Unit 1: Atomic Structure

Energy is Quantized!

• Formula: E = hf = hc/λ
• Wavelength and energy relationships:
  • UV > Vl > IR (in terms of energy and frequency)
  • Wavelength increases in the order: UV < Vl < IR

Photoelectric Effect

• Light photons of specific wavelengths or frequencies are used to excite or eject electrons.
• Formula: Energy of photons = B.E (Binding Energy) + K.E (Kinetic Energy).

Coulombic Force

• Coulomb’s Law: q₁q₂ / r
• Ionic, metallic, Lattice, Hydration

Electron Configuration

• 1s 2s 2p 3s 3p 4s 3d
• Cr , Cu exception

Photoelectron Spectroscopy (PES)

• Intensity (y-axis) corresponds to the number of electrons.
• Binding energy (x-axis) represents electron removal energy.

Periodic Trends

Proton no = Nuclear charge

Shell no = Energy level = Shielding no

Ions – HO¹⁻, NH₄¹⁺, NO₃¹⁻ , CO₃²⁻ , SO₄²⁻ , PO₄³⁻

Ionization energy Excemption

• Ionization energy values decrease:
  1. From Be to B (due to p orbital initiation).
  2. From N to O (due to paired electrons in p orbital).
• Electron Affinity ( Exothermic)

Mass Spectrometer

• Calculates average atomic mass.
• Example: Bromine isotopes
  • ⁷⁹Br (50.5%)
  • ⁸¹Br (49.5%)

Unit 2: Chemical Bonding

Metallic Bond

• Free electrons = delocalized electrons = valence electrons.
• Examples: Li < Mg (bond strength – Coloumb law)
• Metal related to Ionization Energy , Nonmetal related to Electron Affinity
• Malleability and Ductility
• Types:
  • Interstitial Alloy: Different radii, more rigid.
  • Substitutional Alloy: Comparable radii, malleable, ductile.

Ionic Bond

• Transferring electron(s)
• Bond strength proportional to Coulomb’s law.
• Examples: LiF > NaCl , LiF < MgO.
• Lattice Energy:
  • Formula: XY(s) → X⁺(g) + Y⁻(g).
• lattice stucture

Covalent Bond

• Sharing electrons.
• Types:
  • Network Covalent Solids: Ex) C (diamond, graphite), SiO₂ , Si
  • Molecular Substances: Ex) H₂O, CO₂, N₂.
  • allotrope vs isotope vs isomer
• Strength – Intermolecular Forces

Covalent Bonds and Potential Energy

Lewis Struture

CH₄, CH₂=CH₂, NH₃ , NH₄⁺ , CH₃COOH, CH₃COO ⁻ , CH₃COOCH₃, CO₃²⁻ , SO₄²⁻ , PO₄³⁻

from period 3, octet expanded

Bond Energy

• Energy required to dissociate a bond in a gaseous state.
• Endothermic , (Bond Formation – Exothermic)

Bond Length

• Distance between nuclei where potential energy is minimized.

Polar vs. Nonpolar Bonds

• Bond : Electronegativity difference
• Molecular Polar – Asymmetric arrangement, Bond polarities not Cancel out, dipole Moments.

Resonance

• Not dynamic equilibrium
• Equivalent bond length and strength
• Only their average form exists (Bond order / # Resonance)
• Delocalized pi electrons.
• Enhance The Stability
• C₆H₆ , NO₃¹⁻ , CO₃²⁻ , SO₄²⁻ , PO₄³⁻ examples

Formal Charge

• Formal charge = original v.e – (# Lone e + # Bond )
• Zeros are preferred!

VSEPR Theory

• Valence Shell Electron Pair Repulsion.
• Minimizes repulsion:
  • Bonding pair–bonding pair < bonding pair–lone pair < lone pair–lone pair.
  • Molecular geometry vs Electron domain geometry

Hybridization

• Types: sp, sp², sp³.

