CHEM 162 - General Chemistry II

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CHEM 162 Exam 1 Review – Winter2015 page 1 of 4

CHEM 162: Exam 1 Study Guide Chapter 15: Chemical Equilibrium equilibrium: state where the forward and reverse reactions or processes occur at the same rate – Know that concentrations are not changing at equilibrium, but they need not be equal to one another. – Be able to indicate when equilibrium is achieved given concentration vs. time plots Use the law of mass action to write equilibrium expressions for Kc or Kp for homogeneous and heterogeneous reactions. – Include only gases; omit pure liquids and solids. For general reaction, j A + k B l C + m D

[B] [A] [D] [C]

cK = or k j

m l

DC p

P P P P

K =

Extent of reaction – For large values of Kc or Kp (>10

3), the reaction essentially goes to completion. → The equilibrium mixture consists mostly of products (product favored) → The equilibrium lies to the right. – For small values of Kc or Kp (<10

-‐3), the reaction does not occur to any significant degree. → The equilibrium mixture consists mostly of reactants (reactant favored) → The equilibrium lies to the left. – For intermediate values (10-‐3 < Kc or Kp < 10

3), the equilibrium mixture contains appreciable amounts of both reactants and products. → For Kc or Kp > 1, equilibrium lies to the right. Equilibrium positions: Set of equilibrium concentrations or partial pressures of reactants and products for a system. – While Kc or Kp for a reaction are constant for a given temperature, various equilibrium positions are possible for that reaction depending on the initial concentrations of reactants and products. Know Kc or Kp are unitless!

Relating Kp and Kc: Kp=Kc(RT)Δ n

– Be able to solve for Δn for any homogeneous or heterogeneous reaction. – Know that Kp=Kc for Δn=0. Reaction Quotient (Q): instant state of system, not necessarily at equilibrium. – Also determined using law of mass action. – Q < K: too many reactants → System shifts right to make more products. – Q > K: too many products → System shifts left to make more reactants. – Q = K: system at equilibrium – To compare Q and K, Q must be determined in terms of partial pressure for Kp and in concentrations for Kc. Determination of Kp or Kc – Be able to solve for either using a variety of

experimental data. – Given equilibrium concentrations or partial P’s. – Given initial & changes in conc. or partial P’s. – Given total pressure at equilibrium. Equilibrium Problems Solving for Kp or Kc. 1. Get balanced chemical equation 2. Write equilibrium expression 3. Set up equilibrium ICE table. – Let x=change in conc. or partial pressure. 4. Substitute equilibrium conc. or pressures into equilibrium expression for Kp or Kc. 5. Solve for x, using quadratic method if necessary. 6. Substitute value for x into equilibrium conc. or pressures to solve for Kp or Kc. Le Chatelier's Principle – A system at equilibrium will shift (if possible) to

minimize any stress (change in concentration, pressure, volume, or temperature)

– Predict shifts in equilibrium given specific changes. – Recognize that adding an inert gas (e.g. He, Ne, etc.)

to a gaseous system has no effect on the equilibrium. – Only changes in temperature (T) affect Kp or Kc. – For endothermic reactions, Kp and Kc ↑ as T ↑. – For exothermic reactions, Kp and Kc ↑ as T ↓.

CHEM 162 Exam 1 Review – Winter2015 page 2 of 4

Chapter 15: Chemical Equilibrium (Continued) Characteristics of the Equilibrium Expression – For reverse reaction, equilibrium expression is reciprocal of that for forward reaction – Multiplying coefficients by factor, n, raises equilibrium constant to nth power – Multiple Equilibria – When a reaction can be expressed as a sum of two or more reactions, the equilibrium constant for the overall

reaction is simply the product of the equilibrium constants for the individual reactions. – Be able to manipulate a series of reactions to determine the equilibrium constant for an overall reaction. Chapter 16: Acids And Bases • Know the properties of acids and bases.

• Know Arrhenius definitions for acids and bases. – Know the general form of an Arrhenius acid-‐ base neutralization reaction: – acid + base → water + salt • Know terms monoprotic, polyprotic, etc. • Know Brønsted-‐Lowry definitions for acids and bases. – Recognize conjugate acid-‐base pairs.

• Recognize hydronium ion, H3O + = H+ + H2O

• Know the strong acids: HCl, HBr, HI, HNO3, HClO3, HClO4, H2SO4.

• Know the common strong bases: LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2.

• Recognize strong acids and strong bases dissociate or ionize (break up) completely. → Equilibrium lies far to the right.

HNO3(aq) + H2O(l) → H3O +(aq) + NO3

-‐(aq)

Ca(OH)2(aq) → Ca +2(aq) + 2 OH-‐(aq)

and for H2SO4: H2SO4(aq) + H2O(l) → H3O

+(aq) + HSO4 −(aq)

• Recognize weak acids dissociate or ionize (break up) only to a small degree. → Equilibrium lies far to the left. HF(aq) + H2O(l) H3O

+(aq) + F-‐(aq)

• Write balanced equations and equilibrium expressions for the dissociation of any acid. – Omit pure liquids and solids.

