Chapter 13Properties of SolutionsReview

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Chapter 13Properties of Solutions-Review Intermo
lecular Forces Ion-Ion Ion-Dipole Dipole-Dipole
London Dispersion H-bonds The Solubility
Process Energy Changes Solubility formation,
spontaneity and disorder Saturated Solutions and
Solubility SolventSolute Interactions Ways of
expressing concentrations Mass percentage, ppm
and ppb Mole Fraction, Molarity and Molality
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If same mass and shape, polarity will be the
determining factor If large difference in
molecular weight, dispersion forces will be the
determining factor Important exceptions
H-Bonds
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Relative strengths of bonds CovalentgtIonicgtgtH-bon
dsgtion-dipolegtdipole-dipolegtdispersion
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ppm 1 g/103 kg 1 mg/L for water ppb 1
g/106 kg 1 mg/L for water Mass 100 mass
of component in solution/total mass of
solution Mole fraction Moles of
component/total moles of all components Molarit
y Moles solute/liters of solution Molality
Moles solute/kilograms of solvent
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Chapter 14 Factors that affect reaction
rates Physical state of the reactants Concentrat
ions of the reactants Temperature at which the
reaction occurs Presence of a catalyst Reaction
rates Change of rate with time Reaction rates
and stoichiometry Concentration and
rate Exponents in the rate law Units of rate
constants Initial rates to determine rate
laws Change of concentration with
time First-order reactions Second-order
reactions Half-life Temperature and
rate Orientation factor Activation
energy Arrhenius equation Determining the
activation energy Reaction mechanisms Elementary
reactions Multistep mechanisms Rate-determining
step for a multistep mechanism Slow initial
step Fast initial step Catalysis Homogeneous H
eterogeneous Enzymes
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Chapter 15 Concept of equilibrium Equilibrium
constant In terms of concentration, Kc In terms
of partial pressures, Kp Units Directions,
stoichiometries, Hesss law Heterogeneous
equilibria Calculating equilibrium
constants Predicting the direction of
reaction Reaction quotient, Q Calculating
equilibrium concentrations Le Châteliers
principle Change in reactant or product
concentrations Change in volume or
pressure Change in temperature Catalysts
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If K gtgt 1, then equilibrium lies to the right and
products predominate If K ltlt 1, then equilibrium
lies to the left and reactants predominate
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The equilibrium constant of a reaction in the
reverse direction is the inverse of the
equilibrium constant of the reaction in the
forward direction. The equilibrium constant of
a reaction that has been multiplied by a number
is the equilibrium constant raised to a power
equal to that number. The equilibrium constant
for a net reaction made up of two or more steps
is the product of the equilibrium constants for
the individual steps.
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Reaction quotient, Q A number obtained by
substituting starting reactant and product
concentrations or partial pressures into an
equilibrium-constant expression. If Q K,
then the system is at equilibrium If Q gt K,
then the reaction will go to the left If Q lt
K, then the reaction will go to the right If
use concentrations, Q will be written as Qc If
use partial pressures, Q will be written as Qp
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Chapter 16 Acid-Base Equilibria Acid-base
definitions Arrhenius Brønsted-Lewis Conjugate
acid-base pairs Relative strengths of acids and
bases Autoionization of water pH scale Strong
acids and bases Weak acids Calculating Ka from
pH Using Ka to calculate pH Polyprotic
acids Weak bases Types of weak
bases Relationship between Ka and Kb Acid-base
properties of salt solutions Anions with
water Cations with water Combined effect of
cations and anions in solution Acid-base behavior
and chemical structure Binary acids Oxyacids Ca
rboxylic acids Lewis acids and bases Hydrolysis
of metals
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Arhhenius Acids are substances
that, when dissolved in water, increase the
concentration of H ions Bases are
substances that, when dissolved in water,
increase the concentration of OH-
ions Brønsted-Lowry An acid is a
substance that can donate a proton to another
substance A base is a substance that can
accept a proton
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One more definition of acids and bases
A Lewis acid is an electron-pair acceptor A
Lewis base is an electron-pair donor
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Rules on non-metal oxyacids For oxyacids that
have the same number of OH groups and the same
number of O atoms, acid strength increases with
increasing electronegativity of the central
atom. For oxyacids that have the same central
atom, acid strength increases as the number of
oxygen atoms attached to the central atom
increases. (Acidity increases as the oxidation
of the central ion increases)
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Chapter 17 Common ion effect Buffered
solutions Composition and actions of
buffers Calculating the pH of a buffer Buffer
capacity and pH range Addition of strong acids
or bases to buffers Acid-base titrations Strong
acid-strong base Weak acid-strong
base Titrations of polyprotic acids Solubility
equilibria Solubility-product constant,
Ksp Solubility and Ksp Factors that affect
solubility Common ion effect Solubility and
pH Formation of complex ions Amphoterism Precipi
tation and separation of ions Selective
precipitation of ions Qualitative analysis for
metallic elements
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Chapter 19Thermodynamics Spontaneous
Processes Criterion for spontaneity Reversible
and irreversible processes Entropy and the second
law of thermodynamics Entropy change DS for
phase changes The second law of
thermodynamics The molecular interpretation of
entropy Molecular motions and energy Boltzmanns
equation and microstates Making qualitative
predictions about DS The third law of
thermodynamics Entropy changes in chemical
reactions Entropy changes in the
surroundings Gibbs free energy Standard
free-energy changes Free energy and
temperature Free energy and the equilibrium
constant
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Reversible process ?Suniv ?Ssystem
?Ssurroundings 0 Irreversible process
?Suniv ?Ssystem ?Ssurroundings gt 0 Second
Law of Thermodynamics The total entropy of the
universe increases in any spontaneous process
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Entropy is a measure of how many microstates are
associated with a particular macroscopic state S
k ln (W). Entropy increases with the number
of microstates of the system The number of
microstates available to a system increases with
an increase in volume, an increase in
temperature, or an increase in the number of
molecules because any of these changes increases
the possible positions and energies of the
molecules of the system.
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The third law of thermodynamics The entropy of a
pure crystalline substance at absolute zero is
zero S( 0 K) 0
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?G equals the maximum useful work that can be
done by the system on its surroundings in a
spontaneous process occurring at constant
temperature and pressure ?G -wmax
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Chapter 20 Oxidation states Balancing
oxidation-reduction equations Balancing
equations by the method of half-reactions
Balancing equations in acid solution Balancing
equations in basic solution Voltaic cells A
molecular view of electrode processes Cell EMF
under standard conditions Standard reduction
(half-cell) potentials Strength of oxidizing and
reducing agents Free energy and redox
reactions EMF and ?G Cell EMF under nonstandard
conditions The Nernst equation Concentration
cells Batteries and fuel cells Lead-acid
battery Alkaline battery Nickel-cadmium,
nickel-metal-hydride, and lithium-ion
batteries Hydrogen fuel cells Direct methanol
fuel cells Electrolysis Quantitative aspects of
electrolysis Electrical work
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• Rules for balancing redox equations in acids by
  the method of half-reactions (pp 850-851)
• Assign oxidation states to see which atoms are
  gaining electrons and which are losing electrons
• Divide the equation into two half-reactions, one
  for oxidation and the other for reduction
• Balance each half-reaction
• a. First, balance the elements other than H
  and O
• b. Next, balance the O atoms by adding H2O as
  needed
• c. Then, balance the H atoms by adding H as
  needed
• d. Finally, balance the charge by adding e- as
  needed
• e. Check that the number of e- match the
  changes in oxidation states
• Multiply the half-reactions by integers if
  necessary so that the number of electrons lost in
  one half-reaction equals the number gained in the
  other
• Add the two half-reactions and whenever possible
  simplify by canceling species appearing on both
  sides of the combined equation
• Check to make sure that the number of atoms of
  each kind on the left side of the equation is the
  same as the number on the right side

