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n Chemistry Notes
A Comprehensive Study Guide
Covering Core Concepts from Atomic Structure to Organic Chemistry
Topic Page
1. Atomic Structure & the Periodic Table 2
2. Chemical Bonding 3
3. Stoichiometry & The Mole 4
4. States of Matter & Gas Laws 5
5. Thermodynamics 6
6. Chemical Equilibrium 7
7. Acids, Bases & pH 8
8. Electrochemistry & Redox 9
9. Organic Chemistry 10
1. Atomic Structure & The Periodic Table
1.1 Subatomic Particles
Particle Symbol Charge Mass (amu) Location Proton pn +1 1.007 Nucleus Neutron nn 0 1.009 Nucleus Electron en − 1 0.00055 Electron cloud
1.2 Atomic Number, Mass Number & Isotopes
The atomic number (Z) equals the number of protons, which defines the element. The mass number (A)
equals protons + neutrons. Isotopes are atoms of the same element with different numbers of neutrons.
Atomic Number (Z) = Number of protons
Mass Number (A) = Protons + Neutrons
Neutrons = A − Z
H Note: Isotopes have the same Z but different A values. Example: C-12, C-13, C-14 are all carbon.
1.3 Electron Configuration
Electrons occupy orbitals following three rules:
- Aufbau Principle: Fill orbitals from lowest to highest energy (1s → 2s → 2p → 3s...)
- Pauli Exclusion Principle: Each orbital holds max 2 electrons with opposite spins
- Hund's Rule: Fill each orbital singly before pairing electrons
Example — Iron (Fe, Z=26): [Ar] 3d^6 4s^2
1.4 Periodic Trends
Property Across Period ( → ) Down Group ( ↓ ) Atomic Radius Decreases Increases Ionization Energy Increases Decreases Electronegativity Increases Decreases Electron Affinity Generally increases Generally decreases
- Dipole-Dipole Forces: Between polar molecules; stronger than LDF for similar-sized molecules
- Hydrogen Bonding: Special dipole-dipole between H bonded to N, O, or F; strongest IMF (explains
high BP of water)
H Note: IMF strength order: H-bonding > Dipole-Dipole > London Dispersion Forces
3. Stoichiometry & The Mole
3.1 The Mole Concept
The mole is the SI unit for amount of substance. One mole contains exactly 6.022 × 10²³ particles
(Avogadro's number).
1 mol = 6.022 × 10^23 particles (Avogadro's Number)
Molar Mass = mass (g) per 1 mole of substance
n = m / M → moles = mass (g) / molar mass (g/mol)
3.2 Balancing Chemical Equations
The Law of Conservation of Mass states that atoms are neither created nor destroyed in a reaction.
Balanced equations have equal numbers of each atom on both sides.
Unbalanced: H 2 + O 2 → H 2 O
Balanced: 2H 2 + O 2 → 2H 2 O
3.3 Stoichiometric Calculations
Mole ratios from balanced equations act as conversion factors between reactants and products.
- Step 1: Write and balance the chemical equation
- Step 2: Convert given quantity to moles
- Step 3: Use mole ratio from equation
- Step 4: Convert moles to desired units
3.4 Limiting Reagent & Percent Yield
The limiting reagent is the reactant that is completely consumed first and determines the maximum
amount of product formed.
Theoretical Yield = max product from limiting reagent
% Yield = (Actual Yield / Theoretical Yield) × 100%
3.5 Solution Stoichiometry
Concentration is expressed as molarity (M) — moles of solute per liter of solution.
Molarity (M) = moles of solute / liters of solution
C 1 V 1 = C 2 V 2 (Dilution Formula)
4. States of Matter & Gas Laws
4.1 Kinetic Molecular Theory (KMT)
- Gas particles are in constant, random motion
- Particle volume is negligible compared to container volume
- No attractive/repulsive forces between ideal gas particles
- Average kinetic energy is proportional to absolute temperature (K)
4.2 The Gas Laws
Law Variables Formula Condition Boyle's Law P & V PnVn = PnVn Constant T, n Charles's Law V & T Vn/Tn = Vn/Tn Constant P, n Gay-Lussac's Law P & T Pn/Tn = Pn/Tn Constant V, n Avogadro's Law V & n Vn/nn = Vn/nn Constant T, P Combined Gas Law P, V, T PnVn/Tn = PnVn/Tn Constant n
4.3 Ideal Gas Law
PV = nRT
Where: P = pressure (atm), V = volume (L), n = moles, R = 0.08206 L·atm/mol·K, T = temperature (K)
H Note: Always convert temperature to Kelvin: K = °C + 273.