Unit 3: IMF and Gas

Intermolecular Forces

• Strength – Intermolecular Forces determine melting and boiling points, Vapour pressure and Volatility. (Physical state)

London Dispersion Forces (LDF)

• Instantaneous dipoles, Temporary Foreces, Induced Dipoles.
• Present between all molecules, especially nonpolar.
• Increased by:
  • Larger # electrons.
  • Larger electron clouds.
  • Greater polarizability.

Dipole-Dipole Interaction

• Present between polar molecules.
• Stronger than LDF but weaker than hydrogen bonding.
• Permanent Forces

Hydrogen Bonding

• Strongest intermolecular force.
• Examples: H bonded to F, O, or N.

Dipole-Induced Dipole Force

• Present between polar and nonpolar molecules.
• Stronger than LDF but weaker than dipole-dipole forces.

Ideal Gas Law and Partial Pressure

• Formula: PV = nRT
• Partial Pressure : Total pressure x mole Fraction

Kinetic Molecular Theory

• Ideal gas behavior
  1. Random & constant motion
  2. Perfectly elastic collision
  3. No intermolecular forces
  4. Negligible volume of particles (point-like)

Kinetic Energy vs. Speed

• Although AVERAGE kinetic energy values are always constant for all gases
• Speed = v∝√(T/Mr)

Ideal Gas vs. Real Gas

• Ideal Gas
  • Higher T, Lower P, Volum negligible, No Attraction
• Condition for less ideal:
  • High P ➔ The combined volume of the particles becomes significant.
  • Low T ➔ The intermolecular forces become significant due to their slow motion.

Maxwell-Boltzmann Distribution

Unit 4. Stoichiometry

Some Formulas

• Percent Yield: (actual yield / theoretical yield) × 100
• Percent Error: |theoretical − actual| / theoretical × 100
• Average Atomic Mass: Σ(mass of each isotope × % abundance)

Solubility Rule

• Every group 1A elements, NH₄⁺, NO₃⁻ containing salts are always soluble without exception.
• Strong Acid – H ⁺ , OH ⁻

Balancing Redox Reactions

1. Balance element except O, H
2. Balance O using H₂O
3. Balance H using H⁺
4. Balance charge using e⁻

Unit 5: Kinetics

Reaction Rate

• Definition: The slope of the graph of concentration vs. time.
• Change in concentration over time.
• Unit: M/s or mol/L·s.

Collision Theory

1. Must Collide
2. Proper Orientation
3. Sufficient Energy greater than activation energy.→ Leads to effective collision (enables the reaction to occur).

Factors Affecting Reaction Rates

1. Surface Area ( Smaller Size ) : Frequencies of Collision
2. Concentration. – Frequencies of Collision
3. Pressure and Volume. – Frequencies of Collision
4. Temperature. 1) Velocity – Frequencies of Collision , More particles have Higher KE than Activation Energy
5. Catalyst. – Lower Activation Energy → More particles have Higher Kinetic Energy than Activation Energy

Rate Law

• General Form: mA + nB → product Rate = k × [A]ᵐ × [B]ⁿ
• Reaction orders (m, n) are determined by experiment only.
• The unit of rate constant, k, depends on the overall reaction order.
• Rate constant is affected by temperature and the presence of a catalyst.

Rate Law Determination

Experimental rate law (concentration vs. rate).

Rate law Diagram

Integrated rate law (concentration vs. time).

• Zero Order: [A] = -kt + [A]₀
• First Order: ln[A] = -kt + ln[A]₀ (Half-life = 0.693/k).
• Second Order: 1/[A] = kt + 1/[A]₀

Reaction mechanism.

• A multistep reaction occurs as follows: 2A + B ⇌ C (fast, reversible) C + E → D + A (slow).
• Overall Reaction Equation: A + B + E → D Intermediate: **C** Catalyst: **A**
• Rate Law:Rate = k[A]²[B].

