CHEM 150 Midterm I

Engineering Chemistry · Camosun College · Units 17 Reference

Unit 1 Matter & Measurement

Classification, atomic structure, stoichiometry, nomenclature

Classification of Matter
  • Element pure substance; cannot be separated by chemical means
  • Compound two+ elements, fixed ratio; separated only by chemical means
  • Homogeneous mixture uniform (e.g., salt water, glass, air)
  • Heterogeneous mixture non-uniform (e.g., drywall, sandy water)
Tip: Physical changes (melting, dissolving) preserve chemical identity. Chemical changes (burning, rusting) produce new substances.
Dalton's Atomic Theory (1808)
  1. Matter composed of indivisible atoms
  2. All atoms of an element are identical
  3. Compounds = atoms in fixed whole-number ratios
  4. Chemical reactions = rearrangement of atoms; atoms not created/destroyed
Note: Isotopes violate postulate 2; nuclear reactions violate postulate 4.
Atomic Structure
ParticleChargeMass (u)Location
Proton+11.0073Nucleus
Neutron01.0087Nucleus
Electron10.000549Orbitals

Atomic number Z = # protons  |  Mass number A = protons + neutrons  |  Isotopes = same Z, different A

Average Atomic Mass
avg mass = Σ (fractional abundance × isotope mass) e.g., Cl: 0.7577×34.97 + 0.2423×36.97 = 35.45 u
Stoichiometry Essentials
n = m / M     (mol = g ÷ g/mol) N = n × N   (N = 6.022×10²³) M = mol × M   (g = mol × g/mol)

Limiting reagent: convert all reactants to same product; smallest amount is the limit.

% yield = (actual / theoretical) × 100% M = n / V   (mol/L)    MᵢVᵢ = MᶠVᶠ
Nomenclature Quick Reference
  • Ionic: cation name + anion name; variable-charge metals use Roman numerals (Fe² = iron(II))
  • Oxyanions: "Nick the Camel" NO nitrate, NO nitrite, SO² sulfate, SO² sulfite, ClO perchlorate, ClO chlorate, ClO chlorite, ClO hypochlorite
  • Molecular: Greek prefixes (mono, di, tri, tetra, penta )
  • Acids: H + anion; -ide hydro___ic acid; -ate ___ic acid; -ite ___ous acid
Metal Types & Periodic Table Groups

Diatomic elements: H, N, O, F, Cl, Br, I (mnemonic: HOFBrINCl)

Metals (left/centre): ductile, malleable, lustrous, conduct heat & electricity.  Non-metals (right): brittle, poor conductors.  Metalloids: along staircase (B, Si, Ge, As, Sb, Te).

Variable-charge metals: Cu (1+, 2+) Fe (2+, 3+) Co (2+, 3+) Mn (2+, 7+) Sn (2+, 4+)  | Fixed: Na, Mg², Al³, Ca², Ba², K

Unit 2 Chemical Bonding

Bond types, periodic trends, ionization energy, lattice energy, Born-Haber

Types of Chemical Bonds
  • Ionic complete electron transfer; metal + nonmetal; e.g., NaCl. Electrostatic attraction of oppositely charged ions.
  • Covalent electron sharing; nonmetal + nonmetal; e.g., H, CO. Can be polar or nonpolar.
  • Metallic metal cations in a delocalized "sea" of electrons; explains conductivity, malleability.
ΔEN rule: 0 = nonpolar covalent | 01.5 = polar covalent | >1.5 ionic character
Periodic Trends Summary
Ionization Energy (I)
Increases leftright across period
Decreases topbottom in group

Exceptions: B < Be (2p vs 2s); O < N (paired e repulsion in O)

Atomic Radius
Increases topbottom (more shells)
Decreases leftright (Zeff increases)
Electron Affinity
More negative going right and up

More negative = more exothermic = "wants" electrons more. Exceptions: N (half-filled 2p stable)

Effective Nuclear Charge & Ionic Radii
Slater's Rule (simplified)
Z_eff = Z S    (S = # core electrons shielding valence shell) Examples: C (Z=6, 2 core e) Z_eff = 4 N (Z=7, 2 core e) Z_eff = 5 B (Z=5, 2 core e) Z_eff = 3

Ionic radius: Cations are smaller than parent atom (lost electrons, more Zeff per electron). Anions are larger (gained electrons, less Zeff per electron).