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Anode and oxidation both start with
vowels Cathode and reduction both start with
consonants Anions are attracted to the
anode Cations are attracted to the cathode
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One coulomb is the amount of electric charge
transported in one second by a steady current of
one ampere.
1 Coulomb 1 ampere x 1 second
In general, Coulombs amperes x seconds
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Chapter 21Nuclear Chemistry Radioactivity Nucle
ar equations Types of radioactive
decay Patterns of nuclear stability Neutron-to-p
roton ratio Radioactive series Nuclear
Transmutations Accelerating charged
particles Using neutrons Transuranium
elements Rates of radioactive decay Dating Calc
ulations based on half-life Detection of
radioactivityradiotracers Energy changes in
nuclear reactions--Nuclear binding energies
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Nucleon A particle found in the nucleus of an
atom Nuclide nucleus with a specified number
of protons and neutrons Radionuclides nuclides
that are radioactive Radioisotopes atoms
containing radionuclides
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Generalizations Nuclei above the belt tend to
emit beta particles, thus lowering the number of
neutrons and increasing the number of protons
thus moving toward the belt. Nuclei below the
belt can undergo either positron emission or
electron capture. Either increases the number of
neutrons and decreases the number of
protons. Nuclei with atomic numbers gt83 tend to
decay by alpha emission. Two protons and two
neutrons are lost bringing it closer to the belt.

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Chapter 23Metals Metallic bonding Physical
properties of metals Electron-sea model for
metallic bonding Molecular-orbital model for
metals \ Alloys--Intermetallic compounds Transiti
on metals Physical properties Electron
configurations and oxidation states Magnetism Ch
emistry of selected transition metals Chromium I
ron Copper
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Chapter 24Chemistry of Coordination
Compounds Metal Complexes Werners theory The
metal-ligand bond Charges, coordination numbers,
and geometries Ligands with more than one donor
atom Metals and chelates in living
systems Nomenclature of coordination
chemistry Isomerism Structural
isomerism Stereo isomerism