Thermodynamics

Page 1: Atomic Structure and Properties

Key Concepts

1. Ionization Energy and Electron Configuration

   • Higher energy levels = further from nucleus = lower Coulombic attraction = easier removal (lower 1st ionization energy).

   • Across a period, Z_eff increases, attracting valence electrons stronger, leading to decreased atomic radius and increased ionization energy.

2. Photoelectron Spectroscopy (PES)

   • Higher peaks indicate more electrons in sublevels.

   • Larger binding energy = electrons closer to nucleus.

3. Electron Configuration

   • Example: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶.

   • For cations, remove valence electrons in order from s to p, and if necessary, from d orbitals last.

4. Isotopes

   • Same protons, different neutrons.

   • Mass spectroscopy measures atomic masses of isotopes.

5. Periodic Trends

   • Elements in the same group share similar properties.

   • Metals on left side, nonmetals on right.

   • Cations smaller than original atoms, anions larger due to added electrons.

Page 2: Molecular and Ionic Compound Structure and Properties

Bonding Concepts

1. Covalent Bonds

   • Formed by two nonmetals sharing electrons.

2. Ionic Bonds

   • Formed by metal transferring electrons to a nonmetal.

3. Bond Polarity

   • Greater electronegativity difference = more polar bond.

4. Combustion Reactions

   • Products: CO₂ and H₂O.

   • Carbon forms four bonds.

Geometry and Hybridization

1. Bond Angles

   • 4 domains = 109.5°.

   • 3 domains = 120°.

   • 2 domains = 180°.

   • Corresponding hybrid orbitals: sp³, sp², sp.

2. Polarity of Molecules

   • Asymmetrical = polar; symmetrical = nonpolar.

3. Bond Types

   • Single bond = sigma; double bond = sigma + pi; triple bond = sigma + 2 pi.

Page 3: Intermolecular Forces and Properties

Intermolecular Forces (IMFs)

1. Strengths of IMFs

   • From weakest to strongest: London Dispersion, dipole-dipole, hydrogen bonding, ion-dipole.

   • All molecules have London dispersion forces; strength increases with size.

2. Boiling and Melting Points

   • Increase with stronger IMFs; vapor pressure and volatility decrease.

Solid Structures

1. Molecular Solids

   • Low melting/boiling points, non-conductive.

2. Ionic Solids

   • High melting/boiling points; conductive in liquid/aq state.

3. Covalent Network Solids

   • Very high melting/boiling points (e.g., SiO₂, diamonds).

Gases

1. Gas Characteristics

   • Homogeneous mixtures due to constant particle motion; compressibility due to space between particles.

   • Pressure is due to particle collisions with container walls.

2. Gas Laws Relationships

   • P and V inversely related; T and V, T and P directly related.

3. Ideal Gas Constant

   • PV = nRT.

   • 1 mole of ideal gas = 22.4 L at STP.

Page 4: Chemical Reactions

Chemical Reaction Basics

1. Diatomic Elements

   • H₂, N₂, Cl₂, Br₂, I₂, F₂ (e.g., "I Have No Bright Or Clever Friends").

2. Empirical and Molecular Formulas

   • Empirical: simplest ratio.

   • Molecular: multiple of empirical.

3. Percentage Yield and Error

   • % yield = (experimental/theoretical) x 100%.

   • % error = (experimental - theoretical)/theoretical x 100%.

4. Limiting Reactants

   • Product amount determined by limiting reactant.

   • Mass is conserved in all changes.

Page 5: Kinetics

Reaction Rates

1. Collision Theory

   • Particles must collide effectively, overcoming activation energy; height of the "hill" is the activation energy.

2. Rate Law for Elementary Steps

   • Example: 2A + B → C + D; Rate = k[A]²[B].

3. Factors Affecting Reaction Rate

   • Add catalyst; increase concentration, surface area, pressure, and temperature to increase collision frequency.

4. Half-Life

   • 1st order: t₁/₂ = 0.693/k.

   • Constant half-life for 1st order reactions.

Page 6: Thermochemistry

Enthalpy and Reactions

1. Exothermic Reactions

   • Negative ΔH, feels hot; heat is a product.

   • ΔHrxn = Bonds broken - Bonds formed.

2. Changing Reactions

   • ΔH doubles when the reaction is doubled; the sign changes when reversed; add ΔH’s when adding reactions.

Page 7: Equilibrium

Equilibrium Concepts

1. Equilibrium Constant (K)

   • K_eq = [products]⁡x/[reactants]⁡y.

   • Large K_eq means more products at equilibrium, small means more reactants.

2. Le Chatelier's Principle

   • Q > K_eq shifts left; changes in pressure shift if moles differ.

3. Solubility Equilibrium

   • Ksp formulas: 2 ions = Ksp = x²; 3 ions = Ksp = 4x³.

Page 8: Acids and Bases

Acid-Base Properties

1. pH Scale

   • Acids < 7, bases > 7; neutral at 25°C.

2. Ionization

   • Acids donate H⁺, bases accept H⁺.

3. Strong Acids and Bases

   • Strong acids: HNO₃, H₂SO₄, HClO₄, HBr, HI, HCl.

   • Strong bases: Group 1 and soluble Group 2 hydroxides.

Page 9: Applications of Thermodynamics

Thermodynamic Favorability

1. Spontaneous Reactions

   • Favorable if (−)ΔG.

   • Reactions with (−)ΔH and (+)ΔS are always favorable.

2. Equilibrium Constant and Free Energy

   • ΔG = 0 at equilibrium; ΔGo = −RT lnK.

Electrochemistry

1. Oxidation Numbers and Reactions

   • LEO goes GER; oxidation occurs at the anode.

2. Batteries

   • Electrons flow anode (−) to cathode (+); cations to cathode, anions to anode.

3. Electrochemical Calculations

   • ΔGo = −nFEo; n = electrons transferred.