Isoelectronic series (same # electrons): more protons smaller. E.g., Al³ < Na < F < O² (all have 10 electrons).

Lattice Energy

Energy released when gaseous ions form 1 mol of ionic crystal. Always negative (exothermic).

Factors: Lattice energy with higher ionic charges and smaller ions (shorter distance).

CompoundLattice E (kJ/mol)mp (°C)
LiF1049848
NaCl787801
KBr691734
MgO37952825
AlO159162054
Born-Haber Cycle (NaCl example)
Na(s) Na(g) +107 kJ (sublimation) Na(g) Na(g) + e +496 kJ (IE) ½Cl(g) Cl(g) +121 kJ (½ bond energy) Cl(g) + e Cl(g) 349 kJ (EA) Na(g) + Cl(g) NaCl(s) 787 kJ (lattice E) Na(s) + ½Cl(g) NaCl(s) 412 kJ (ΔH°f)

Use Hess's Law: ΔH°f = sublimation + IE + ½ bond + EA + lattice E

IE Trend Practice (memorize order)

Increasing IE:  Sr < Ca < Se < Br  (same group ordering + period position)

Increasing ionic size:  Ti < K < S² < Se²  (isoelectronic Ti, K, S² all have 18 e; Se² has 36 e = bigger shell)

Increasing atomic radius:  Br < Se < Ca < Sr

Unit 3 Lewis Structures

Drawing rules, formal charges, resonance, exceptions to octet

Lewis Structure Drawing Steps
  1. Count total valence electrons (VE). For anions add n e; for cations subtract n e.
  2. Choose central atom: usually least electronegative (but not H or F). In oxoanions, the non-O atom is central.
  3. Place one bond between central atom and each terminal atom.
  4. Complete octets on terminal atoms first using lone pairs.
  5. Place remaining electrons on central atom.
  6. If central atom is short of octet, convert lone pairs on terminals to double/triple bonds.
  7. Verify: total electrons used = VE counted in step 1.
Quick VE check: VE = (sum of group numbers) ± charge. E.g., NO = 5+6+6+1 = 18 VE
Formal Charge Formula
FC = (valence e) ½(bonding e) (lone pair e) e.g., in NNO (best NO structure): Left N: 5 ½(6) 2 = 0 Right N: 5 ½(6) 0 = +2? check each structure

Best Lewis structure:

  1. Lowest formal charges (ideally zero)
  2. Negative FC on most electronegative atom
  3. Same-sign FCs on adjacent atoms = very bad
Resonance

When two or more valid Lewis structures can be drawn, resonance exists. The true structure is the resonance hybrid intermediate between all structures.

Key rule: Only electrons move in resonance. Atoms never move. Use double-headed arrow () between structures.

Equivalent resonance: NO, NO, CO² (all structures identical by symmetry contribute equally)

Non-equivalent resonance: NO structures NOT identical; best (lowest FC) contributes most to hybrid.

