Essential Reference Tables for General Chemistry

Reference Tables in General Chemistry

Introduction

Reference tables are crucial tools in General Chemistry, providing standardized data for calculations, problem-solving, and conceptual understanding. The following study notes summarize key appendices commonly found in chemistry textbooks, including thermodynamic data, bond enthalpies, solubility products, formation constants, acid-base constants, reduction potentials, ionization energies, electron affinities, and the periodic table of elements.

Thermodynamic Data

Standard Enthalpy of Formation ($\Delta H_f^\circ$), Standard Entropy ($S^\circ$), and Standard Gibbs Free Energy of Formation ($\Delta G_f^\circ$)

These values are essential for calculating the energy changes in chemical reactions and predicting spontaneity.

• Standard Enthalpy of Formation ($\Delta H_f^\circ$): The heat change when one mole of a compound is formed from its elements in their standard states at 25°C and 1 atm.

• Standard Entropy ($S^\circ$): The absolute entropy of a substance at 25°C and 1 atm.

• Standard Gibbs Free Energy of Formation ($\Delta G_f^\circ$): The change in free energy when one mole of a compound is formed from its elements in their standard states.

These values are used in the following equations:

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

• Standard reaction free energy: $\Delta G_{rxn}^\circ = \sum \Delta G_f^\circ (\text{products}) - \sum \Delta G_f^\circ (\text{reactants})$

• Relationship between free energy, enthalpy, and entropy: $\Delta G^\circ = \Delta H^\circ - T\Delta S^\circ$

Average Bond Enthalpies (25°C)

Definition and Application

Bond enthalpy is the energy required to break one mole of a specific type of bond in a gaseous molecule. Average bond enthalpies are used to estimate the enthalpy change of reactions:

• Bond breaking is endothermic (energy absorbed).

• Bond forming is exothermic (energy released).

The enthalpy change for a reaction can be estimated as:

• $\Delta H_{rxn} \approx \sum \text{(Bond enthalpies of bonds broken)} - \sum \text{(Bond enthalpies of bonds formed)}$

Bond	Bond enthalpy (kJ/mol)
C–H	413
O–H	463
C=O	799
H–H	436

Solubility Products ($K_{sp}$) at 25°C

Definition and Use

The solubility product constant ($K_{sp}$) quantifies the equilibrium between a solid and its ions in a saturated solution. It is used to predict precipitation and calculate solubility.

• For $\text{AgCl (s)} \rightleftharpoons \text{Ag}^+ (aq) + \text{Cl}^- (aq)$, $K_{sp} = [\text{Ag}^+][\text{Cl}^-]$

• Low $K_{sp}$ values indicate low solubility.

Salt	Solubility equilibrium	$K_{sp}$
AgCl	AgCl(s) ⇌ Ag+ + Cl-	1.8 × 10-10
BaSO4	BaSO4(s) ⇌ Ba2+ + SO42-	1.1 × 10-10

Cumulative Formation Constants of Complexes ($\beta$) at 25°C

Definition and Application

Formation constants ($\beta$) describe the equilibrium for the stepwise formation of complex ions from a metal ion and ligands. They are used to predict the stability of complexes in solution.

• For $\text{Ag}^+ + 2\text{NH}_3 \rightleftharpoons [\text{Ag(NH}_3)_2]^+$, $\beta_2 = \frac{[\text{Ag(NH}_3)_2^+]}{[\text{Ag}^+][\text{NH}_3]^2}$

• High $\beta$ values indicate stable complexes.

Complex	$\beta$
[Ag(NH3)2]+	1.6 × 107
[Cu(NH3)4]2+	2.1 × 1013

Acidity and Basicity Constants for Weak Acids and Bases (25°C)

Acid Dissociation Constant ($K_a$) and Base Dissociation Constant ($K_b$)

These constants measure the strength of weak acids and bases in water.

• $K_a$ for acids: $\text{HA} \rightleftharpoons \text{H}^+ + \text{A}^-$, $K_a = \frac{[\text{H}^+][\text{A}^-]}{[\text{HA}]}$

• $K_b$ for bases: $\text{B} + \text{H}_2\text{O} \rightleftharpoons \text{BH}^+ + \text{OH}^-$, $K_b = \frac{[\text{BH}^+][\text{OH}^-]}{[\text{B}]}$

• Smaller $K_a$ or $K_b$ values indicate weaker acids or bases.

Acid/Base	Name	$K_a$ or $K_b$
CH3COOH	acetic acid	1.8 × 10-5
NH3	ammonia	$K_b = 1.8 × 10^{-5}$

Standard Reduction Potentials (25°C)

Electrochemical Series

Standard reduction potentials ($E^\circ$) indicate the tendency of a species to gain electrons (be reduced) under standard conditions. They are used to predict the direction of redox reactions and calculate cell potentials.

• Standard hydrogen electrode (SHE) is assigned $E^\circ = 0.00$ V.

• Cell potential: $E_{cell}^\circ = E_{cathode}^\circ - E_{anode}^\circ$

Half-reaction	$E^\circ$ (V)
Cu2+ + 2e- → Cu(s)	+0.34
Zn2+ + 2e- → Zn(s)	-0.76

Ionization Energies and Electron Affinities (First 36 Elements at 25°C)

Trends and Definitions

• Ionization energy (IE): The energy required to remove an electron from a gaseous atom or ion.

• Electron affinity (EA): The energy change when an electron is added to a gaseous atom.

• Trends: IE generally increases across a period and decreases down a group. EA is more negative for nonmetals.

Element	IE1 (kJ/mol)	EA (kJ/mol)
H	1312	-73
O	1314	-141

Table of Molar Masses and the Periodic Table of the Elements

Atomic Masses and Periodic Trends

• Molar mass: The mass of one mole of a substance, expressed in g/mol.

• Used for stoichiometric calculations and conversions between mass and moles.

• The periodic table organizes elements by increasing atomic number and groups elements with similar chemical properties.

Element	Symbol	Atomic Number	Molar Mass (g/mol)
Hydrogen	H	1	1.008
Carbon	C	6	12.011
Oxygen	O	8	15.999

Summary Table: Key Reference Data in General Chemistry

Appendix	Main Purpose
Thermodynamic Data	Calculate reaction enthalpy, entropy, and free energy
Bond Enthalpies	Estimate reaction enthalpy from bond energies
Solubility Products	Predict precipitation and calculate solubility
Formation Constants	Assess stability of complex ions
Acid/Base Constants	Calculate pH and equilibrium for weak acids/bases
Reduction Potentials	Predict redox reaction direction and cell voltage
Ionization Energies/Electron Affinities	Understand periodic trends and reactivity
Molar Masses/Periodic Table	Stoichiometry and element identification

Additional info: These tables are essential for solving quantitative and conceptual problems in General Chemistry, including thermochemistry, equilibrium, electrochemistry, and descriptive chemistry.