Untitled

CHEM 124: Review For Exam

Chapter 14 (Pages 8-9)

Review Key Terms

Spontaneous thermodynamic processes
Nonspontaneous thermodynamic processes
Entropy
Enthalpy
Free energy (or Gibbs energy)
Standard free (Gibbs) energy of formation
Standard free (Gibbs) energy of reaction
Reaction quotient, Q
Equilibrium constant, K
Nonstandard free energy of reaction
Vapor pressure

Main Points and Concepts

Spontaneous Processes
- Chemical systems always proceed spontaneously towards zero nonstandard free (Gibbs) energy, corresponding to the reaction equilibrium.
Equilibrium Constant Relation
- The relationship between the equilibrium constant and the standard free energy change of the system can be expressed as:
  $\Delta G^{\circ} = -RT \ln(K)$
  where R is the universal gas constant and T is the temperature in Kelvin.
Sign of ΔG and K Relationship
- If $\Delta G^{\circ} > 0$ and $K < 1$ , the reaction favors the reactants.
- Conversely, if $\Delta G^{\circ} < 0$ and $K > 1$ , the reaction favors the products.
Relation of ΔG, Q, and K
- Understand how the sign of $\Delta G$ relates to the relative values of Q and K.
- Cases:
 - If $Q > K$ , then $\Delta G > 0$ (non-spontaneous process).
 - If $Q < K$ , then $\Delta G < 0$ (spontaneous process).
Calculating ΔG under Nonstandard Conditions
- The nonstandard free energy change can be calculated using:
  $\Delta G = \Delta G^{\circ} + RT \ln(Q)$
- Notes: Requires ability to write Q expressions and calculate them based on concentrations or partial pressures.
Van’t Hoff’s Equation
- Important for calculating changes in the equilibrium constant with temperature:
 $\ln(K) = - \frac{\Delta H{rxn}^{\circ}}{R} \cdot \frac{1}{T} + \frac{\Delta S{rxn}^{\circ}}{R}$
 or:
 $\ln(K) = - \frac{\Delta H{rxn}^{\circ}}{R} (\frac{1}{T1} - \frac{1}{T2}) + \frac{\Delta S{rxn}^{\circ}}{R} \ln(K1)$ where $\Delta H{rxn}^{\circ}$ is standard enthalpy change, $\Delta S_{rxn}^{\circ}$ is standard entropy change.

Chapter 15 (Pages 1-8)

Review Key Terms

Amphiprotic
Arrhenius acids and bases
Bronsted-Lowry acids and bases
Proton donor
Proton acceptor
Conjugate acid-base pairs
Acid-base reaction
Strong acids
Strong bases
Autoionization
Ion-product constant for water
Hydronium ion, H₃O⁺
Hydroxide ion, OH⁻
Kw
pH
pOH
pKw
Acid ionization constant (Kₐ)
Base ionization constant (K𝑏)
pKₐ
pK𝑏
Weak acid
Weak base
Oxoacid
Binary acid
Monoprotic acid
Diprotic acid
Triprotic acid
Polyprotic acid
% ionization of an acid

