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 $CO<em>2$ and $H</em>2O$.
• 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.