Exceptions to the Octet Rule

1. Odd-electron species (radical)

  • NO, NO odd total VE, always one atom with 7e
  • Very reactive; paramagnetic

2. Incomplete octets

  • BF (6e on B), BeCl (4e on Be)
  • Lewis acids can accept a lone pair

3. Expanded valence shell

  • Only elements in period 3+ (Si, P, S, Cl, As, Se, Br, I, Xe)
  • PCl (10e), SF (12e), XeF (12e)
  • SO², ClO expanded shells minimize formal charges
Never give C, N, O, F more than 8 valence electrons.
Bond Dissociation Energies (D°, kJ/mol)
BondD° (kJ/mol)BondD° (kJ/mol)BondD° (kJ/mol)
HH436CC347NN946
CH414C=C611O=O498
CO360CC837HO464
C=O799CN305HN391
ΔH reaction using bond energies
ΔH = Σ D°(bonds broken) Σ D°(bonds formed) (bonds broken = reactants; bonds formed = products)

Unit 4 Polarity & VSEPR

Electronegativity, dipole moments, molecular geometry

Electronegativity & Bond Polarity

Electronegativity (EN) = ability of atom in molecule to attract shared electrons (Pauling scale).

Increases up and to the right on periodic table. F is most electronegative (4.0).

μ = Q × r    (dipole moment, debyes D) 1 D = 3.34×10³ C·m % ionic character = (measured μ / theoretical μ) × 100%
Polarity: Bond dipoles must add as vectors. CO has polar bonds but is nonpolar (linear, dipoles cancel). HO has polar bonds AND is polar (bent, net dipole).
VSEPR Algorithm
  1. Draw Lewis structure
  2. Count Electron Groups (EG) around central atom: each single bond, double bond, triple bond, and lone pair = 1 EG
  3. Determine Electron Group Geometry (EGG) from # EG
  4. Determine Molecular Shape from bonding pairs only
  5. Determine if molecule is polar (net dipole 0)
Lone pairs repel more than bonding pairs bond angles are compressed. LPLP > LPBP > BPBP repulsion.
VSEPR Shape Reference Table
EGEGGLPMolecular ShapeAnglesExamplePolar?
2Linear0Linear180°CO, HCNNo / Yes
3Trig. planar0Trigonal planar120°BF, SONo
3Trig. planar1Bent (120°)~117°SO, NOYes
4Tetrahedral0Tetrahedral109.5°CH, CClNo
4Tetrahedral1Trig. pyramidal107°NH, PHYes
4Tetrahedral2Bent (109°)104.5°HO, HSYes
5Trig. bipyramidal0Trig. bipyramidal90°/120°PCl, PFNo
5Trig. bipyramidal1See-saw~90°/120°SFYes
5Trig. bipyramidal2T-shaped~90°BrF, IFYes
5Trig. bipyramidal3Linear180°XeF, INo
6Octahedral0Octahedral90°SFNo
6Octahedral1Square pyramidal~90°BrFYes
6Octahedral2Square planar90°XeFNo
TBP lone pairs: Always go equatorial (120° from each other) because equatorial sites have fewer 90° neighbors than axial sites.

Unit 5 Intermolecular Forces

London dispersion, dipole-dipole, hydrogen bonding, effects on properties

Types of IMFs (weakest strongest)
ForceBetweenStrengthExample
London DispersionALL molecules (induced dipoles)Weakest; increases with molar mass & surface areaHe, Ar, CH, I
DipoleDipolePolar moleculesModerate; depends on polarityHClHCl, acetone
Hydrogen BondingH bonded to F, O, or N near another F/O/NStrong (1540 kJ/mol)HO, NH, HF, alcohols
IonDipoleIons + polar solventsStrongest of Van der WaalsNa in water
Remember: IMFs are 10100× weaker than covalent bonds. Boiling water breaks H-bonds between HO molecules NOT the OH bonds within them.
Effects on Physical Properties
  • Boiling point: with stronger/more IMFs. Same mass polar > nonpolar. Same polarity heavier (more surface area) = higher bp.
  • Vapour pressure: with stronger IMFs (harder to escape)
  • Viscosity: with stronger IMFs; with higher temperature
  • Surface tension: net inward force on surface molecules; with stronger IMFs
BP order: He(269°C) < O(183°C) < Cl(34°C) < acetone(56°C) (London disp. only) (London) (London) (London + dipole-dipole)
H-Bonding Special Cases
  • Ice floats: H-bonding creates rigid open lattice ice less dense than liquid water (4°C = maximum density)
  • Anomalous bp: HO (100°C), HF (19°C), NH (33°C) are all higher than expected from their group trend
  • DNA: H-bonds hold the double helix together
  • Kevlar: strength comes from H-bonding between amide groups
H-bonding requires: H directly bonded to F, O, or N. The H must be attracted to a nearby F, O, or N on another molecule.
"Like Dissolves Like" Solubility

