Chemical Reactions and Equations: Energy, Balancing, and Types

Chemical Reactions: Fundamentals

Collision Theory and Activation Energy

Chemical reactions occur when reactant molecules collide with sufficient energy to break existing bonds and form new ones. The minimum energy required for a reaction to proceed is called activation energy.

• Collision Theory: Reactants must collide with enough energy and proper orientation for a reaction to occur.

• Activation Energy ($E_a$): The energy barrier that must be overcome for reactants to transform into products.

• Energy Diagram: Shows the energy changes during a reaction, including activation energy and the energy difference between reactants and products.

Example: The reaction between nitrogen and oxygen to form nitrogen monoxide requires sufficient energy to break the N≡N and O=O bonds.

Energy Changes in Chemical Reactions

Exothermic vs. Endothermic Reactions

Chemical reactions can either release or absorb energy, typically in the form of heat.

• Exothermic Reaction: Releases heat; the energy of the products is lower than that of the reactants. Heat is a product.

• Endothermic Reaction: Absorbs heat; the energy of the products is higher than that of the reactants. Heat is a reactant.

Examples:

• Exothermic: $\mathrm{C(s) + O_2(g) \rightarrow CO_2(g) + 394\ kJ}$

• Endothermic: $\mathrm{N_2(g) + O_2(g) + 43\ kJ \rightarrow 2NO(g)}$

Additional info: Energy diagrams for exothermic reactions show the products at a lower energy level than reactants, while endothermic diagrams show the opposite.

Law of Conservation of Mass and Chemical Equations

Law of Mass Conservation

The Law of Conservation of Mass states that matter is neither created nor destroyed in a chemical reaction. The total mass of reactants equals the total mass of products.

• Each element must have the same number of atoms on both sides of the equation.

• Balancing equations ensures this law is obeyed.

Law of Constant Composition

The Law of Constant Composition states that a chemical compound always contains the same proportion of elements by mass.

• Never change the subscripts in a chemical formula when balancing equations; only adjust coefficients.

Symbols Used in Chemical Equations

Chemical equations use specific symbols to indicate the physical state and reaction conditions.

Symbol	Meaning
(s)	Solid
(l)	Liquid
(g)	Gas
(aq)	Aqueous (dissolved in water)
Δ	Heat is added
→	Yields/products

Balancing Chemical Equations

Steps to Balance Equations

Balancing chemical equations ensures the same number of each atom on both sides of the reaction.

• List the number of atoms of each element in reactants and products.

• Adjust coefficients to balance atoms; do not change subscripts.

• Start with the most complex molecule or the element that appears in the fewest compounds.

• Check your work by counting atoms again.

Example:

• $\mathrm{N_2 + O_2 \rightarrow N_2O}$ (Not balanced)

• $\mathrm{N_2 + O_2 \rightarrow 2NO}$ (Balanced)

Types of Chemical Reactions

Classification of Reaction Types

Chemical reactions are classified based on how reactants transform into products.

• Combination Reaction: Two or more reactants combine to form one product. Example: $\mathrm{2Mg(s) + O_2(g) \rightarrow 2MgO(s)}$

• Decomposition Reaction: One substance splits into two or more simpler substances. Example: $\mathrm{2KClO_3(s) \rightarrow 2KCl(s) + 3O_2(g)}$

• Single Replacement Reaction: One element replaces another in a compound. Example: $\mathrm{Zn(s) + 2HCl(aq) \rightarrow ZnCl_2(aq) + H_2(g)}$

• Double Replacement Reaction: Two elements in different compounds exchange places. Example: $\mathrm{AgNO_3(aq) + NaCl(aq) \rightarrow AgCl(s) + NaNO_3(aq)}$

Combustion Reactions

Combustion is a special type of combination reaction involving oxygen and producing heat and light.

• General pattern: Hydrocarbon + O2 → CO2 + H2O

• Example: $\mathrm{CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(g)}$

Oxidation-Reduction (Redox) Reactions

Definition and Electron Transfer

Oxidation-reduction reactions involve the transfer of electrons between substances.

• Oxidation: Loss of electrons; increase in oxidation state.

• Reduction: Gain of electrons; decrease in oxidation state.

Example: $\mathrm{4Fe(s) + 3O_2(g) \rightarrow 2Fe_2O_3(s)}$ Fe is oxidized, O is reduced.

Oxidation States

Oxidation state (or oxidation number) is a value assigned to an atom to indicate its degree of oxidation or reduction.

• Element in its free form: oxidation state = 0

• Oxygen in compounds: usually -2

• Hydrogen in compounds: usually +1

Example: $\mathrm{Zn(s) + 2HCl(aq) \rightarrow ZnCl_2(aq) + H_2(g)}$ Zn is oxidized (0 to +2), H is reduced (+1 to 0).

Summary Table: Types of Chemical Reactions

Type	General Pattern	Example
Combination	$A + B \rightarrow AB$	$2H_2 + O_2 \rightarrow 2H_2O$
Decomposition	$AB \rightarrow A + B$	$2KClO_3 \rightarrow 2KCl + 3O_2$
Single Replacement	$A + BC \rightarrow AC + B$	$Zn + 2HCl \rightarrow ZnCl_2 + H_2$
Double Replacement	$AB + CD \rightarrow AD + CB$	$AgNO_3 + NaCl \rightarrow AgCl + NaNO_3$
Combustion	Hydrocarbon + $O_2 \rightarrow CO_2 + H_2O$	$CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$

Important Reminders for Balancing Equations

• Do not change subscripts in chemical formulas.

• Adjust only coefficients to balance equations.

• Balance polyatomic ions as single units if they appear unchanged on both sides.

• Multiply all coefficients to get the simplest whole number ratio.

Additional info: Practice problems and "Learning Check" questions throughout the notes reinforce understanding of reaction types, balancing, and energy changes.