Chemistry Exam Review

Chemical Equations and Reactions

Balancing Equations and Identifying Reaction Types

• Coefficients and Reaction Types:
  • Page 14 questions focus on coefficients and reaction types.
  • Examples:
    • Question 1: "blank 3 3 blank single"
    • Question 2: "balanced double"
    • Question 3: "blank 6 4 synthesis"
    • Question 4: "2 blank 2 blank single"
    • Question 5: "blank 2 2 blank double"
    • Question 6: "2 3 3 blank double"
    • Question 7: "2 3 6 blank double"
    • Question 8: "2 6 3 2 single"
    • Question 9: "blank 3 blank 3 double"
    • Question 10: "blank blank blank 2 double"
    • Question 11: "2 blank blank 2 double"
    • Question 12: "2 blank blank 2 double"
    • Question 13: "balanced single"
    • Question 14: "blank or blank or double"
    • Question 15: "2 blank blank blank double"

Synthesis Reactions (Questions 1-5)

• Synthesis reactions involve combining two or more substances into one compound.
• Key Point: Correctly combine ions based on their charges from the periodic table.

Example 1: Sodium and Oxygen

• Reactants: Sodium (Na) and Oxygen (O)
• Note: Oxygen is diatomic, so it exists as O_2 by itself. This is only when it's by itself, not in a compound.
• Ions: Sodium is Na^{+1}, Oxygen is O^{-2}
• Product: Sodium Oxide (Na_2O)
• Balanced Equation: 4Na + O2 \rightarrow 2Na2O

Example 2: Magnesium and Fluorine

• Reactants: Magnesium (Mg) and Fluorine (F)
• Ions: Magnesium is Mg^{+2}, Fluorine is F^{-1}
• Product: Magnesium Fluoride (MgF_2)
• Balanced Equation: Mg + F2 \rightarrow MgF2

General Rule for Synthesis Reactions with Single Elements

• If two single elements react, they will combine to form a compound.

Other Synthesis Examples

• Sodium and Chlorine: 2Na + Cl_2 \rightarrow 2NaCl
• Aluminum and Sulfur: 2Al + 3S \rightarrow Al2S3
• Calcium and Phosphorus: 3Ca + 2P \rightarrow Ca3P2

Decomposition Reactions (Questions 6-8)

• Decomposition reactions involve breaking down a compound into simpler substances.
• Key Point: Oxygen, Hydrogen, Nitrogen, Fluorine, Chlorine, Bromine, Iodine are always diatomic when by themselves.

Example 1: Mercury (II) Oxide

• Reactant: Mercury (II) Oxide (HgO)
• Products: Mercury (Hg) and Oxygen (O_2)
• Balanced Equation: 2HgO \rightarrow 2Hg + O_2

Example 2: Aluminum Oxide

• Reactant: Aluminum Oxide (Al2O3)
• Products: Aluminum (Al) and Oxygen (O_2)
• Balanced Equation: 2Al2O3 \rightarrow 4Al + 3O_2

Double Replacement Reactions (Questions 9-12)

• Double replacement reactions involve the exchange of ions between two compounds.
• Positive ions swap with positive ions and negative ions swap with negative ions.

Key Concepts

• Cations (positive ions) swap with cations.
• Anions (negative ions) swap with anions.

Example 1: Nickel (II) Sulfate and Aluminum Fluoride

• Reactants: Nickel (II) Sulfate (NiSO4) and Aluminum Fluoride (AlF3)
• Products: Nickel (II) Fluoride (NiF2) and Aluminum Sulfate (Al2(SO4)3)
• Balanced Equation: 3NiSO4 + 2AlF3 \rightarrow 3NiF2 + Al2(SO4)3

Example 2: Sodium Chloride and Fluorine

• Reactants: Sodium Chloride (NaCl) and Fluorine (F_2)
• Products: Sodium Fluoride (NaF) and Chlorine (Cl_2)
• Balanced Equation: 2NaCl + F2 \rightarrow 2NaF + Cl2

Example 3: Potassium Chloride and Sodium Hydroxide

• Reactants: Potassium Chloride (KCl) and Sodium Hydroxide (NaOH)
• Products: Potassium Hydroxide (KOH) and Sodium Chloride (NaCl)
• Balanced Equation: KCl + NaOH \rightarrow KOH + NaCl
• This reaction is already balanced.

