Stoichiometry Notes

Chemical Equations

Chemical equations use symbols to represent chemical reactions.

  • Reactants are on the left side of the arrow.
  • Products are on the right side of the arrow.
  • Labels indicate physical states: (g) gas, (l) liquid, (s) solid, (aq) aqueous.

Balancing Chemical Equations

Equations must be balanced to obey the law of conservation of mass. Use stoichiometric coefficients to balance.

  • Balance elements that appear in only one reactant and one product first.
  • Balance elements in two or more reactants/products next.

Reaction Types

  • Combination: Two or more reactants form one product.
  • Decomposition: One reactant forms two or more products.
  • Combustion: Substance burns in oxygen, often producing CO2 and H2O.
  • Single displacement: One solid metal exchanges with another.
  • Double displacement (Metathesis): Anions in two ionic compounds exchange cations.
  • Neutralization: Acid + Base -> Salt + Water.
  • Condensation: Two molecules combine, releasing water.

Ionic Equations

  • Molecular Equation: Compounds represented as molecules.
  • Complete Ionic Equation: Compounds that exist as ions in solution are represented as ions.
  • Net Ionic Equation: Includes only species involved in the reaction; spectator ions are excluded.

Solubility Rules

  • Memorize solubility rules to determine if a compound is aqueous or solid.

Electrolytes and Nonelectrolytes

  • Electrolyte: Dissolves in water to conduct electricity (ionic).
  • Strong Electrolyte: Dissociates completely.
  • Weak Electrolyte: Produces ions upon dissolving but exists mainly as non-ionized molecules.
  • Nonelectrolyte: Dissolves in water but does not conduct electricity (molecular).

Precipitation Reactions

An insoluble product (precipitate) separates from the solution.

Acid-Base Reactions

  • Arrhenius Acid: Produces H^+ in water.
  • Arrhenius Base: Produces OH^- in water.
  • Brønsted Acid: Proton donor.
  • Brønsted Base: Proton acceptor.
  • Monoprotic Acid: Has one proton to donate.
  • Polyprotic Acid: Has more than one acidic hydrogen.
  • Neutralization: Reaction between acid and base, typically producing water and a salt. The net ionic equation of a strong acid–strong base reactions is: H^+(aq) + OH^–(aq) \longrightarrow H_2O(l)

Oxidation-Reduction (Redox) Reactions

Electrons are transferred between reactants.

  • Oxidation: Loss of electrons (OIL).
  • Reduction: Gain of electrons (RIG).
  • Reducing Agent: The compound that causes reduction.
  • Oxidizing Agent: The compound that causes oxidation.

Oxidation Numbers

Rules for assigning oxidation numbers:

  1. Elemental form: 0
  2. Monatomic ion: charge of the ion
  3. Group 1A metals: +1
  4. Group 2A metals: +2
  5. Hydrogen: +1 with nonmetals, -1 with metals
  6. Oxygen: -2 (except -1 in peroxides)
  7. Fluorine: -1
  8. Sum of oxidation numbers = charge on molecule/ion

Activity Series

List of metals in order of decreasing ease of oxidation. A metal will be oxidized by ions of any element below it.

Balancing Redox Equations

Use the half-reaction method:

  1. Balance the Electrons
  2. Make sure to have both mass balance and charge balance

Other Types of Redox Reactions

  • Combination
  • Decomposition
  • Disproportionation
  • Combustion

Mole and Chemical Reactions

Use mole ratios from balanced equations to convert between reactants and products.

Limiting Reactants

The reactant used up first. Calculate product formed from each reactant; the one producing less is the limiting reactant.

Reaction Yield

\% Yield = \frac{Actual Yield}{Theoretical Yield} \times 100

Gravimetric Analysis

Analytical technique based on mass measurement.

Molarity and Reactions in Aqueous Solution

Use molarity and stoichiometry to find volumes/moles of reactants and products.

Acid-Base Titrations

  • Titration: A volumetric technique that uses burettes.
  • Equivalence Point: The point in the titration where the acid has been neutralized.
  • Indicator: Signalled by a colour change.