Chapter 7: Chemical Reactions – Study Notes

Chapter 7: Chemical Reactions

Chemical Reactions

Chemical reactions describe the process by which substances (reactants) are transformed into new substances (products). Chemical equations are used to represent these changes, showing the reactants, products, and their physical states.

• Reactants are the starting substances in a chemical reaction.

• Products are the substances formed as a result of the reaction.

• The arrow (→ or =) means "goes to," "yields," or "produces."

• A plus sign (+) separates multiple reactants or products.

• The double arrow (⇌) indicates equilibrium, meaning the reaction is reversible.

• Physical states are indicated as follows: (aq) for aqueous, (g) for gas, (s) for solid or precipitate, (l) for liquid.

Balancing Equations

Balancing chemical equations ensures the law of conservation of mass is obeyed, meaning the number of atoms of each element is the same on both sides of the equation.

• Conservation of mass: No atoms are lost or gained during a chemical reaction.

• Do not change the subscripts in chemical formulas, as this alters the chemical identity.

• Use coefficients to balance equations.

• Example: $\_\mathrm{H}_2(g) + \_\mathrm{O}_2(g) \rightarrow \_\mathrm{H}_2\mathrm{O}(l)$

Balancing Reactions

Balancing is typically done by inspection, following a systematic approach.

• Write the unbalanced equation first.

• Leave counting elements (like Fe, O2, H2) until last.

• Count polyatomic ions (e.g., NO3-, SO42-) as units if they appear unchanged on both sides.

• Sometimes, subscripts can suggest coefficients needed for balancing.

• Example: $\_\mathrm{N}_2(g) + \_\mathrm{H}_2(g) \rightarrow \_\mathrm{NH}_3(g)$

Balancing Examples

Practice balancing the following reactions:

• $\_\mathrm{Fe} + \_\mathrm{Cl}_2 \rightarrow \_\mathrm{FeCl}_3$

• $\_\mathrm{Zn} + \_\mathrm{HCl} \rightarrow \_\mathrm{ZnCl}_2 + \_\mathrm{H}_2$

• $\_\mathrm{Ca}_3\mathrm{P}_2 + \_\mathrm{H}_2\mathrm{O} \rightarrow \_\mathrm{Ca(OH)}_2 + \_\mathrm{PH}_3$

Types of Chemical Reactions

Chemical reactions are classified into several types based on the changes that occur:

• Dissociation and Ionization: Formation of ions in water.

• Combination: Two or more reactants form a single product.

• Combustion: Rapid reaction with oxygen producing heat and light.

• Decomposition: A single compound breaks down into simpler substances.

• Single Replacement: One element replaces another in a compound.

• Double Replacement: Exchange of ions between two compounds.

Dissolution

Dissolution: Forming Ions in Water

Dissolution refers to the process of a substance dissolving in water to form ions. A substance is considered soluble if a significant amount dissolves.

• A solution is a homogeneous mixture of a solute (the dissolved substance) and a solvent (the substance doing the dissolving, often water).

Dissociation and Ionization

Ionic compounds dissociate in water to produce ions, while acids ionize to form ions. Molecular compounds may simply dissolve without forming ions.

• Dissociation: $\mathrm{NaCl}(s) \rightarrow \mathrm{Na}^+(aq) + \mathrm{Cl}^-(aq)$

• $\mathrm{Na}_2\mathrm{SO}_4(s) \rightarrow 2\mathrm{Na}^+(aq) + \mathrm{SO}_4^{2-}(aq)$

• Ionization (for acids): $\mathrm{HNO}_3(g) \rightarrow \mathrm{H}^+(aq) + \mathrm{NO}_3^-(aq)$

• Molecular compounds: $\mathrm{C}_{12}\mathrm{H}_{22}\mathrm{O}_{11}(s) \rightarrow \mathrm{C}_{12}\mathrm{H}_{22}\mathrm{O}_{11}(aq)$

• Aqueous (aq): Indicates ions or molecules are dissolved in water.

Saturated Solutions and Equilibrium

When more solute is added than can dissolve, a saturated solution is formed and equilibrium is established between dissolved and undissolved solute.

• Saturated solution: Maximum amount of compound dissolved.

• Example equilibrium: $\mathrm{NaCl}(s) \rightleftharpoons \mathrm{Na}^+(aq) + \mathrm{Cl}^-(aq)$

Dissolution Examples

Equations for dissolution of common compounds:

• $\mathrm{BaCl}_2 \rightarrow \mathrm{Ba}^{2+} + 2\mathrm{Cl}^-$

• $\mathrm{HF} \rightarrow \mathrm{H}^+ + \mathrm{F}^-$

• $\mathrm{NH}_4\mathrm{OH} \rightarrow \mathrm{NH}_4^+ + \mathrm{OH}^-$

Key Terms and Concepts

• Solute: The substance that is dissolved in a solution.

• Solvent: The substance that dissolves the solute (commonly water).

• Dissociation: The process by which ionic compounds separate into ions in water.

• Ionization: The process by which acids form ions in water.

• Saturated Solution: A solution containing the maximum amount of dissolved solute at a given temperature.

• Equilibrium: The state in which the rate of dissolution equals the rate of precipitation.

Example Table: Physical States in Chemical Equations

Symbol	Meaning
(aq)	Aqueous (dissolved in water)
(g)	Gas
(s)	Solid or precipitate
(l)	Liquid

Example Table: Types of Chemical Reactions

Type	Description	General Equation
Combination	Two or more reactants form one product	$A + B \rightarrow AB$
Decomposition	One compound breaks into two or more products	$AB \rightarrow A + B$
Single Replacement	One element replaces another in a compound	$A + BC \rightarrow AC + B$
Double Replacement	Exchange of ions between two compounds	$AB + CD \rightarrow AD + CB$
Combustion	Reaction with oxygen producing heat and light	$C_xH_y + O_2 \rightarrow CO_2 + H_2O$

Example: Dissociation of sodium sulfate in water: $\mathrm{Na}_2\mathrm{SO}_4(s) \rightarrow 2\mathrm{Na}^+(aq) + \mathrm{SO}_4^{2-}(aq)$

Example: Combination reaction of aluminum and bromine: $2\mathrm{Al}(s) + 3\mathrm{Br}_2(l) \rightarrow 2\mathrm{AlBr}_3(s)$

Example: Combustion of methane: $\mathrm{CH}_4(g) + 2\mathrm{O}_2(g) \rightarrow \mathrm{CO}_2(g) + 2\mathrm{H}_2O(g)$

Additional info: These notes expand on the brief points in the slides, providing definitions, examples, and tables for clarity and completeness. The content is suitable for introductory college chemistry students preparing for exams.