Unit 6: Thermodynamics

Key Topics:

• Exothermic vs. Endothermic Reactions
• Standard Enthalpy Change (ΔH°) Calculations:
  • Using calorimetric data: ΔH° = -q / mol , q = mct
  • Hess’s Law: Sum of enthalpy changes of individual reactions
  • Using bond energies, Combustaion energies: ΔH° = ΣB.E(reactants) – ΣB.E(products)
  • Using formation energies **ΔH° = Σproducts - ΣReactnats**

Entropy change (ΔS) Calculations

1. Distribution higher → entropy change increases
2. Calculation : **ΔH° = Σproducts - ΣReactnats**

Spontaneity of Reactions

• ΔH°: Exothermic reactions (ΔH° < 0) are preferred.
• ΔS°: Higher entropy (ΔS° > 0) is preferred.
• ΔG°: Determines spontaneity (ΔG° = ΔH° – TΔS°).

Gibbs Free Energy (ΔG°):

• If ΔG° < 0: Reaction is thermodynamically favored. ΔG° = ΔH° – TΔS° ΔG° = -RT lnK ΔG° = -nFE°
• Temperature-dependent spontaneity:
  • High T: ΔS-driven.
  • Low T: ΔH-driven.

Unit 7: Equilibrium

• Meaning of Equilibrium
  • Concentrations Constant
  • Forward and reverse reaction rates are equal.

• Reversible vs. Irreversible Reactions
  • Irreversible reactions (e.g., combustion, precipitation and gas produced)
• Expression:Kₑq = [C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ
  • Includes only gases and aqueous phases.

Factors Affecting Equilibrium

• Temperature:The only factor that changes Keq.

Relationship Between Keq and Q

• Q > Keq: Shifts left (toward reactants).
• Q < Keq: Shifts right (toward products).
• Q = Keq: The system is at equilibrium.

Le Chatelier’s Principle

• Types of Stress:
  1. Concentration Changes: Adjusts to restore equilibrium.
  2. Pressure Changes (via Volume):Affects gases; shifts to favor fewer or more moles of gas.
  3. Temperature Changes:Alters Keq (forward reaction is endothermic or exothermic reactions).
  4. Catalysts:Speed up the both reactions equally but do not affect Keq.
• Examples:
  • N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH < 0

Manipulation of Equilibrium – rule

Ksp

• Definition of Ksp
  • Soluble salt: AB₍aq₎ → A⁺₍aq₎ + B⁻₍aq₎ (Completely dissociated)
  • Insoluble salt: AB₍ₛ₎ ⇌ A⁺₍aq₎ + B⁻₍aq₎ / Kₛₚ = [A⁺][B⁻]
• Basic Calculation
  • If molar solubility = x:
    • Kₛₚ of AB compound = x²
    • Kₛₚ of AB₂ compound = 4x³
• Common Ion Effect

Unit 8 : Acids and Bases

Common Acids and Bases

• Strong Acids: HCl, HBr, HI, H₂SO₄, HNO₃, HClO₄,
• Strong Bases: NaOH, KOH, Group 1A, 2A hydroxides.
• Weak Bases: NH₃, Fe(OH)₂, Al(OH)₃, Transition metal + OH

pH Calculation Formulas

1. Strong Acids and Bases pH (Strong Acid)=−log[H+] pOH (Strong Base)=−log[OH−] pH + pOH = 14
2. Weak Acids pH = -log√(Ka × C)

pOH = -log√(Kb × C)

Percent Dissociation

%Dissociation=√(Ka/C) × 100

Acids and Bases with Water

• HCl + H₂O → H₃O⁺ + Cl⁻
• Cl⁻ + H₂O → HCl + OH⁻
• HF + H₂O ⇌ H₃O⁺ + F⁻
• F⁻ + H₂O ⇌ HF + OH⁻

Neutralization (Acid + Base , Carbonate and Metal)

• HCl + NaOH → H₂O + NaCl
  • H⁺ + OH⁻ → H₂O
• HF + NaOH → H₂O + NaF
  • HF + OH⁻ → H₂O + F⁻
• HCl + Na₂CO₃ → H₂O + CO₂ + 2NaCl Ionic Equation: 2H⁺ + CO₃²⁻ → H₂O + CO₂
• HCl + Zn → ZnCl₂ + H₂ Ionic Equation: 2H⁺ + Zn → Zn²⁺ + H₂

