Stoichiometry: Calculations with Chemical Formulas and Equations

Stoichiometry: Calculations with Chemical Formulas and Equations

Chemical Equations

Chemical equations are symbolic representations of chemical reactions, showing the reactants and products, their physical states, and the relative quantities involved. They are fundamental tools for describing chemical changes and for performing quantitative calculations in chemistry.

• Reactants are substances consumed in the reaction; products are substances formed.

• Chemical equations use formulas to represent substances and arrows to indicate the direction of the reaction.

• Coefficients in front of formulas indicate the relative number of molecules or moles of each substance involved.

• A balanced chemical equation has equal numbers of each type of atom on both sides, reflecting the Law of Conservation of Mass: Matter is neither created nor destroyed.

Coefficients vs. Subscripts

When balancing chemical equations, it is crucial to distinguish between coefficients and subscripts:

• Coefficients change the amount of a substance but not its identity.

• Subscripts define the identity of a compound; changing them alters the substance itself.

• Rule: Never change subscripts when balancing an equation.

• Example: H2O (water) vs. H2O2 (hydrogen peroxide) are different substances.

Balancing Chemical Equations

Balancing ensures the same number of each atom on both sides of the equation. The process involves adjusting coefficients only.

• Start by balancing atoms that appear in only one reactant and one product.

• Balance hydrogen and oxygen atoms last, as they often appear in multiple compounds.

• Example: Combustion of methane

CH4 + 2 O2 → CO2 + 2 H2O C: 1 | 1 H: 4 | 4 O: 4 | 4

• Check each element to confirm the equation is balanced.

Practice: Balancing Equations

Practice problems help reinforce the balancing process. For each equation, adjust coefficients to ensure atom balance.

• Example: 6 Li (s) + N2 (g) → 2 Li3N (s)

• Example: TiCl4 (l) + 2 H2O (l) → TiO2 (s) + 4 HCl (aq)

• Example: 2 NH4NO3 (s) + N2 (g) → 3 N2 (g) + O2 (g) + 2 H2O (l)

Chemical Reactivity: Types of Reactions

Chemical reactions can be classified into several types based on the rearrangement of atoms and molecules.

Combination (Synthesis) Reactions

• Two or more reactants combine to form a single product.

• General form: A + B → C

• Example: C (s) + O2 (g) → CO2 (g)

Decomposition Reactions

• A single reactant breaks apart to form two or more products.

• General form: C → A + B

• Example: 2 KClO3 (s) → 2 KCl (s) + 3 O2 (g)

Single Replacement (Displacement) Reactions

• An element replaces another element in a compound.

• General form: A + BX → AX + B

• Example: Fe (s) + CuSO4 (aq) → FeSO4 (aq) + Cu (s)

Double Replacement (Metathesis) Reactions

• Two compounds exchange ions to form two new compounds.

• General form: AY + BX → AX + BY

• Example: Na2CO3 (aq) + 2 AgNO3 (aq) → 2 NaNO3 (aq) + Ag2CO3 (s)

Combustion Reactions

• Rapid reactions with O2 that produce heat and light.

• Hydrocarbons combust to form CO2 and H2O.

• Example: CH4 (g) + 2 O2 (g) → CO2 (g) + 2 H2O (g)

*Additional info: The notes continue with stoichiometric calculations, formula weights, mole concept, and quantitative relationships, which are standard in General Chemistry and would be included in a full set of study notes for this chapter.*