HA(aq) + H2O(l) H3O +(aq) + A−(aq)

Ka = [H

3 O+]*[A,]

[HA]

• Recognize the strength of an acid is inversely related to the strength of its conjugate base. → Strong acids have conjugate bases that are weaker than H2O. → Weak acids have conjugate bases that are stronger than H2O. → Be able to write the net ionic equation for the conjugate base reacting with H2O to form the conjugate acid and OH−. • Know Lewis definitions for acids and bases. – Be able to identify the Lewis acid or base in a given reaction. – Know why nitrogen-‐containing compounds are generally Lewis bases. – Know highly charged cations (Al3+, Cu2+, etc.) can act as Lewis acids and why. • Recognize how structure, bond strength, and

bond polarity influence the properties of an acid. – Be able to explain: – Why HF is weak while HCl, HBr, HI are strong – Why CH4 does not display acidic properties – The ranking of ternary oxyacids (e.g. HClO3

vs HClO2 vs HBrO2, etc.) in terms of strength – Ka values for different acids, including ternary oxyacids, based on structure

CHEM 162 Exam 1 Review – Winter2015 page 3 of 4

Chapter 16: Acids And Bases (Continued) • Water is amphoteric; it can act as an acid or a base.

• Autoionization of water:

2 H2O(l) H3O +(aq) + OH−(aq)

Kw = [H3O +][OH–]=1.0x10–14

Kw = water’s ion-‐product or dissociation constant – When [H3O

+] = [OH–], the solution is neutral. – When [H3O

+] > [OH–], the solution is acidic. – When [H3O

+] < [OH–], the solution is basic. – Be able to solve for [H3O

+] or [OH–] given the other then classify the substance as acidic, basic, neutral. pH scale: pH = 7: neutral and pH < 7: acidic and pH > 7: basic (or alkaline) [H3O

+] = 10−pH [OH-‐] = 10−pOH pH = −log [H3O

+] pOH = −log [OH−] pH + pOH = 14.00 Know # of sig figs in [H3O

+] or [OH−] determines the # of decimal places in pH and pOH and vice versa. Calculate pH of Strong Acids and Strong Bases – Strong acids and bases ionize completely. – Accounting for all the H3O

+ and OH– ions, → [H3O

+] = original molarity of strong acid → [OH–] = molarity of base × (# of OH– in base) pH Calculations for Weak Acids – Weak acids remain mostly undissociated. – Write the equation for the dissociation of an acid. – Set up ICE table with [HA] given or calculated. – Calculate pH of weak acid solution. – Use percent ionization of solution to get Ka. – Use the approximation that

[HX] – x ≈ [HX] for [HA] x

<5%

– Use quadratic method or the method of successive

approximations for: [HA] x

≥5%

Percent ionization = [H3O

+]&at&equilibrium

&[HA] ×100%

Convert Ka and Kb: Kw = Ka ·∙ Kb = 1.0×10

–14 pH Calculations for Weak Bases – Write equation for reaction of base with H2O. – Write equilibrium expression for the weak base. – The weaker the base, smaller the Kb. – Set up ICE table with [A–] given or calculated. – Solve for x to calculate [OH–] and pOH of weak base solution. Acid-‐Base Properties of Salts – salt = ionic compound – Soluble salts dissociate into ions in water. – Classify a given salt as acidic, basic, neutral. Ions that produce acidic solutions – NH4

+ – Highly charged metal ions: Al3+, Zn2+, etc., except cations of strong bases Ions that produce basic solutions – anions that are conjugate bases of weak acids

o all anions except anions of strong acids – Note: SO4

2-‐ + H2O ↔ HSO4 − + OH−

Ions that produce neutral solutions – do not react with H2O to make H

+ or OH− Salts with an acidic cation and basic anion – Classify a given salt as acidic, basic, neutral. – If Ka>Kb, salt is acidic. – If Kb>Ka, salt is basic. – Account for different ion concentrations by multiplying initial concentration with Ka or Kb. pH Calculations for Weak Bases – Write equation for reaction of base with H2O. – Write equilibrium expression for the weak base. – The weaker the base, smaller the Kb. – Set up ICE table with [A–] given or calculated. – Solve for x to calculate [OH–] and pOH of a weak base solution.

CHEM 162 Exam 1 Review – Winter2015 page 4 of 4

Chapter 17: Acids And Bases (Continued) Polyprotic Acids – Write the stepwise dissociation of any polyprotic acid. – Recognize the [H+] is determined only by 1st dissociation step for most polyprotic acids. – Solve for pH and equilibrium ion concentrations.

The Common Ion Effect – The shift in equilibrium caused by the addition of a salt with an ion in common with the dissolved substances – Given a compound added to an acid or base at equilibrium, predict equilibrium shifts (to the left or right), if [H+] ↑ or ↓, if [OH–] ↑ or ↓, and if pH ↑ or ↓, – Calculate the pH of an acid or base solution when a common ion is added.

Be prepared to solve problems combining concepts from Chapters 15 and 16, as well as

other concepts and calculations from CHEM 161 (e.g. Gas Laws).

You will be given a Periodic Table with atomic numbers, atomic masses, Solubility Rules, the Activity Series, and constants and conversions.