Chemical Reactions: Types, Calculations, and Stoichiometry

Basics of Chemical Reactions

Physical vs. Chemical Changes

Chemical reactions involve the transformation of one or more substances into new substances with different properties. Physical changes, in contrast, do not alter the chemical identity of a substance.

• Physical Change: Alters the physical state (e.g., melting, boiling) without changing composition.

• Chemical Change: Involves bond breaking and formation, resulting in new substances.

• Indicators of Chemical Change: Color change, formation of a precipitate, gas production, heat/flame emission or absorption.

Example: The tarnishing of silver (Ag) to silver sulfide (Ag2S) is a chemical change.

Chemical Equations

A chemical equation uses formulas and symbols to represent the reactants and products in a reaction. It must be balanced to obey the Law of Conservation of Mass.

• Reactants: Substances present before the reaction (left side).

• Products: Substances formed by the reaction (right side).

• Coefficients: Numbers in front of formulas indicating the number of molecules or moles.

• Physical States: Indicated as (s), (l), (g), or (aq).

Law of Conservation of Mass: Atoms are neither created nor destroyed in a chemical reaction; equations must be balanced using whole-number coefficients.

Types of Chemical Reactions

Combination (Synthesis) Reactions

Two or more elements or compounds combine to form a single product.

• General Form: A + B → AB

• Example: 2Mg(s) + O2(g) → 2MgO(s)

Decomposition Reactions

A single compound breaks down into two or more simpler substances.

• General Form: AB → A + B

• Example: 2HgO(s) → 2Hg(l) + O2(g)

Single Replacement Reactions

One element replaces another in a compound.

• General Form: A + BC → B + AC

• Example: Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)

Double Replacement Reactions

Two compounds exchange ions to form two new compounds.

• General Form: AB + CD → AD + CB

• Example: ZnS(s) + 2HCl(aq) → ZnCl2(aq) + H2S(g)

Summary Table: Types of Reactions

Type of Reaction	Examples	Comment
Single replacement	2NaCl + Br2 → 2NaBr + Cl2 Fe + CuSO4 → FeSO4 + Cu	Element replaces another in a compound
Double replacement	AgNO3 + NaCl → AgCl + NaNO3 HCl + NaOH → H2O + NaCl	Ions exchange between compounds

The Mole and Avogadro’s Number

Definition and Use

The mole (mol) is a counting unit in chemistry, representing 6.02 × 1023 particles (Avogadro’s Number, NA).

• 1 mole of atoms, molecules, ions, or formula units = 6.02 × 1023 entities

• Used to relate macroscopic amounts to the number of particles

Example: 1 mole of CO2 molecules = 6.02 × 1023 CO2 molecules

Using Avogadro’s Number in Calculations

• To convert moles to particles: multiply by NA

• To convert particles to moles: divide by NA

Example Calculation: How many molecules are in 5.0 moles of CO2? $5.0 \text{ mol} \times 6.02 \times 10^{23} \text{ molecules/mol} = 3.01 \times 10^{24} \text{ molecules}$

Subscripts and Moles in Chemical Formulas

Subscripts in a chemical formula indicate the number of atoms of each element in one molecule and the mole ratio in one mole of the compound.

Mass to Mole Conversions

Molar Mass

The molar mass is the mass of one mole of a substance, expressed in grams per mole (g/mol). It is numerically equal to the formula weight in atomic mass units (amu).

• Formula: $\text{Molar mass (g/mol)} = \text{Sum of atomic masses of all atoms in the formula}$

Relating Grams to Moles

• To convert grams to moles: divide by molar mass

• To convert moles to grams: multiply by molar mass

Example: How many moles are in 36.5 g of CaCO3? $\text{Moles} = \frac{36.5 \text{ g}}{100.09 \text{ g/mol}} = 0.365 \text{ mol}$

Mole Calculations in Chemical Equations

Stoichiometry

Balanced chemical equations provide the mole ratios needed to relate reactants and products. Calculations often require converting grams to moles, using the mole ratio, and then converting back to grams if needed.

• Example: Using the equation 2C2H6 + 5O2 → 4CO + 6H2O, how many moles of CO are produced from 3.5 moles of C2H6? $3.5 \text{ mol C}_2\text{H}_6 \times \frac{4 \text{ mol CO}}{2 \text{ mol C}_2\text{H}_6} = 7.0 \text{ mol CO}$

Limiting Reactants

Concept and Identification

The limiting reactant is the substance that is completely consumed first, thus limiting the amount of product formed. The other reactant is in excess.

• To identify the limiting reactant, calculate the amount of product formed from each reactant; the one producing less product is limiting.

Theoretical Yield and Percent Yield

Definitions

• Theoretical Yield: Maximum amount of product predicted by stoichiometry.

• Actual Yield: Amount of product actually obtained from the reaction.

• Percent Yield: Measures reaction efficiency.

Formula: $\text{Percent yield} = \frac{\text{actual yield (g)}}{\text{theoretical yield (g)}} \times 100\%$

Example: If the theoretical yield is 23 g and the actual yield is 15 g, then: $\text{Percent yield} = \frac{15}{23} \times 100\% = 65.2\%$

Summary Table: Reaction Types and Examples

Type	General Equation	Example
Combination	A + B → AB	2Na + Cl2 → 2NaCl
Decomposition	AB → A + B	2KClO3 → 2KCl + 3O2
Single Replacement	A + BC → B + AC	Zn + 2HCl → ZnCl2 + H2
Double Replacement	AB + CD → AD + CB	AgNO3 + NaCl → AgCl + NaNO3