Polar solvents dissolve polar/ionic solutes. Nonpolar solvents dissolve nonpolar solutes.

Soluble in water (polar):

  • NH H-bonds with water
  • HF, HCl polar, H-bond/dipole
  • SO polar molecule
  • CHOH OH group H-bonds
  • NaCl, ionic compounds (if lattice E < solvation E)

Insoluble in water:

  • CH (acetylene) nonpolar
  • CCl nonpolar (symmetric)
  • Octane, hexane nonpolar
  • S nonpolar
  • MgO lattice E too large

Unit 6 Solutions & Colligative Properties

Raoult's Law, Henry's Law, vapour pressure, osmosis

Raoult's Law Vapour Pressure Lowering
P_A = x_A × P°_A ΔP = x_solute × P°_A x_A = mole fraction of A = n_A / n_total

Example: 1 mol sucrose + 15 mol water. x_water = 15/16 = 0.9375. P = 0.9375 × 31.26 mbar = 29.3 mbar (pure water = 31.26 mbar).

Benzene/toluene mixture (0.5 mol each): Total P = x_benz×P°_benz + x_tol×P°_tol. Vapour is enriched in the more volatile (higher vapour pressure) component.

Henry's Law Gas Solubility
S_g = k × P_g S_g = solubility of gas (mol/L) k = Henry's constant (mol/(L·kPa) or mol/(L·atm)) P_g = partial pressure of gas above liquid

Key facts:

  • Gas solubility increases with pressure
  • Gas solubility decreases with temperature (carbonated drinks more bubbly when warm)
  • Larger molecules (stronger London forces) higher solubility
Lake Nyos (1986): Deep lake CO suddenly outgassed under pressure release 1700 deaths.
Colligative Properties (depend on # particles only)
PropertyFormulaNotes
Vapour pressure loweringΔP = x_solute × P°Raoult's Law
Boiling point elevationΔT_b = K_b × m × im = molality (mol/kg)
Freezing point depressionΔT_f = K_f × m × iAntifreeze, salting roads
Osmotic pressureπ = MRTM = molarity, R = 8.314, T in K

van 't Hoff factor i: # particles per formula unit. NaCl i=2; MgCl i=3; glucose i=1.

Osmotic pressure example
NaCl at 0.15 M, 25°C: π = 2 × 0.15 × 8.314×10³ × 298 = 0.744 kPa×10³... Actually: π = MiRT = (0.15)(2)(0.08206)(298) = 7.33 atm enormous!
Osmosis: Solvent moves from low solute concentration to high solute concentration across semipermeable membrane. Isotonic = no net flow; hypertonic = cells shrink (crenation); hypotonic = cells burst (haemolysis).

Unit 7 Gases

Gas laws, ideal gas law, Dalton's Law, KMT, real gases

Pressure Units & Conversions
1 atm = 760 mmHg = 760 torr = 101.325 kPa = 14.7 psi 1 Pa = 1 N/m²     STP: 0°C (273.15 K), 1 atm, molar volume = 22.4 L/mol
Simple Gas Laws
Boyle's: PV = PV    (const T, n) Charles': V/T = V/T (const P, n; T in Kelvin!) Gay-Lussac: P/T = P/T (const V, n) Avogadro: V n (const T, P) Combined: PV/T = PV/T
Always convert T to Kelvin! T(K) = T(°C) + 273.15
Ideal Gas Law & Dalton's Law
Ideal Gas Law: PV = nRT R = 8.314 L·kPa/(mol·K) [or] R = 0.08206 L·atm/(mol·K) Dalton's Law: P_total = P + P + P + ... P_i = x_i × P_total (x_i = mole fraction) Molar mass from ideal gas: M = mRT / PV

