB cC + dD

Equilibrium constant: K = [C]c[D]

d

[A]a[B]

b

o Pure solids--do not appear in expression—you’ll see this in Ksp problems soon!

o Pure liquids--do not appear in expression—H2O (l) is pure, so leave it out of the calculation

o Water--as a liquid or reactant, does not appear in the expression. (55.5M will not change significantly)

CHANGING STOICHIOMETRIC COEFFICIENTS

o when the stoichiometric coefficients of a balanced equation are multiplied by some factor, the K is raised to the

power of the multiplication factor (Kn). 2x is K squared; 3x is K cubed; etc.

o REVERSING EQUATIONS

o take the reciprocal of K ( 1/K)

o ADDING EQUATIONS

o multiply respective K=s (K1 x K2 x K3 …)

Kc & Kp--NOT INTERCHANGEABLE! Kp = Kc(RT)Δn

where

o Δn is the change in the number of moles of gas going from reactants to products:

o R = universal gas law constant 0.0821 L atm/ mol K

o T = temperature in Kelvin

Kc = Kp if the number of moles of gaseous product = number of moles of gaseous reactant.

THE REACTION QUOTIENT

For use when the system is NOT at equilibrium.

For the general reaction

aA + bB cC + dD

Reaction quotient = Qc = [C]c[D]

d

[A]a[B]

b

Qc has the appearance of K but the concentrations are not necessarily at equilibrium.

1. If Q<K, the system is not at equilibrium:

2. If Q = K, the system is at equilibrium.

3. If Q>K, the system is not at equilibrium:

EXTERNAL FACTORS AFFECTING EQUILIBRIA

Le Chatelier=s Principle: If a stress is applied to a system at equilibrium, the position of the equilibrium will shift in the

direction which reduces the stress.

Temperature—exothermic: heat is a product; endothermic: heat is a reactant.

Adding or removing a reagent--shift tries to reestablish Q.

Pressure--increase favors the side with the least # of gas moles; the converse is also true.

catalysts--NO EFFECT on K; just gets to equilibrium faster!

14&15 Acids and Bases, Aqueous Equilibria

ACID-BASE THEORIES:

ARRHENIUS DEFINITION

acid--donates a hydrogen ion (H+) in water

base--donates a hydroxide ion in water (OH-)

This theory was limited to substances with those "parts"; ammonia is a MAJOR exception!

BRONSTED-LOWRY DEFINITION

acid--donates a proton in water

base--accepts a proton in water

conjugate acid-base pair--A pair of compounds that differ by the presence of one H+ unit. This idea is critical

when it comes to understanding buffer systems. Pay close attention here!

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LEWIS DEFINITION

acid--accepts an electron pair

base--donates an electron pair

This theory explains all traditional acids and bases + a host of coordination compounds and is used widely in

organic chemistry. Uses coordinate covalent bonds

monoprotic--acids donating one H+

(ex. HC2H3O2)

diprotic--acids donating two H+'s (ex. H2C2O4)

polyprotic--acids donating many H+'s (ex. H3PO4)

polyprotic bases--accept more than one H+; anions with -2 and -3 charges (ex. PO4

3- ; HPO4

2-)

ACIDS ONLY DONATE ONE PROTON AT A TIME!!!

RELATIVE STRENGTHS OF ACIDS AND BASES

Strength is determined by the position of the "dissociation Do Not confuse concentration with strength!

STRONG ACIDS:

Hydrohalic acids: HCl, HBr, HI

Nitric: HNO3

Sulfuric: H2SO4

Perchloric: HClO4

STRONG BASES:

Hydroxides OR oxides of IA and IIA metals

THE STRONGER THE ACID THE WEAKER ITS CB, the converse is also true.

WEAK ACIDS AND BASES: are in equilibrium

HA + H2O H3O+ + A

-

Ka = [H3O+][A

-] < 1

[HA]

for weak base reactions:

B + H2O HB+ + OH

-

Keq[H2O]2 = Kw = [H3O

+][OH

-]

Kw = 1.0 x 10-14

( Kw = 1.008 x 10

-14 @ 25 degrees Celsius)

Kw = Ka x Kb (another very beneficial equation)

The pH Scale

Used to designate the [H+] in most aqueous solutions where H

+ is small.

pH = - log [H+]

pOH = - log [OH-]

pH + pOH = 14

Calculating pH of Weak Acid Solutions

Calculating pH of weak acids involves setting up an equilibrium. Always start by writing the equation, setting up the

acid equilibrium expression (Ka), defining initial concentrations, changes, and final concentrations in terms of X,

substituting values and variables into the Ka expression and solving for X. (use the RICE diagram learned in general

equilibrium!)

Determination of the pH of a Mixture of Weak Acids

Only the acid with the largest Ka value will contribute an appreciable [H+]. Determine the pH based on this acid and

ignore any others.

Calculating pH of polyprotic acids

Acids with more than one ionizable hydrogen will ionize in steps. Each dissociation has its own Ka value.

The first dissociation will be the greatest and subsequent dissociations will have much smaller equilibrium constants. As

each H is removed, the remaining acid gets weaker and therefore has a smaller Ka. As the negative charge on the acid

increases it becomes more difficult to remove the positively charged proton.

