Balancing Chemical Equations and Reaction Types

Balancing Equations & Conservation of Mass

• Balancing chemical equations ensures the same number of atoms of each element are on both sides of the equation.
• Balancing chemical equations follows the Law of Conservation of Mass.
• The Law of Conservation of Mass states that matter cannot be created or destroyed in a chemical reaction.
• The mass of reactants must equal the mass of the products.
• Steps for balancing equations:
  • Write the unbalanced equation.
  • Count atoms of each element on both sides.
  • Adjust coefficients to balance atoms.
  • Repeat until balanced.
  • Ensure the simplest coefficient ratio.

Classifying Reactions (5 Types)

• Five types of reactions include: synthesis, decomposition, single replacement, double replacement, and combustion.

Synthesis (Combination) Reaction

• Two or more reactants combine to form a single product (A + B → AB).
• Example: 2H₂ + O₂ → 2H₂O

Decomposition Reaction

• A single compound breaks down into two or more simpler products (AB → A + B).
• Example: 2H₂O₂ → 2H₂O + O₂

Single Replacement (Displacement) Reaction

• One element replaces another in a compound (A + BC → AC + B).
• Example: Zn + 2HCl → ZnCl₂ + H₂

Double Replacement (Displacement) Reaction

• The ions of two compounds swap places (AB + CD → AD + CB).
• Example: NaCl + AgNO₃ → NaNO₃ + AgCl

Combustion Reaction

• A substance reacts with oxygen, releasing energy as light and heat.
• Often involves hydrocarbons (CxHy + O₂ → CO₂ + H₂O).
• Example: CH₄ + 2O₂ → CO₂ + 2H₂O

Ionic Compounds: Naming and Properties

• Ionic compounds form when metals combine with nonmetals, involving a transfer of electrons and formation of ions.
• Positive ions are cations, and negative ions are anions.

Naming Ionic Compounds

• The cation (metal) comes first, followed by the anion (non-metal).
• Metals with multiple charges indicate the charge in parentheses with Roman numerals.
• NaCl is Sodium chloride.
• FeCl₂ is Iron(II) chloride.
• FeCl₃ is Iron(III) chloride.

Properties of Ionic Compounds

• They have high melting and boiling points due to strong electrostatic forces.
• They conduct electricity when molten or dissolved because ions are free to move.
• They are brittle due to their rigid structure.

Factors Affecting Rates of Reaction

Concentration

• Higher reactant concentrations increase the reaction rate due to more frequent collisions.

Temperature

• Increasing temperature speeds up reactions by increasing the frequency and energy of collisions.

Surface Area

• Smaller reactant pieces increase the reaction rate by increasing the surface area exposed.

Catalysts

• Catalysts speed up reactions without being consumed by lowering activation energy.

Pressure (for gases)

• Increased pressure increases the frequency of collisions, speeding up reactions.

Phase Transitions (Endothermic and Exothermic)

• Phase transitions are physical changes between states of matter (solid, liquid, gas).

Endothermic

• Endothermic transitions absorb energy (require heat).
  • Melting (solid to liquid)
  • Boiling (liquid to gas)
  • Sublimation (solid to gas)

Exothermic

• Exothermic transitions release energy (give off heat).
  • Freezing (liquid to solid)
  • Condensation (gas to liquid)
  • Deposition (gas to solid)

Molarity

• Molarity (M) measures solute concentration in a solution.
• Molarity is defined as moles of solute per liter of solution.
• M = (moles of solute) / (liters of solution)
• Example: 2 moles of NaCl in 1 liter of water has a molarity of 2 M.

Acids and Bases

Arrhenius Definition

• Acid: Increases H⁺ concentration in aqueous solution.
  • Example: HCl → H⁺ + Cl⁻
• Base: Increases OH⁻ concentration in aqueous solution.
  • Example: NaOH → Na⁺ + OH⁻

Bronsted-Lowry Definition

• Acid: A proton (H⁺) donor.
• Base: A proton (H⁺) acceptor.

Conjugates

• Acid loses a proton and forms its conjugate base.
  • Example: HCl (acid) → Cl⁻ (conjugate base) + H⁺
• Base gains a proton and forms its conjugate acid.
  • Example: NH₃ (base) + H⁺ → NH₄⁺ (conjugate acid)