Chemical Reactions and Equations

Chemical Equations

• Chemical equations use chemical symbols to represent reactants and products.
• They illustrate the relative amounts of reactants and products.
• A skeletal equation shows the chemical formulas of reactants and products (e.g., $H_2 + O_2 \rightarrow H_2O$).
• A balanced equation demonstrates the correct and relative amounts of reactants and products (e.g., $2H_2 + O_2 \rightarrow 2H_2O$).

Stoichiometry

• Stoichiometry relates the quantities of reactants and products to each other.
• It adheres to the Law of Conservation of Mass.

Reaction Types

• Combination Reaction: Two or more substances combine to form a single product (A + B → C; e.g., $S + O_2 \rightarrow SO_2$).
• Decomposition Reaction: A single compound breaks down into two or more products (C → A + B; e.g., $CaCO_3 \rightarrow CaO + CO_2$).
• Combustion Reaction: A rapid reaction that produces a flame, often involving oxygen (e.g., $CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$).

Formula Weights

• Formula weight is the sum of the atomic weights of each atom in a chemical formula.
• It's also known as molecular weight when referring to molecules (e.g., $H_2SO_4$ = 2(1.0 amu) + (32.1 amu) + 4(16.0 amu) = 98.1 amu).

Percent Composition

• Percent composition indicates the percentage by mass of each element in a substance.
• The formula is: % Element = $\frac{(number , of , atoms)(atomic , weight , of , element)}{(formula , weight , of , compound)} \times 100$
• For example, in $H_2O$: %H = 11% and %O = 89%.

Avogadro's Number and the Mole

• Mole: a fixed number of particles, specifically, 1 mol = $6.022142 \times 10^{23}$ particles.
• Avogadro's Number ($N_A$) is approximately $6.022 \times 10^{23}$.
• Molar mass is the mass of 1 mole of a substance.
• Molar mass is numerically equal to the formula weight, expressed in g/mol (e.g., 1 mol $H_2O$ = 18.0 g $H_2O$).

Molar Mass as a Conversion Factor

• Molar mass is used to convert between mass and moles.
• Grams to Moles: grams $\times$ $\frac{1 , mol}{# , grams}$ = moles
• Moles to Grams: moles $\times$ $\frac{# , grams}{1 , mol}$ = grams

Empirical Formulas

• Empirical formulas represent the smallest whole number ratio of atoms in a substance (e.g., $HO$).
• To determine empirical formulas: assume 100 g sample, convert grams to moles, divide each mole quantity by the smallest mole quantity, and multiply by an integer if necessary to get the smallest whole number ratio.

Molecular Formulas

• Molecular formulas indicate the actual number of atoms of each element in a molecule (e.g., $H_2O_2$).
• The molecular formula is a whole number multiple of the empirical formula, where Molecular weight = n $\times$ Empirical Formula weight.

Stoichiometric Coefficients

• Stoichiometric coefficients in a balanced equation indicate the relative amounts of reactants and products.
• Coefficients can be interpreted as the relative number of moles of each substance involved in the reaction (e.g., in $2H_2(g) + O_2(g) \rightarrow 2H_2O(g)$, 2 mol $H_2$ reacts with 1 mol $O_2$ to produce 2 mol $H_2O$).

Limiting Reactants

• The limiting reactant is the reactant that is completely consumed in a reaction.
• It determines the amount of product that can be formed.
• Reactants that are not completely consumed are in excess.

Theoretical Yield

• Theoretical yield is the quantity of product calculated to form when all of the limiting reactant reacts.

Actual Yield

• Actual yield is the amount of product actually obtained in a reaction.
• Actual yield is always less than the theoretical yield.

Percent Yield

• Percent yield relates the actual yield to the theoretical yield using the formula:
• $% Yield = \frac{Actual , Yield}{Theoretical , Yield} \times 100$