Example 4: Barium Chloride and Sodium Phosphate

• Reactants: Barium Chloride (BaCl2) and Sodium Phosphate (Na3PO_4)
• Products: Barium Phosphate (Ba3(PO4)_2) and Sodium Chloride (NaCl)
• Balanced Equation: 3BaCl2 + 2Na3PO4 \rightarrow Ba3(PO4)2 + 6NaCl

Example 5: Silver Nitrate and Calcium Chloride

• Reactants: Silver Nitrate (AgNO3) and Calcium Chloride (CaCl2)
• Products: Silver Chloride (AgCl) and Calcium Nitrate (Ca(NO3)2)
• Balanced Equation: 2AgNO3 + CaCl2 \rightarrow 2AgCl + Ca(NO3)2

Combustion Reactions (Questions 13-15)

• Combustion reactions involve burning a substance in oxygen to produce carbon dioxide and water.
• General Form: Hydrocarbon + Oxygen → Carbon Dioxide + Water
• CxHy + O2 \rightarrow CO2 + H_2O
• Key Point: Always add oxygen (O2) as a reactant and carbon dioxide (CO2) and water (H_2O) as products.
• Example: CH4 + 2O2 \rightarrow CO2 + 2H2O

Balancing Combustion Reactions

• 1. Count the number of carbon, hydrogen, and oxygen atoms on both sides of the equation.
• 2. Balance the carbon atoms first.
• 3. Balance the hydrogen atoms.
• 4. Balance the oxygen atoms last.
• 5. If balancing oxygen results in an odd number, double the hydrocarbon and rebalance.

Example 1: Combustion of Methane (CH_4)

• CH4 + O2 \rightarrow CO2 + H2O
• Balanced Equation: CH4 + 2O2 \rightarrow CO2 + 2H2O

Example 2: Combustion of Acetylene (C2H2)

• C2H2 + O2 \rightarrow CO2 + H_2O
• Balanced Equation: 2C2H2 + 5O2 \rightarrow 4CO2 + 2H_2O

Example 3: Combustion of Heptane (C7H{16})

• C7H{16} + O2 \rightarrow CO2 + H_2O
• Balanced Equation: C7H{16} + 11O2 \rightarrow 7CO2 + 8H_2O

Factors Affecting Reaction Rate

Factors that affect the speed of reaction, not equilibrium or yield.

• Temperature: Higher temperature increases reaction rate.
  • Increases kinetic energy and collisions between particles.
• Concentration: Higher concentration increases reaction rate.
  • More particles in solution lead to more collisions.
• Surface Area: Smaller particle size (greater surface area) increases reaction rate.
• Catalyst: Speeds up reaction by lowering activation energy.
  • Not a reactant or product.
• Pressure: Higher pressure increases reaction rate for gases only.
  • Reduces volume and increases particle collisions.
  • Liquids and solids are negligibly affected.

Acids and Bases

Acids

• Release hydrogen ions (H^+) in water.
• Accept electron pairs.
• Examples:
  • Sulfuric Acid (H2SO4)
  • Hydrochloric Acid (HCl)
  • Phosphoric Acid (H3PO4)
• Common Acids:
  • Stomach acid: Hydrochloric acid (HCl)
  • Soda: Phosphoric acid (H3PO4)
  • Vinegar: Acetic acid

Bases

• Release hydroxide ions (OH^-) in solution.
• Donate electron pairs.
• Examples:
  • Ammonium Hydroxide (NH_4OH)
  • Potassium Hydroxide (KOH)
• Common Bases:
  • Soap: Potassium hydroxide (KOH)
  • Milk of Magnesia: Magnesium Hydroxide (Mg(OH)_2)
  • Drain Cleaner: Sodium Hydroxide (NaOH)
  • Baking Soda: Sodium Bicarbonate (NaHCO_3)

Properties of Acids and Bases

Neutralization Reactions

• Acids and bases react to form a salt and water.
• Salt: An ionic compound formed from the reaction of an acid and a base where the Hydrogen of the acid is replaced by a metal.
• Example:
  • HCl + NaOH \rightarrow NaCl + H_2O
  • H2SO4 + 2KOH \rightarrow K2SO4 + 2H_2O

pH Scale

• Measures the hydrogen ion concentration (H^+).
• 7 is neutral.
• Less than 7 is acidic.
• Greater than 7 is basic.
• The further from 7, the stronger the acid or base.
• Strong Acid: pH less than 3.
• Strong Base: pH greater than 11.

Indicators

• Acids turn red.
• Bases turn blue.