Gas over water: P_gas = P_atm P_HO(vapour)

Kinetic-Molecular Theory (5 Postulates)
  1. Gases consist of a large number of molecules in continuous random motion
  2. Volume of individual molecules is negligible compared to total volume
  3. No attractive or repulsive forces between molecules
  4. Average kinetic energy is proportional to absolute temperature: KE_avg = ½mv² = (3/2)k_B T
  5. Collisions between molecules are perfectly elastic (energy is transferred but not lost)

At same T: all gases have the same average KE. Lighter molecules move faster (effuse/diffuse faster).

Effusion: escape through a tiny hole. Diffusion: spread through space. Both faster for lighter gases.

Real Gases van der Waals
[P + n²a/V²][V nb] = nRT a = correction for intermolecular attractions (reduces observed pressure) b = correction for molecular volume (reduces available volume)

Real gases deviate most from ideal at high pressure and low temperature (conditions that bring molecules closer together).

CO at 273K: ideal 756.6 kPa; van der Waals 727.1 kPa (a=364, b=0.0427)

Common Gas Calculation Tips
Units: Use R = 8.314 when P is in kPa; use R = 0.08206 when P is in atm. Never mix.

Limiting reagent + gas: Find limiting reagent first, then use moles of product to calculate volume with ideal gas law.

Mixture pressures: Calculate moles of each gas, find total moles, then P_i = n_i × RT/V for each component.

Molar mass of gas from density
M (g/mol) = ρ × R × T / P (ρ in g/L)

Key Formulas & Constants

Everything you need on exam day

Constants
N = 6.022 × 10²³ mol¹ 1 u = 1.660 × 10² kg R = 8.314 L·kPa/(mol·K) R = 0.08206 L·atm/(mol·K) STP molar volume = 22.4 L/mol 1 atm = 101.325 kPa = 760 torr
All Key Formulas
n = m/M MᵢVᵢ = MᶠVᶠ % yield = actual/theoretical × 100% FC = VE ½(BE) LP Z_eff = Z S (core electrons) μ = Qr (dipoles) ΔH = Σ D°(broken) Σ D°(formed) P_A = x_A × P°_A (Raoult) ΔP = x_solute × P°_A S_g = k × P_g (Henry) π = MRT (osmotic pressure) PV = nRT PV/T = PV/T P_total = Σ P_i P_i = x_i × P_total [P + n²a/V²][V nb] = nRT (van der Waals)
Common Ion Charges
IonChargeIonCharge
NH+1OH1
NO1NO1
SO²2SO²2
CO²2PO³3
ClO1ClO1
ClO1ClO1
MnO1CrO²2
Naming Acids Quick Chart
AnionAcid nameAnionAcid name
Cl (chloride)HCl hydrochloric acidNO (nitrate)HNO nitric acid
SO² (sulfate)HSO sulfuric acidNO (nitrite)HNO nitrous acid
PO³ (phosphate)HPO phosphoric acidClO (chlorate)HClO chloric acid
CO² (carbonate)HCO carbonic acidClO (hypochlorite)HClO hypochlorous acid
Stoichiometry Worked Example Limiting Reagent
Reaction: 2Nb + 5Cl 2NbCl Given: 22.2 g Nb (M=92.91) and excess Cl mol Nb = 22.2 / 92.91 = 0.2390 mol Nb mol NbCl = 0.2390 mol Nb × (2 mol NbCl / 2 mol Nb) = 0.2390 mol mass NbCl = 0.2390 × 270.17 g/mol = 64.6 g (theoretical yield) If actual yield = 55.4 g % yield = 55.4/64.6 × 100 = 85.7%