Chem honors unit 1

1. SI Units (Base Units)

• Time → second (s)
  
  

• Length → meter (m)
  
  

• Mass → kilogram (kg)
  
  

• Temperature → kelvin (K)
  
  

• Amount of substance → mole (mol)
  
  

• Electric current → ampere (A) → 1 A = 1 C/s
  
  

• Luminous intensity → candela (cd)
  
  

Why important: Every measurement in chemistry reduces to these.

2. Derived Units

• Force (Newton, N) → kg·m/s²
  
  

• Energy (Joule, J) → kg·m²/s²
  
  

• Pressure (Pascal, Pa) → N/m² = kg/(m·s²)
  
  

• Charge (Coulomb, C) → A·s
  
  

• Density (ρ) → kg/m³ (≈ 1 g/cm³ for water)
  
  

3. Metric Prefixes

• Giga (G) = 10⁹
  
  

• Mega (M) = 10⁶
  
  

• Kilo (k) = 10³
  
  

• Centi (c) = 10⁻²
  
  

• Milli (m) = 10⁻³
  
  

• Micro (μ) = 10⁻⁶
  
  

• Nano (n) = 10⁻⁹
  
  

• Pico (p) = 10⁻¹²
  
  

4. Dimensional Analysis (Factor-Label Method)

Process: Write conversion factors as fractions so units cancel.

• Example 1: Convert 175 lb to kg (1 lb = 0.454 kg)
  175 lb × (0.454 kg / 1 lb) = 79.5 kg
  
  

• Example 2: Convert 2.50 km to cm
  2.50 km × (1000 m / 1 km) × (100 cm / 1 m) = 2.50 × 10⁵ cm
  
  

5. Significant Figures

• Multiplication/division → keep fewest sig figs
  
  

• Addition/subtraction → keep fewest decimals
  
  

Example: 6.82 + (3.1 × 5.76) = 6.82 + 17.9 = 24.7 → 25 (2 sig figs)

6. Atomic Structure & Ions

• Atomic number Z = # protons
  
  

• Mass number A = protons + neutrons
  
  

• Electrons = protons (neutral atom) ± charge
  
  

Examples:

• ¹⁵O²⁻ → p=8, n=7, e=10
  
  

• ⁶⁰Co²⁺ → p=27, n=33, e=25
  
  

• ⁴⁰K⁺ → p=19, n=21, e=18
  
  

Isoelectronic species: Cl⁻, Ar, K⁺, Ca²⁺ (all 18 e⁻).

7. Average Atomic Mass

Formula: Σ(fraction × mass)

Weird probability stuff

Examples:

• Copper → 0.73(63) + 0.27(65) = 63.54 amu
  
  

• Neon → 0.906(20) + 0.094(22) = 20.18 amu
  
  

8. The Mole

• 1 mol = 6.022 × 10²³ particles
  
  

Example: 12.0 g of C-14 → protons

• n = 12.0 ÷ 14.0 = 0.857 mol
  
  

• atoms = 0.857 × 6.022×10²³ = 5.16×10²³ atoms
  
  

• protons = 6 × 5.16×10²³ = 3.10×10²⁴ protons
  
  

9. Electrostatics

Coulomb’s Law: F = k (q₁q₂) / r²

• Opposite charges → attract
  
  

• Like charges → repel
  
  

• Halve r → F × 4
  
  

Example: q₁=+2 μC, q₂=−3 μC, r=0.50 m
F = 0.216 N (attractive)

10. Quantum Numbers & Orbitals

• n = principal shell (1, 2, 3…)
  
  

• l = subshell (0=s, 1=p, 2=d, 3=f)
  
  

• mₗ = orbital orientation (−l … 0 … +l)
  
  

• mₛ = spin (+½ or −½)
  
  

Rules:

• Pauli Exclusion → no two e⁻ same 4 numbers
  
  

• Hund’s Rule → fill singly before pairing
  
  

• Aufbau → fill lowest energy first
  
  

Exceptions:

• Cr = [Ar]4s¹3d⁵
  
  

