Chemistry Test 3

Solving Equilibrium Problems

  1. Balance the equation
  2. Write the equilibrium expression
  3. List the initial concentratsions
  4. Calculate Q and determine the shift to equilibrium
  5. Define equilibrium concentrations
  6. Substitute equilibrium concentrations into equilibrium expression and solve
  7. Check calculated concentrations by calculating K

Quadratic and Simplification
5% test: if a simplification like this changes the area of simplification by less than 5% then the simplification is valid. If the change is 5% or greater use the quadratic equation.
Le Chatelier’s Principle
when a system at equilibrium is subjected to a change in temperature, pressure or concentration of a reacting species, the system responds by attaining a new equilibrium that offsets the impact of the change.
Effects of Changes on the System

Concentration: the system will shift away from the added component.


Temperature: K will change depending upon the temperature (treat the enery change as a reactant).


Pressure: a) adding of inert gas does ont affect the equilibrium position. b) decreasing the volume shifts the equilibrium toward the side with fewer moles.

Arrhenius acid


Arrhenius base


Bronsted-Lowry acid


Bronsted-Lowry base

produces H+ in aqueous solutions.


produces OH- in aqueous solutions.


H+ donor


H+ acceptor



Strong acid

Ionization equilibrium lies far to the right (products).

Yields a weak conugate base

K = large

Completely breaks apart in water

Weak acid

Ionization equilibrium lies far to the left (reactants).

Weaker the acid, the stronger its conjugate base.

Does not break apart much.

Strong Acids

HCl hydrocloric acid

HBr hydrobromic acid


HClO4 perchloric acid

HNO3 nitric acid

H2SO4 sulfuric acid

Strong Bases









Self Ionization of Water

has a reaction with itself


Amphoteric: can behave as either an acid or base


H2O + H2O –> H3O+ + OH


Kw = 1.0 x 10-14






H+ > OH–      pH < 7


OH > H+     pH > 7


H+ = OH–       pH = 7


Larger Ka value = stronger acid

Smaller pKa value = stronger acid


will have H+ bound to F N or O atoms that can be donated

pH Scale

pH = -log[H+]


pH decreases as H+ increases


The acid(base) dominates pH (pOH) if:

  • its concentration is greater than 10-6M
  • (pure water would donate 10-7M H+)
  • its Ka(Kb) is larger than Kw

Percent Ionization aka Percent Dissociation

[H3O+]/[HA] x100%


more concentrated   more diluted

<—acid conc —–

—-percent dissociation–>

<—[H+] conc——-

pKw = pH + pOH = 14

Product of the ionization constants of an acid conjugate base pair is the ionization constant of water


Ka x Kb = Kw


pKa + pKb = pKw


Larger Kb = Stronger base

Smaller pKb = Stronger base


will have N or O atoms that have lone pair electrons that can attract H+


pOH = -log [OH-]


pOH > 7 acidic

pOH < 7 basic

pOH = 7 neutral

Polyprotic Acid

an acid that contatins more than one ionizable H atom per molecule


the first proton is the easiest to remove (largest Ka)


After that there is an anion that the proton would be removed from.


It gets progressively more difficult to removed succeeding protons

Acid-Base Properties of Salts

cations may act as acids in water

except: Li+, Na+, K+, Rb+, Cs+, Ca2+, Sr2+

(strong bases with hydroxide removed)


these are the conjugate acids of strong bases and are such super weak acids that they will not act as an acid in water


these are pH neutral

Acid-Base Properties of Salts

Anions may act as bases in water

Except: Cl-, Br-, I-, NO3-, HSO4-, ClO4-, BrO4-, IO4

(strong acids without H+)


these are the conjugate bases of strong acids and are such super weak bases that they will not act as a base in water


these are pH neutral

The Effect of Structure on Acid Base Properties

Factors effecting acid strength:

bond polarity (high is good)

bond strength (low is good)


  • Contains the group H-O-X (x=everything else)
  • for a given series the acid strength increases with an increase in the number of oxygen atoms attached to the central atom.
  • The greater the ability of X to draw electrons toward itself, the greater the acidity of the molecule.
  • the greater the electron density, the easier the proton will leave. The stronger the acid.

Molecules containing the grouping H-O-X can behave as acids


  • the acid strength increases with increasing electron-withdrawing ability of X. This enables the H+ to be released. Ex. HClO4
  • If X has a very low electronegativity, the OH- can be lost and the solution will be basic Ex. NaOH

Lewis Acid


Lewis Base

electron pair acceptor


electron pair donor

Common Ion Effect


The suppression of the ionization of a weak electrolyte caused by the addition of an ion that is also a product of the ionization equilibrium of the weak electrolyte.