Dissociation of a Polyprotic Acid

• H₂CO₃ + H₂O ⇌ H₃O⁺ + HCO₃⁻ Kₐ₁
• HCO₃⁻ + H₂O ⇌ H₃O⁺ + CO₃²⁻ Kₐ₂

Key Point:

• Kₐ₁ >> Kₐ₂

Neutralization of a Polyprotic Acid

• H₂CO₃ + 2NaOH → 2H₂O + Na₂CO₃
  • Step 1: H₂CO₃ + NaOH → H₂O + NaHCO₃
  • Step 2: NaHCO₃ + NaOH → H₂O + Na₂CO₃

pH at the equivalence point

• Strong acid + strong base → neutral salt Example: HCl + NaOH → NaCl + H₂O / H⁺ + OH⁻ ⇌ H₂O
• Strong acid + weak base → acidic salt Example: HCl + NH₃ → NH₄Cl / H⁺ + NH₃ → NH₄⁺
• Weak acid + strong base → basic salt Example: HC₂H₃O₂ + NaOH → NaC₂H₃O₂ + H₂O HC₂H₃O₂ + OH⁻ → C₂H₃O₂⁻ + H₂O
• Weak acid + weak base → X
• CV = CV (Volume and Concentration)
• Whether acid is strong or weak, the volume of base is identical

Hydrolysis of Salt

Why NH₄Cl is acidic

• Cl⁻ ion cannot react with water molecules, but NH₄⁺ does to a small extent.Example:
• NH₄⁺ + H₂O ⇌ NH₃ + H₃O⁺

Why NaCl is neutral

• Na⁺, Cl⁻ ions cannot react with water molecules.

Why NaC₂H₃O₂ is basic

• Na⁺ ion cannot react with water molecules, but C₂H₃O₂⁻ does to a small extent.Example:
• C₂H₃O₂⁻ + H₂O ⇌ HC₂H₃O₂ + OH⁻

How to prepare a buffer

Composition (by 1:1)

1. Weak acid + conjugate base HF + NaF / HC₂H₃O₂+NaC₂H₃O₂
2. Weak base + conjugate acid NH₃ + NH₄Cl

Composition (by 2:1)

1. Weak acid + strong base HF + NaOH / HC₂H₃O₂ + NaOH
2. Weak base + strong acid NH₃ + HCl

Half-equivalence point

• Best buffer ratio!
• pH = pKa

Handerson-Hasselbalch equation

• pH = pKa + log [A⁻]/[HA]
• Used for only a buffered solution.

Choosing a good indicator

• pKa of the indicator needs to be close to the pH at the equivalence point.
• Lechatelier’s priciple

W.A. and S.B. (buffering region)

1. pH < pKa: [HC₂H₃O₂] > [C₂H₃O₂⁻ ]
2. pH = pKa: [HC₂H₃O₂] = [C₂H₃O₂⁻ ]
3. pH > pKa: [HC₂H₃O₂] < [C₂H₃O₂⁻ ]
4. higher Ka , higher Dissociation(Ionization)

Unit 9 : Cell and Others

Reaction:

• Cu² ⁺ + 2e⁻ → Cu , E° (Reduction Potential) =+0.34V
• Zn² ⁺ + 2e⁻ → Zn , E° (Reduction Potential) =−0.76V

Zn + Cu² ⁺ → Zn²⁺ + Cu, E° cell (Cell Potential) = 1.10 V

How Ecell changes with concentrations:

1. Cell #1:
   • [Cu²⁺] = 1.0 M, [Zn²⁺] = 1.0 M
   • E° cell = 1.10 V
2. Cell #2:
   • [Cu²⁺] = 0.1 M, [Zn²⁺] = 1.0 M
   • Ecell < 1.10 V
3. Cell #3:
   • [Cu²⁺] = 1.0 M, [Zn²⁺] = 0.1 M
   • Ecell > 1.10 V
4. Cell #4:
   • [Cu²⁺] = 0.1 M, [Zn²⁺] = 0.1 M
   • E° cell = 1.10 V, but it runs for a shorter time.