4.4 Real Gases & Van der Waals Equation
Real gases deviate from ideal behavior at high pressures and low temperatures. Van der Waals equation
corrects for intermolecular attractions (a) and molecular volume (b):
(P + an²/V²)(V − nb) = nRT
4.5 Phase Changes
Process Direction Energy Change Melting / Fusion Solid → Liquid Endothermic (+∆H) Freezing Liquid → Solid Exothermic (−∆H) Vaporization Liquid → Gas Endothermic (+∆H)
Condensation Gas → Liquid Exothermic (−∆H) Sublimation Solid → Gas Endothermic (+∆H) Deposition Gas → Solid Exothermic (−∆H)
6. Chemical Equilibrium
6.1 Dynamic Equilibrium
At equilibrium, the rates of forward and reverse reactions are equal, and the concentrations of reactants
and products remain constant (but not necessarily equal).
6.2 Equilibrium Constant (K)
For the reaction: aA + bB n cC + dD
Keq = [C]c[D]d^ / [A]a[B]b
- K >> 1: Equilibrium favors products (forward reaction)
- K << 1: Equilibrium favors reactants (reverse reaction)
- K = 1: Roughly equal concentrations of reactants and products H Note: Pure solids and liquids are NOT included in K expressions
6.3 Le Chatelier's Principle
When a system at equilibrium is disturbed, it shifts to minimize the disturbance and re-establish
equilibrium.
Stress Applied Direction of Shift Add reactant Toward products (→) Remove reactant Toward reactants (←) Add product Toward reactants (←) Increase pressure Toward fewer moles of gas Decrease pressure Toward more moles of gas Increase temperature Toward endothermic direction Decrease temperature Toward exothermic direction Add catalyst No shift (rate increases both ways)
6.4 Reaction Quotient (Q)
Q has the same form as K but uses current (non-equilibrium) concentrations. Comparing Q to K predicts
reaction direction.
Q < K → Reaction proceeds forward ( → ) to reach equilibrium
Q > K → Reaction proceeds backward ( ← ) to reach equilibrium
Q = K → System is at equilibrium
8. Electrochemistry & Redox Reactions
8.1 Oxidation & Reduction
- Oxidation: Loss of electrons (LEO — Loss of Electrons is Oxidation)
- Reduction: Gain of electrons (GER — Gain of Electrons is Reduction)
- Oxidizing agent: Causes oxidation; itself is reduced
- Reducing agent: Causes reduction; itself is oxidized H Note: Mnemonic: OIL RIG — Oxidation Is Loss, Reduction Is Gain
8.2 Assigning Oxidation Numbers
Rule Oxidation Number Pure element 0 Monatomic ion Equals the charge Oxygen in compounds −2 (except peroxides: −1) Hydrogen with nonmetals + Hydrogen with metals − 1 Fluorine Always − 1 Sum in neutral compound 0 Sum in polyatomic ion Equals ionic charge
8.3 Galvanic (Voltaic) Cells
- Convert chemical energy into electrical energy through spontaneous redox reactions
- Anode: Oxidation occurs (negative electrode in galvanic cell)
- Cathode: Reduction occurs (positive electrode in galvanic cell)
- Salt bridge: Maintains electrical neutrality by allowing ion flow
E°cell = E°cathode − E°anode
H Note: Spontaneous reaction: E°cell > 0, ∆ G < 0
8.4 Nernst Equation & Faraday's Law
E = E° − (RT/nF) × ln(Q) → At 25°C: E = E° − (0.0592/n) × log(Q)
Faraday's Law: m = (M × I × t) / (n × F)
Where: m = mass deposited, M = molar mass, I = current (A), t = time (s), n = electrons transferred, F =
96,485 C/mol
Isomers are compounds with the same molecular formula but different structural arrangements:
- Structural (constitutional) isomers: Different connectivity of atoms (e.g., n-butane vs isobutane)
- Stereoisomers — cis/trans: Same connectivity, different spatial arrangement around a double bond
or ring
- Enantiomers: Mirror-image isomers around a chiral center (non-superimposable); rotate
plane-polarized light
H Note: Chirality is crucial in pharmacology — different enantiomers can have very different biological effects.