ACID-BASE PROPERTIES OF SALTS: HYDROLYSIS

Salts are produced from the reaction of an acid and a base. (neutralization)

Salts are not always neutral. Some hydrolyze with water to produce acidic and basic solutions.

Neutral Salts- Salts that are formed from the cation of a strong base and the anion of a strong acid form neutral

solutions when dissolved in water. A salt such as NaNO3 gives a neutral solution.

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Basic Salts- Salts that are formed from the cation of a strong base and the anion of a weak acid form basic

solutions when dissolved in water. The anion hydrolyzes the water molecule to produce hydroxide ions and thus

a basic solution. K2S should be basic since S-2

is the CB of the very weak acid HS-, while K

+ does not hydrolyze

appreciably. S2-

+ H2O OH- + HS

-

Acid Salts- Salts that are formed from the cation of a weak base and the anion of a strong acid form acidic

solutions when dissolved in water. The cation hydrolyzes the water molecule to produce hydronium ions and

thus an acidic solution. NH4Cl should be weakly acidic, since NH4+ hydrolyzes to give an acidic solution, while

Cl- does not hydrolyze. NH4

+ + H2O H3O

+ + NH3

If both the cation and the anion contribute to the pH situation, compare Ka to Kb.

If Kb is larger, basic! The converse is also true.

1. Strong acid + strong base = neutral salt

2. Strong acid + weak base = acidic salt

3. Weak acid + strong base = basic salt

4. Weak acid + weak base = ? ( must look at K values to decide )

THE LEWIS CONCEPT OF ACIDS AND BASES

acid--can accept a pair of electrons to form a coordinate covalent bond

base--can donate a pair of electrons to form a coordinate covalent bond

6 & 16 Thermochemistry

ENERGY AND WORK

E = q(heat) + w(work)

Signs of q

+q if heat absorbed

–q if heat released

Signs of w

+ w if work done on the system (i.e., compression)

-w if work done by the system (i.e., expansion)

When related to gases, work is a function of pressure

w = -PV

NOTE: Energy is a state function. (Work and heat are not.)

ENTHALPY

H is a state function

H = q at constant pressure (i.e. atmospheric pressure)

Enthalpy can be calculated from several sources including:

Coffee-cup calorimetry. q = H @ these conditions.

Bomb calorimetry – weighed reactants are placed inside a steel container and ignited.

Heat capacity – energy required to raise temp. by 1 degree (Joules/ C)

Specific heat capacity (Cp) – same as above but specific to 1 gram of substance

change) re temperatuof (degrees material) of (

ferredheat trans ofquantity heat specific

g

Molar heat capacity -- same as above but specific to one mole of substance

(J/mol K or J/mol C )

Energy (q) released or gained -- q = mCpT

q = quantity of heat ( Joules or calories)

m = mass in grams

ΔT = Tf - Ti (final – initial)

Cp = specific heat capacity ( J/gC)

Specific heat of water (liquid state) = 4.184 J/gC ( or 1.00 cal/g C)

Heat lost by substance = heat gained by water

Enthalpy of a Reaction

Hrxn = Hf (products) - Hf (reactants)

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Hess’s Law

sum up the H’s for the individual reactions to get the overall Hrxn.

First decide how to rearrange equations so reactants and products are on appropriate sides of the arrows.

If equations had to be reversed, reverse the sign of H

If equations had be multiplied to get a correct coefficient, multiply the H by this coefficient since H’s are in

kJ/MOLE (division applies similarly)

Check to ensure that everything cancels out to give you the exact equation you want.

Hint** It is often helpful to begin your work backwards from the answer that you want!

Bond Energies

Energy must be added/absorbed to BREAK bonds (endothermic). Energy is released when bonds are FORMED

(exothermic).

H = sum of the energies required to break old bonds (positive signs) plus the sum of the energies released in the

formation of new bonds (negative signs).

H = bonds broken – bonds formed

H = + reaction is endothermic

H = - reaction is exothermic (favored – nature tends towards lower energy)

ENTHALPY (H)

heat content (exothermic reactions are generally favored)

ENTROPY (S)

disorder of a system (more disorder is favored) Nature tends toward

chaos! Think about your room at the end of the week! Your mom will love this law.

S = - H

T

S = + MORE DISORDER (FAVORED CONDITION)

S = - MORE ORDER

Calculating Entropy

Srxn = S (products) - S (reactants)

FREE ENERGY

G = H - TS

G = G + RT ln (Q)

G = free energy not at standard conditions

G = free energy at standard conditions

R = universal gas constant 8.3145 J/molK

T = temp. in Kelvin

ln = natural log

Q = reaction quotient: Q = [products]

[reactants]

“RatLink”: G = -RTlnK

Grxn = G (products) - G (reactants)

SUMMARY OF FREE ENERGY:

G = + NOT SPONTANEOUS

G = - SPONTANEOUS

Conditions of G:

H S Result

negative positive spontaneous at all temperatures

positive positive spontaneous at high temperatures

negative negative spontaneous at low temperatures

positive negative not spontaneous, ever

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17 Electrochemistry

OIL RIG – oxidation is loss, reduction is gain (of electrons)

Oxidation – the loss of electrons, increase in charge

Reduction – the gain of electrons, reduction of charge

Oxidation number – the assigned charge on an atom

Oxidizing agent (OA) – the species that is reduced and thus causes oxidation

Reducing agent (RA) – the species that is oxidized and thus causes reduction

GALVANIC CELLS

Parts of the voltaic or galvanic cell:

Anode--the electrode where oxidation occurs. After a period of time, the anode may appear to become smaller as

it falls into solution.