Equilibrium and Dead Battery

• Q = 1: E° cell = Ecell
• Q > 1:
  • More Products (P), Less Reactants (R)
  • Ecell < E° cell
  • Closer to equilibrium
• Q < 1:
  • Less Products (P), More Reactants (R)
  • Ecell > E° cell
  • Farther from equilibrium

Electrolysis

Electrolysis of molten salt:

• Reaction: 2NaCl → 2Na + Cl₂
  • Na⁺ + e⁻ → Na (E° = -2.71 V)
  • Cl₂ + 2e⁻ → 2Cl⁻ (E° = 1.36 V)
  • Overall cell potential: E° cell = -4.07 V
• External electrical current > 4.07 V is needed.
• Electrolysis is non-spontaneous and requires energy.

Reactivity Comparison

1. Order of reactivity for metals X, Y, Z:
   • If X doesn’t react with Y(NO₃)₂ but reacts with Z(NO₃)₂:
     • Reactivity: X < Z < Y (The highest Reducing agent)
   • Tendency to lose electrons determines reactivity.

Metal Deposition

• Equation : q = It = nF
  1. How many grams of Al can be deposited from Al(NO₃)₃ solution by 2.0 A for 10 minutes?
     • Answer: 0.11 g
  2. How long does it take to deposit 10.0 g of Cu from Cu(NO₃)₂ solution by 5.0 A?
     • Answer: 6100 seconds or 1.0 x 10² minutes

Laboratory

Solution Preparation (1): How to prepare 100 mL of 1.00 M NaCl solution

1. Measure 5.85 g of NaCl using an electronic balance.
2. Add the NaCl to a 100 mL volumetric flask.
3. Add a small amount of water and swirl to dissolve completely.
4. Fill the flask with water until the meniscus reaches the 100 mL mark.

Solution Preparation (2): How to prepare 500 mL of 0.10 M H₂SO₄ solution

1. Wear safety goggles and gloves.
2. Use M₁V₁ = M₂V₂ to calculate the volume of concentrated H₂SO₄ needed (25 mL).
3. Add water to a volumetric flask first.
4. Carefully add the acid to the flask.
5. Add more water until the total volume reaches 500 mL.

Beer’s Law

• A = abc
  • A: Absorbance (no unit)
  • a: Absorptivity constant (M⁻¹cm⁻¹)
  • b: Path length (cm)
  • c: Concentration (M)
• Graph: Absorbance vs. Concentration
• Chromatography As solvent rises onto the paper, solute A~D reaches specific points depending on their intermolecular forces with the solvent molecules. In order of increasing interaction with the solvent: A < B < C < D
• Physical Separation Methods
  • Filtration: Insoluble solid from liquid
  • Distillation: Soluble solid from liquid or miscible liquids using boiling point differences

Lab Safety

• Never pour water to acid, as it might cause vigorous splattering due to rapid heat increase.
• Pour acid to water slowly to ensure proper absorption of heat.
• Never return residue of chemicals to the reagent container.

Titration

• Be careful with the buret:
  • Rinse it with titrant twice before use. – water , titrant
• Choose an indicator with a pKa close to the equivalence point pH.

Gravimetric Analysis

• Example: A student determines the mass % of silver in a copper-silver alloy.
  1. Add saturated NaCl(aq) until no more precipitate forms.
  2. Filter, wash, dry, and weigh the precipitate.
  3. Repeat until a constant mass is achieved.

Dehydration of Hydrated Salt

• Cover the crucible to avoid mass loss.
• Measure repeatedly to ensure complete dehydration.

Gas Collecting via Water Displacement

• Adjust the graduated cylinder to equalize water levels.
• Consider vapor pressure, determined by temperature.

Significant Figures

• Examples:
  • 0.010 g
  • 100. mL
  • log 0.010 = 2.00
• Rules:
  1. Addition/Subtraction: Use the fewest decimal places.
  2. Multiplication/Division: Use the fewest significant figures.