Honors Chemistry Unit 4B - Predicting Products, Net Ionic Equations, & Redox Reactions

1. Predicting Products of Chemical Reactions

When predicting reaction products, first identify the reaction type. Each type has specific rules and patterns to follow.

a. Synthesis (Combination) Reactions

  • Definition: Two or more reactants combine to form a single product.

  • General Formula: A + B → AB

Common Synthesis Reactions to Memorize

  1. Metal + Nonmetal → Ionic Compound

    • Example: 2Na+Cl2→2NaCl2Na + Cl_2 → 2NaCl2Na+Cl2​→2NaCl

    • Sodium reacts with chlorine gas to form sodium chloride (table salt).

  2. Nonmetal Oxide + Water → Acid

    • Example: CO2+H2O→H2CO3CO_2 + H_2O → H_2CO_3CO2​+H2​O→H2​CO3​

    • Carbon dioxide dissolves in water to form carbonic acid.

  3. Metal Oxide + Water → Base (Metal Hydroxide)

    • Example: CaO+H2O→Ca(OH)2CaO + H_2O → Ca(OH)_2CaO+H2​O→Ca(OH)2​

    • Calcium oxide reacts with water to form calcium hydroxide.

🔥 Practice Problem:
Predict the product of the reaction: Mg+O2Mg + O_2Mg+O2​

b. Decomposition Reactions

  • Definition: A single reactant breaks down into two or more simpler substances.

  • General Formula: AB → A + B

Common Decomposition Reactions to Memorize

  1. Metal Carbonates → Metal Oxide + CO₂

    • Example: CaCO3→CaO+CO2CaCO_3 → CaO + CO_2CaCO3​→CaO+CO2​

    • Calcium carbonate decomposes into calcium oxide and carbon dioxide.

  2. Metal Hydroxides → Metal Oxide + H₂O

    • Example: Mg(OH)2→MgO+H2OMg(OH)_2 → MgO + H_2OMg(OH)2​→MgO+H2​O

  3. Metal Chlorates → Metal Chloride + O₂

    • Example: 2KClO3→2KCl+3O22KClO_3 → 2KCl + 3O_22KClO3​→2KCl+3O2​

🔥 Practice Problem:
Predict the products of: H2O2→?H_2O_2 → ?H2​O2​→?

c. Single Replacement Reactions

  • Definition: A more reactive element replaces a less reactive one in a compound.

  • General Formula: A + BC → B + AC

Activity Series (Memorize!):
💡 A metal higher on the series will replace a metal lower in a compound.

  • Most Reactive Metals: Li > K > Ca > Na > Mg > Al > Zn > Fe > Ni > Sn > Pb > H > Cu > Ag > Au

  • Most Reactive Halogens: F₂ > Cl₂ > Br₂ > I₂

🔥 Practice Problem:
Will this reaction occur?
Cu+HCl→?Cu + HCl → ?Cu+HCl→? (Use the activity series!)

d. Double Replacement Reactions

  • Definition: Two ionic compounds swap ions to form two new compounds.

  • General Formula: AB + CD → AD + CB

💡 Reaction Occurs If:
✅ A precipitate (solid), gas, or water is formed.
✅ Use Solubility Rules to determine precipitates.

Solubility Rules (Key to Memorize)

  • Soluble (Dissolves in Water):

    • Alkali metals (Na⁺, K⁺, Li⁺)

    • Nitrates (NO₃⁻), Acetates (C₂H₃O₂⁻), Ammonium (NH₄⁺)

    • Halides (except Ag⁺, Pb²⁺, Hg₂²⁺)

    • Sulfates (except Ba²⁺, Pb²⁺, Sr²⁺)

  • Insoluble (Forms a Precipitate):

    • Carbonates (CO₃²⁻), Phosphates (PO₄³⁻)

    • Hydroxides (OH⁻), except with Na⁺, K⁺, NH₄⁺

🔥 Practice Problem:
Predict the products & state whether a precipitate forms:
Na2SO4+BaCl2→?Na_2SO_4 + BaCl_2 → ?Na2​SO4​+BaCl2​→?

e. Combustion Reactions

  • Definition: A hydrocarbon reacts with oxygen to produce CO₂ + H₂O.

  • General Formula: CxHy + O₂ → CO₂ + H₂O

🔥 Practice Problem:
Balance the combustion of propane:
C3H8+O2→CO2+H2OC_3H_8 + O_2 → CO_2 + H_2OC3​H8​+O2​→CO2​+H2​O

2. Net Ionic Equations

  • Shows only the reacting ions (removes spectator ions).

  • Steps to Determine Net Ionic Equation:

    1. Write the balanced equation.

    2. Break aqueous compounds into ions.

    3. Cancel spectator ions.

    4. Write the final net ionic equation.

Example:
Na2SO4+BaCl2→BaSO4(s)+2NaClNa_2SO_4 + BaCl_2 → BaSO_4 (s) + 2NaClNa2​SO4​+BaCl2​→BaSO4​(s)+2NaCl

  • Ionic: 2Na++SO42−+Ba2++2Cl−→BaSO4(s)+2Na++2Cl−2Na^+ + SO_4^{2-} + Ba^{2+} + 2Cl^- → BaSO_4 (s) + 2Na^+ + 2Cl^-2Na++SO42−​+Ba2++2Cl−→BaSO4​(s)+2Na++2Cl−

  • Net Ionic: Ba2++SO42−→BaSO4(s)Ba^{2+} + SO_4^{2-} → BaSO_4 (s)Ba2++SO42−​→BaSO4​(s)

3. Redox (Oxidation-Reduction) Reactions

  • Oxidation: Loss of electrons (LEO = Lose Electrons Oxidation)

  • Reduction: Gain of electrons (GER = Gain Electrons Reduction)

Oxidation Number Rules

  1. Pure elements: 0 (e.g., O₂, Na, Cl₂)

  2. Ions: Charge of ion (e.g., Na⁺ = +1)

  3. Oxygen: -2 (except peroxides: -1)

  4. Hydrogen: +1 (except metal hydrides: -1)

  5. Sum in neutral compound = 0, in polyatomic ion = charge of ion

Balancing Redox Equations (Half-Reaction Method)

  1. Split into oxidation & reduction half-reactions.

  2. Balance atoms except H & O.

  3. Balance O using H₂O.

  4. Balance H using H⁺ (acidic) or OH⁻ (basic).

  5. Balance charge with electrons.

  6. Multiply half-reactions to equalize electrons.

  7. Add reactions together.

Example (Acidic Solution):
MnO4−+Fe2+→Mn2++Fe3+MnO_4^- + Fe^{2+} → Mn^{2+} + Fe^{3+}MnO4−​+Fe2+→Mn2++Fe3+
Balanced:
MnO4−+8H++5Fe2+→Mn2++5Fe3++4H2OMnO_4^- + 8H^+ + 5Fe^{2+} → Mn^{2+} + 5Fe^{3+} + 4H_2OMnO4−​+8H++5Fe2+→Mn2++5Fe3++4H2​O

🔥 Practice Problem:
Balance in acidic solution:
Cr2O72−+Fe2+→Cr3++Fe3+Cr_2O_7^{2-} + Fe^{2+} → Cr^{3+} + Fe^{3+}Cr2​O72−​+Fe2+→Cr3++Fe3+