Shift in equilibrium position that occurs because of the addition of an ion already involved in the equilibrium reaction


an application of Le Chatelier’s principle


the strong acid causes less of the weak acid to dissociate than would be expected

Buffered Solutions

resist a change in pH


they are weak acids or bases containing a common ion


after addition of strong acid or base, deal with stoichiometry first, then the equilibrium

Buffer Solutions
a solution with appreciable amounts of a weak acid and its conjugate base that resist change in pH upon addition of an acid or base
Henderson-Hasslebalch  Equation for BUFFERS
pH = pKa = log nbase/nacid
Characteristics of Buffered Solutions

  • buffers contian relatively large amounts of weak acid and corresponding conjugate base
  • added H+ reacts to completion with the conjugate base
  • Added OH- reacts to completion with the weak acid
  • The pH in the buffered solution is deteremined by the ratio of the concentrations of the weak acid and weak base. As long as this ratio remains virtually constant. This will be the case as long as the concentrations of the buffering materials (HA and A- or B and BH+) are large compared with amounts of H+ or OH- added

Buffering Capacity

the amount of acid or base that a buffer can neutralize before its pH changes appreciable


when the ratio nbase/nacid is close to 1, the buffer has its maximum buffer capacity.

Buffer Region
a weak acid/conjugate base pair acts best as a buffer around the pH region equal to the pKa. Usually within 1 pH unit of the pKa
Titration Curve

Plotting the pH of the solution being analyzed as a function of the amount of titrant added


equivalence (stoichiometric) point- point in the titration when enough titrant has been added to react exactly with the substance in solution being titrated

Acid Base Indicators

a substance whose color depends on the pH of the solution to which it is added.


The indicator is a weak acid. It has a different color than its conjugate base.


The equivalence point is not necessarily the same as the end point (but they are ideally as close as possible)


the color change will become apparent when about 10% of the initial from of the indicator has been converted

how much of a substance that will dissolve in a given amount of solvent at a given temperature
Solubility Product
the equilibrium constant expression for a salt dissolving in water
ksp values

increase the number, more it breaks apart


decrease the number, less it breaks apart


lower ksp more solid that will dissolve


mixing two solutions of ions


Q>Ksp : precipitation occurs and will continue until the concentrations are reduced to the point that they satisfy Ksp: too many products


Q< Ksp : no precipitation occurs: too many reactants, make more products

complex ion
a charged species consistening of a metal ion surrounded by ligands
a lewis base (molecular ion having a lone electron pair that can be donated to an empty orbital on the metal ion to form a covalent bond)
Formation (stability) constant K
Equilibrium constant for each step of the formation of a complex ion by the addition of an individual ligand to a metal ion or complex ion in aqueous solution
two strategies for dissolving a water-insoluble ionic solid

  • if the anion of the solid is a good base, the solubility is greatly increased by acidifying the solution
  • in cases where the anion is not sufficiently basic, the ionic solid often can be dissolved in a solution containing a ligand that forms stable complex ions with its cation.

Spontaneous Process

a change that occurs in a system left to itself; once started no external action is necessary to make the process continue.


exothermic heat given off; combustion


one that occurs without outside intervention

nonspontaneous process
will not occur unless some external action is continuously applied

lets us predict whether a process will occur but gives no information about the amount of time required for the process




thermodynamic property related to the degree of disorder in a system


nature tends toward disorder.


increase if:

  • liquisd are formed from solids
  • gases are formed from either solids or liquids
  • the number of molecules of gas increases as a result of a chemical reaction
  • the temperature of a substance increases

Entropy change, deltaS
the difference in entropy between two states
Positional Entropy

a gas expands into a vacuum because the expanded state has the highest positional probability of states available to the system


S solid < S liquid << S gas


the change in positional entropy is dominated by the relative numbers of molecules of gaseous reactants and products

Second Law of Thermodynamics

in any spontaneous proces there is always an increase in the entropy of the universe


the entropy of the universe is increases


delta Suniv> 0

where delta Suniv = delta Ssys + delta Ssurr


for a spontaneous process


the sign of delta Ssurr depends on the direction of the heat flow.

