Acids, Bases, Buffers, and Thermochemistry: General Chemistry Study Notes

Salts and Their Effect on pH

Classification and Properties of Salts

Salts are ionic compounds formed from the reaction of acids and bases. Their constituent ions can affect the pH of aqueous solutions depending on their origin.

• Cations/Anions: Can be conjugates of acids or bases.

• Ionic Compounds: Composed of metal/nonmetal ions or polyatomic ions.

• Effect on pH: Some ions hydrolyze water, altering solution pH.

Examples of Conjugate Pairs:

• $\text{NH}_4^+/\text{NH}_3$ (conjugate acid/base)

• $\text{HF}/\text{F}^-$ (conjugate acid/base)

• $\text{CH}_3\text{NH}_3^+/\text{CH}_3\text{NH}_2$

Key Equations:

• $K_a \times K_b = K_w$ (for conjugate acid-base pairs)

• $pK_a + pK_b = 14$

Cations: Can act as weak acids if derived from weak bases.

• Example: $\text{NH}_4^+$ can donate a proton to water, forming $\text{NH}_3$ and $\text{H}_3\text{O}^+$.

Anions: Can act as weak bases if derived from weak acids.

• Example: $\text{F}^-$ (from HF) can accept a proton, forming HF.

• Conjugate base of strong acid (e.g., $\text{Cl}^-$ from HCl) is neutral.

Table: Acidic, Basic, and Neutral Salts

Salt	Cation	Anion	Origin	pH Effect
NaCl	Na+	Cl-	SB/SA	Neutral
NaF	Na+	F-	SB/WA	Basic
NH4Br	NH4+	Br-	WB/SA	Acidic
CH3NH3Cl	CH3NH3+	Cl-	WB/SA	Acidic
Ca(CH3COO)2	Ca2+	CH3COO-	SB/WA	Basic

Buffers and Buffer Solutions

Definition and Composition

A buffer is a solution that resists changes in pH upon addition of small amounts of acid or base. Buffers consist of a weak acid and its conjugate base, or a weak base and its conjugate acid.

• Weak acid + conjugate base (e.g., CH3COOH/CH3COO-)

• Weak base + conjugate acid (e.g., NH3/NH4+)

Buffer Capacity: Most effective when $[\text{HA}] \approx [\text{A}^-]$ (ratio close to 1).

Buffer Action and Calculations

• Buffers neutralize added acids/bases via equilibrium reactions.

• Example: $\text{CH}_3\text{COOH} + \text{CH}_3\text{COO}^- \rightleftharpoons \text{CH}_3\text{COO}^- + \text{CH}_3\text{COOH}$

Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation relates the pH of a buffer to the concentrations of acid and base:

• $\text{pH} = \text{p}K_a + \log \left( \frac{[\text{A}^-]}{[\text{HA}]} \right)$

Example Calculation:

• Given $[\text{A}^-] = 0.150$ M, $[\text{HA}] = 0.050$ M, $pK_a = 4.74$

• $\text{pH} = 4.74 + \log \left( \frac{0.150}{0.050} \right) = 5.22$

Buffer Formation

• Mixing a weak acid/base with its conjugate.

• Reacting a weak acid/base with a strong base/acid to generate the conjugate.

Titration and Equivalence Point

Titration of Weak Acid with Strong Base

Titration involves the gradual addition of a titrant to a solution to determine concentration or pH changes. The equivalence point is reached when moles of acid equal moles of base.

• Initial pH: Calculated using weak acid equilibrium.

• After addition of base: Use ICE table and buffer equations.

• At equivalence: All weak acid converted to conjugate base; pH determined by base hydrolysis.

Example:

• Titration of 25 mL 0.1 M HCOOH with 0.1 M NaOH.

• Initial: $\text{HCOOH} + \text{H}_2\text{O} \rightleftharpoons \text{HCOO}^- + \text{H}_3\text{O}^+$

• After addition: $\text{HCOOH} + \text{OH}^- \rightarrow \text{HCOO}^- + \text{H}_2\text{O}$

• At equivalence: $\text{HCOO}^-$ acts as a weak base.

Thermochemistry and Energy Exchange

Introduction to Thermodynamics

Thermodynamics is the study of heat, work, and energy exchanges in chemical processes.

• Kinetic Energy: Due to motion.

• Potential Energy: Due to position or composition.

• Thermal Energy: Associated with temperature.

• Chemical Energy: Associated with chemical bonds.

Units of Energy

• Joule (J): $1 \text{J} = 1 \text{kg} \cdot \text{m}^2 / \text{s}^2$

• Calorie (cal): $1000 \text{cal} = 1 \text{C}$

Energy Exchange: System and Surroundings

• System: Part of universe under study.

• Surroundings: Everything else.

• Types: Open, closed, isolated systems.

First Law of Thermodynamics

The total energy of the universe is constant. Energy cannot be created or destroyed, only transferred.

• $\Delta E_{\text{universe}} = 0 = \Delta E_{\text{system}} + \Delta E_{\text{surroundings}}$

• $\Delta E_{\text{system}} = q + w$

• q: Heat (J)

• w: Work (J)

Work and Heat Calculations

• Work due to volume change: $w = -P_{\text{ext}} \Delta V$

• Heat transfer: $q = m C_s \Delta T$

• Specific Heat Capacity ($C_s$): Amount of heat required to raise temperature of 1 g by 1°C.

Enthalpy ($H$)

Enthalpy is the total internal energy plus pressure-volume work at constant pressure.

• $H = E + PV$

• $\Delta H = \Delta E + P \Delta V$

• At constant pressure: $\Delta H = q$

Calorimetry

Calorimetry measures heat changes in chemical reactions.

• $q_{\text{reaction}} = m C_s \Delta T$

• Example: Dissolving MgSO4 in HCl and measuring temperature change.

Standard Enthalpy of Formation ($\Delta H_f^\circ$)

Definition and Calculation

Standard enthalpy of formation is the enthalpy change when 1 mol of a compound forms from its elements in their standard states (1 atm, 25°C).

• Standard state for gases: pure gas at 1 atm.

• Standard state for solids/liquids: pure substance at 1 atm, 25°C.

• Standard state for solutions: 1 M concentration.

Calculation:

• $\Delta H_{\text{rxn}}^\circ = \sum \Delta H_f^\circ (\text{products}) - \sum \Delta H_f^\circ (\text{reactants})$

Example:

• For $4\text{NH}_3(g) + 5\text{O}_2(g) \rightarrow 4\text{NO}(g) + 6\text{H}_2\text{O}(g)$, use tabulated $\Delta H_f^\circ$ values to calculate $\Delta H_{\text{rxn}}^\circ$.

Note: $\Delta H_f^\circ$ for elements in their most stable form is zero.