Cathode-- the electrode where reduction occurs. After a period of time it may appear larger, due to ions from

solution plating onto it.

inert electrodes—used when a gas is involved OR ion to ion involved such as Fe3+

being reduced to Fe2+

rather

than Fe0. Made of Pt or graphite.

Salt bridge -- a device used to maintain electrical neutrality in a galvanic cell. This may be filled with agar which

contains a neutral salt or it may be replaced with a porous cup.

Electron flow -- always from anode to cathode. (through the wire)

Standard cell notation (line notation) -

anode/solution// cathode solution/ cathode Ex. Zn/Zn2+

(1.0 M) // Cu2+

(1.0M) / Cu

Voltmeter

measures the cell potential (emf) . Usually is measured in volts.

cell potential

Ecell, Emf, or cell—it is a measure of the electromotive force or the “pull” of the electrons as they travel from the

anode to the cathode [more on that later!]

volt (V)

The unit of electrical potential; equal to 1 joule of work per coulomb of charge transferred

voltmeter

measures electrical potential; some energy is lost as heat [resistance] which keeps the voltmeter reading a tad

lower than the actual or calculated voltage. Digital voltmeters have less resistance. If you want to get picky and

eliminate the error introduced by resistance, you attach a variable-external-power source called a potentiometer.

Adjust it so that zero current flows—the accurate voltage is then equal in magnitude but opposite in sign to the

reading on the potentiometer.

STANDARD REDUCTION POTENTIALS

Each half-reaction has a cell potential

Each potential is measured against a standard which is the standard hydrogen electrode [consists of a piece of inert

Platinum that is bathed by hydrogen gas at 1 atm]. The hydrogen electrode is assigned a value of ZERO volts.

standard conditions—1 atm for gases, 1.0M for solutions and 25C for all (298 K)

naught, Ecell, Emf, or cell become Ecello , Emf

o , or cell

o when measurements are taken at standard conditions.

Calculating Standard Cell Potential

Decide which element is oxidized or reduced using the table of reduction potentials. Remember: THE MORE

POSITIVE REDUCTION POTENITAL GETS TO BE REDUCED.

Write both equations AS IS from the chart with their voltages.

Reverse the equation that will be oxidized and change the sign of the voltage [this is now Eoxidation]

Balance the two half reactions **do not multiply voltage values**

Add the two half reactions and the voltages together.

Ecell = Eoxidation + Ereduction means standard conditions: 1atm, 1M, 25C

AN OX – oxidation occurs at the anode (may show mass decrease)

RED CAT – reduction occurs at the cathode (may show mass increase)

FAT CAT – The electrons in a voltaic or galvanic cell ALWAYS flow From the Anode To the CAThode

Ca+hode – the cathode is + in galvanic cells

Salt Bridge – bridge between cells whose purpose is to provide ions to balance the charge. Usually made of a salt

filled agar (KNO3) or a porous cup.

EPA--in an electrolytic cell, there is a positive anode.

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CELL POTENTIAL, ELECTRICAL WORK & FREE ENERGY

V = work (J)/charge (C)

The work that can be accomplished when electrons are transferred through a wire depends on the “push” or emf which is

IF work flows OUT it is assigned a MINUS sign

When a cell produces a current, the cell potential is positive and the current can be used to do work THEREFORE and

work have opposite signs!

= - w

q therefore -w = q

faraday(F)—the charge on one MOLE of electrons = 96,485 coulombs

q = # moles of electrons x F

For a process carried out at constant temperature and pressure, wmax [neglecting the very small amount of energy

that is lost as friction or heat] is equal to G, therefore….

ΔGo = -nFE

o

G = Gibb’s free energy.

n = number of moles of electrons.

F = Faraday constant 9.6485309 x 104 J/V mol

-Eo implies nonspontaneous.

+Eo implies spontaneous (would be a good battery!)

Strongest Oxidizers are weakest reducers.

As Eo reducing strength .

As Eo oxidizing strength .

DEPENDENCE OF CELL POTENTIAL ON CONCENTRATION

Voltaic cells at NONstandard conditions: LeChatlier’s principle can be applied. An increase in the concentration of a

reactant will favor the forward reaction and the cell potential will increase. The converse is also true!

0.0592

NERNST EQUATION: E = Eo - ---------- log Q @ 25C (298K)

n

As E declines with reactants converting to products, E eventually reaches zero.

Zero potential means reaction is at equilibrium [dead battery]. Also, Q =K AND G = 0 as well.