The magnitude of delta Ssurr depends on the temperature

delta Ssurr = qrev/T = delta H/T


qrev = heat gained in a reversible process -Joules


T = temperature (kelvin)


the change to the system is the reverse of the change to the surroundings

delta Ssurr = delta H/T = -delta H sys/T

Free Energy

delta G = delta H – TdeltaS (from the standpoint of the system)


a process (at constant T, P) is spontaneous in the direction in which free energy decreases


a [-] delta G means [+] deltas Suniv

Gibbs Free Energy, G

delta G negative spontaneous

delga G positive nonspontaneous

delta G zero equilibrium

Third Law of Thermodynamics

  • the entropy of a perfect crystal at 0K is zero
  • because S is explicitly known (=0) at 0K, S values at other temps can be calculated

The more complex the molecule, the larger the entropy

Dependence of Free Energy on Pressure

delta G = delta G zero + RT ln(Q)


R gas law constant 8.3145 J/k*mol

T: temperature in Kelvin

Q: reaction quotient (in partial pressures)

Delta G zero: free energy change at the standard state


Q = K

Free Energy and Work

maximum possible useful work obtainable from a process at constant temperature adn pressure is equal to the change in free energy.


Wmax = deltaG

  • Achieving the maximum work available from a spontaneosu process can occur only via a hypothetical pathway. Any real pathway wastes energy.
  • All real processes are irreversible
  • First law: cant have more than start with
  • Second law: b/c disorder (entropy)



the study of the interchange of chemical and electrical energy
oxidation reduction (redox) reaction

involves a transfer of electrons from the reducing agent to the oxidixing agent.






reducing agent


oxidizing agent

loss of electrons


gain of electrons


electron donor


electron acceptor

balancing by half-reaction method in Acid

Write a separate reduction, oxidation reactions:

For each half reaction:

  • Balance elements except H, O
  • balance O using H2O
  • balance H using H+
  • Balance charge using electrons
  • multiply by interger to equalize electron count
  • add half reactions
  • check that elements and charges are balanaced

Half reaction Method- balancing in base

  • balance as in acid
  • Add OH- that equals H+ ions (both sides)
  • Form water by cominbing H+, OH-
  • Check elements and charges for balance

Galvanic Cells

a device in which chemical energy is changed into electrical energy


this is done with a oxidation-reduction (redox) reaction

Salt bridge
a U-tube filled with an electrolyte that connects the two compartments of a galvanic cell, allowing ion flow without extensive mixing of the diff solutions




the electrode where oxidation occurs


the electrode where reduction occurs

Voltaic Cell: cathode reaction


Voltaic Cell: anode reaction

mass increases, electrode larger


mass decrease, electrode smaller

Cell Potential
or electromotive Force (emf) the Pull or driving force on the electrons
Volt and Volt Meter

the unit of electrical potential defined as one joule of work per coulomb of charge transferred


an instrument that measures cell potential by drawing electrical current through a known resistance

Cell Potential

a galvanic cell consists of an oxidizing agent in one compartment that pulls electrons through a wire from a reducing agent in the other compartment


the pull or driving force on the electrons is called the cell potential (Ecell) or emf of the cell.

unit of electrical potential is the volt V- 1 joule of work per coulomb of charge transferred

Galvanic cell

all half reactions are given as reduction processes in standard tables


when a half reaction is reversed, the sign of Ezero is reversed


when a half reaction is multiplied by an integer, E zero remains the same


a galvanic cell runs spontaneously in the direction that gives a positive value for E zero cell

Line Notation

anode components are listed on the left. Cathode components are listd on the right.


anode and cathode are separated by double vertical lines. a phase difference is indivated by a single vertical line.


work is never the maximum possible if any current is flowing


in any real, spontaneous process some energy is always wasted- the actual work realized is always less than the calculated maximum

maximum cell potential

directly related to the free energy difference between the reactants and the products in the cell


delta G zero = -nFEzero


n=number of moles of electrons

F=Faraday= 96,485 coulombs per mole of electrons

Ion-selective electrodes
an electrode senstive to the concentration of a particular ion in solution
glass electrode
an electrode for measuring pH from the potential difference that develops when it is dipped into an aqueous solution containing H+ ions
a galvanic cell or a group of galvanic cells connected in series, where the potentials of the individual cells added to give the total battery potential
cathodic protection
a method in which an active metal, such as magnesium is connected to steel in order to protect it from corrosion
the process by which metals are oxidized in the atmosphere
process that involves forcing a current through a cell to cause a nonspontaneous chemical reaction to occur

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