Main Points and Concepts

Acid and Base Recognition
- Recognize compounds as either acids or bases from both the Arrhenius and Bronsted-Lowry definitions.
- Bronsted-Lowry Theory: Acids donate H+ ions, while bases accept H+ ions.
Acid-Base Reaction
- In any acid-base reaction, a transfer of H+ occurs from the acid on the reactant side to the base on the reactant side.
- There are always two conjugate acid-base pairs in any acid-base reaction; be able to identify them.
Predict Products and Direction of Reaction
- From the relative strengths of acids or bases, predict the expected products of an acid-base reaction and the direction it will proceed (K > 1 or < 1).
pH Calculations
- Convert between $[H₃O⁺]$ , $[OH⁻]$ , pH, and pOH using: $Kw = 10^{-14}$ at 25°C or $pKw = 14$ at 25°C.
- Calculate these quantities for any strong acid or strong base in solution.
Common Strong Acids and Bases
- Strong acids include three hydrohalic acids: HCl, HBr, HI, and three oxoacids: HNO₃, HClO₄, H₂SO₄. Note: Oxoacids are strong when the number of oxygen atoms exceeds the number of acidic hydrogen atoms by two or more.
- Strong bases are group I metal hydroxides and Ba(OH)₂, Sr(OH)₂, Ca(OH)₂ from group II metal hydroxides.
- Strong acids and bases are completely ionized in water, whereas weak acids and bases are typically less than 5% ionized in water.
Electron Pairs in Bronsted-Lowry Theory
- A lone pair(s) of electrons is necessary for a molecule to act as a base, according to the Bronsted-Lowry definition.
Pure Water at 25°C
- In pure water, $[H₃O⁺] = [OH⁻] = 10^{-7}$ , thus pH = 7.00.
- Adding acid to pure water increases $[H₃O⁺]$ and decreases $[OH⁻]$ based on the autoionization equilibrium of water (Le Chatelier’s principle).
 - Outcome: $[H₃O⁺] > [OH⁻]$ results in pH < 7 and pOH > 7.
- In contrast, adding a base decreases $[H₃O⁺]$ and increases $[OH⁻]$ : $[H₃O⁺] < [OH⁻]$ results in pH > 7 and pOH < 7.
Weak Acid or Base Equilibria
- Be able to write and identify the products of weak acid or weak base equilibria reactions in water, and set up and solve weak acid or weak base equilibrium problems to determine equilibrium concentrations, pH, or pOH, given equilibrium constant information.
Relationship of Acid Strength to Conjugate Base Strength
- Qualitatively and quantitatively recognize how the strength of the acid correlates with its conjugate base strength (toward neutralization). Quantitatively, the relation is expressed as: $Kw = Ka \times K_b$
 for a conjugate acid-base pair.
Percent Ionization
- Understand how % ionization changes with initial weak acid concentration and learn to calculate it based on equilibrium constant information. Also, know to calculate the equilibrium constant from percent ionization and initial concentration.
Polyprotic Acids
- Recognize and write the stepwise acid-ionization equilibrium reactions for polyprotic acids, which can be related to: $K{a1} > K{a2} > K{a3}$ .
- Qualitatively relate molecular structure to the relative acid strength in binary acids or oxoacids.
Salt Solutions
- Determine if a salt solution is neutral, acidic, or basic based on the cations and anions of the salt along with their respective $Ka$ or $Kb$ .
- Calculate the pH of salt solutions according to the type of cation and anion present:
 - (1) Anion of weak acid and cation of strong base yields pH > 7.
 - (2) Cation of strong acid and anion of weak base yields pH < 7.
- Note no pH calculations for salts of weak acid-weak base or salts of polyprotic acids.

Chapter 16 (Pages 2-3)

Review Key Terms

Buffer solution
Buffer range
Buffer capacity

Main Points and Concepts

Recognition of Buffer Systems
- Identify a buffer system based on its components, which typically consist of a weak acid and its conjugate base or a weak base and its conjugate acid.
Resistance to pH Changes
- Understand how buffer solutions are resistant to pH changes when relatively small amounts of either a strong acid or a strong base are added.
Conjugate Acid-Base Pair Selection
- Identify the best conjugate acid-base pair for a buffer to achieve a desired pH based on their $K_a$ value.
Calculating Ratios for Specific pH
- Calculate the ratio of the amounts (mol or M) of conjugate base to weak acid needed to maintain a specific pH.
Buffer Capacity
- Determine buffer solutions with higher capacity; the effectiveness of a buffer typically is greatest at $pH = pK_a \pm 1$ .
Chemical Equations Writing
- Write the chemical equations for neutralization reactions of any acid or base interacting with a buffer.
- Write the chemical equations for the water dissociation equilibrium reaction by any weak acid and weak base, including conjugate bases and acids as per the Bronsted-Lowry theory as they relate to buffer solutions and titrations.
Henderson-Hasselbalch Equation
- Solve for pH (or equilibrium concentrations or pOH) of buffer solutions using the Henderson-Hasselbalch equation, which assumes the small x approximation and non-zero initial concentrations of the acid and its conjugate base.
pH Changes from Strong Acid/Base Addition
- Calculate the change in pH from a small addition of strong acid or strong base to a buffer solution.

Untitled

CHEM 124: Review For Exam

Chapter 14 (Pages 8-9)

Review Key Terms

Main Points and Concepts

Suggested End of Chapter Questions/Problems

Chapter 15 (Pages 1-8)

Review Key Terms

Main Points and Concepts

Suggested End of Chapter Questions/Problems

Chapter 16 (Pages 2-3)

Review Key Terms

Main Points and Concepts

Suggested End of Chapter Questions/Problems