• Cu = [Ar]4s¹3d¹⁰
  
  

11. Light & Photons

• λν = c
  
  

• E = hν = hc/λ
  
  

Example: λ = 510 nm
E = (6.626×10⁻³⁴ × 3.00×10⁸) ÷ (5.10×10⁻⁷) = 3.90×10⁻¹⁹ J

Spectrum order: radio < microwave < IR < visible < UV < x-ray < gamma

12. Periodic Trends

• Zeff (effective nuclear charge): increases →, constant ↓
  
  

• Atomic radius: decreases →, increases ↓
  
  

• Ionization energy: increases →, decreases ↓
  
  

• Electron affinity: more negative → (highest for halogens)
  
  

• Ionic size: cations smaller, anions larger
  
  

13. Bonding

• Bonds form to lower energy
  
  

• Bond length = balance of attraction & repulsion
  
  

• Bond energy = energy needed to break bond
  
  

• Bond order ↑ → length ↓, energy ↑
  
  

📘 FORMULA BANK WITH DEFINITIONS + EXAMPLES

Fundamental Constants

• c = 3.00 × 10⁸ m/s → speed of light
  
  

• h = 6.626 × 10⁻³⁴ J·s → Planck’s constant
  
  

• k = 8.99 × 10⁹ N·m²/C² → Coulomb’s constant
  
  

• NA = 6.022 × 10²³ mol⁻¹ → Avogadro’s number
  
  

• R = 0.0821 L·atm/mol·K = 8.314 J/mol·K → Gas constant
  
  

Mole Relationships

• n = m / M
  
  

  • n = moles
    
    

  • m = mass (g)
    
    

  • M = molar mass (g/mol)
    
    

  • Ex: 36 g H₂O ÷ 18 g/mol = 2.0 mol
    
    

• N = n × NA
  
  

  • N = particles
    
    

  • n = moles
    
    

  • NA = Avogadro’s number
    
    

  • Ex: 0.25 mol Ne × 6.022×10²³ = 1.51×10²³ atoms
    
    

Isotopes

• Average atomic mass = Σ (fraction × isotope mass)
  
  

  • Ex: 40% X-10, 60% X-12 → 11.2 amu
    
    

Electrostatics

• F = k (q₁ q₂) / r²
  
  

  • F = force (N)
    
    

  • q₁, q₂ = charges (C)
    
    

  • r = separation (m)
    
    

  • k = Coulomb’s constant
    
    

  • Ex: q₁=+1 μC, q₂=−2 μC, r=0.30 m → F = 0.20 N (attractive)
    
    

Light & Photons

• λν = c
  
  

  • λ = wavelength (m)
    
    

  • ν = frequency (Hz, s⁻¹)
    
    

  • c = 3.00×10⁸ m/s
    
    

• E = hν = hc/λ
  
  

  • E = energy (J)
    
    

  • h = 6.626×10⁻³⁴ J·s
    
    

  • ν = frequency (Hz)
    
    

  • λ = wavelength (m)
    
    

  • c = 3.00×10⁸ m/s
    
    

  • Ex: λ=400 nm → E=4.97×10⁻¹⁹ J
    
    

Quantum Numbers

• n = principal shell
  
  

• l = subshell (0=s, 1=p, 2=d, 3=f)
  
  

• mₗ = orbital orientation
  
  

• mₛ = spin (+½ or −½)
  
  

• Ex: last electron of P (3p³): n=3, l=1, mₗ=0, mₛ=+½
  
  

Gas Law

• PV = nRT
  
  

  • P = pressure
    
    

  • V = volume
    
    

  • n = moles
    
    

  • R = gas constant
    
    

  • T = temp (K)
    
    

  • Ex: 1.0 mol at 1 atm, 273 K → V = 22.4 L
    
    

Trends

• Zeff = Z – shielding
  
  

• IE (ionization energy): ΔE to remove e⁻ (always +, endothermic)

• EA (electron affinity): ΔE when atom gains e⁻ (often −, exothermic for halogens)