Classification of Chemical Reactions

CHM 100: Classification of Chemical Reactions and Balancing Chemical Equations

Learning Objectives

• Understand the law of conservation of mass and how it applies to chemical equations.

• Write balanced chemical equations.

• Identify, describe and classify different chemical reaction types.

• Apply solubility rules to predict if a precipitation reaction will occur and write the appropriate balanced reaction.

• Predict the products of an acid-base neutralization reaction.

• Write net ionic equations for double replacement reactions to distinguish spectator ions from ions involved in the chemical reaction.

Chemical Equations

• Definition: A chemical equation is an expression using symbols and formulas to represent a chemical reaction. It acts as a "chemical recipe."

Features of a Chemical Equation

• Example of Chemical Reaction:

  • $2HgO(s) <br>ightleftharpoons 2Hg(l) + O_2(g)$

• Key Features:

  • Physical States: Shown in parentheses
    a. (s) = solid
    b. (l) = liquid
    c. (g) = gas
    d. (aq) = aqueous (solution)

  • The reaction arrow indicates the direction of the reaction and shows energy requirements.

  • Reactants: Located on the left of the arrow.

  • Products: Located on the right of the arrow.

  • Coefficients: Numbers placed in front of formulas indicating the number of units of a substance involved in the reaction.

Law of Conservation of Mass

• In a balanced chemical equation:

  • The law of conservation of mass states that matter is conserved.

  • No atoms are lost or gained.

  • The number of atoms for each element is equal on both sides of the reaction arrow.

Balancing Chemical Equations

• A balanced chemical equation indicates the precise amounts of reactants needed to generate a given amount of product.

• Requirements for a Balanced Equation:

  • The numbers and types of atoms must be equal on both sides of the reaction arrow.

  • Changes can only be made with coefficients (numbers), not by changing subscripts of formulas.

• Example: $2NaHCO_3(s) <br>ightleftharpoons Na_2CO_3(s) + H_2O(l) + CO_2(g)$

• Balancing Steps:

  1. Write an unbalanced equation with correct formulas for all reactants and products.

  2. Count the number of atoms of each element on both sides.

  3. Identify unbalanced elements (those not equal on both sides).

  4. Balance one element at a time by adjusting coefficients.

  5. Ensure the law of conservation of mass is maintained and coefficients are expressed in the lowest whole number.

Examples of Balancing Chemical Equations

• Balancing Example:

  • From $H_2 + O_2 <br>ightarrow H_2O$ (not balanced)

  • To balance:
    $2H_2 + O_2 → 2H_2O$ (balanced, obeys law of conservation)

  • Incorrect Change: Changing subscripts, e.g., $H_2 + O_2 o H_2O_2$ is incorrect.

Practice Problems: Writing Balanced Equations

1. Iron and Steam Equation:

   • Reaction: $Fe(s) + H_2O(g) → Fe_3O_4(s) + H_2(g)$

2. Haber Process Equation:

   • Reaction: $N_2(g) + 3H_2(g) → 2NH_3(g)$

3. Burning Ethane:

   • Reaction: $C_2H_6(g) + 7O_2(g) → 4CO_2(g) + 6H_2O(g)$

Types of Reactions

• Classification of chemical reactions includes:

  • Combination

  • Decomposition

  • Single Replacement

  • Double Replacement

  • Combustion

Combination Reactions

• Two or more elements or compounds combine to form one product.

• Example: $2Mg(s) + O_2(g) → 2MgO(s)$

Decomposition Reactions

• A single substance splits into two or more simpler substances.

• Example: $2HgO(s) → 2Hg(l) + O_2(g)$

Single Replacement Reactions

• One element replaces another in a compound.

• Example: $Zn(s) + 2HCl(aq) → ZnCl_2(aq) + H_2(g)$

Double Replacement Reactions

• The positive ions in reactants switch places.

• Example: $AgNO_3(aq) + NaCl(aq) → AgCl(s) + NaNO_3(aq)$

Combustion Reactions

• A substance reacts with oxygen, producing energy, carbon dioxide, and water.

• Example: $CH_4(g) + 2O_2(g) → CO_2(g) + 2H_2O(g) + ext{energy}$

Precipitation Reactions and Solubility Guidelines

• Definition: Precipitation reactions occur when an insoluble solid (precipitate) forms from solutions.

• Guidelines on Solubility:

  • Generally, many compounds are soluble but precipitate based on solubility rules.

  • Low solubility leads to precipitation; high solubility implies no precipitation.

General Rules on Solubility in Water

• Soluble compounds include:

  1. All nitrates, nitrites, and perchlorates.

  2. Alkali metal and ammonium salts.

  3. Most halogen salts.

  4. Most sulfate salts, with exceptions.

• Insoluble compounds include:

  1. Sulfides (many are insoluble).

  2. Most carbonates, phosphates, and chromates (many are insoluble).

Acid-Base Neutralization Reactions

• Neutralization involves reacting an acid with a base to produce salt and water.

• Removes H+ and OH– from solution, yielding neutral H2O.

• Common neutralization: Acid + Base → Water + Salt.

• Example reactions include:

  1. $HBr(aq) + Ba(OH)<em>{2}(aq) → BaBr</em>{2}(aq) + 2 H_2O(l)$

  2. Acid + Carbonate: produces water, salt, and carbon dioxide.

Writing Net Ionic Equations

• Net ionic equations exclude spectator ions and focus on the reactive species.

• Example:

  • Overall equation: $KOH(aq) + HNO_3(aq) → H_2O(l) + KNO_3(aq)$

  • Net Ionic: $K^+(aq) + OH^-(aq) → H_2O(l)$

Study Checks

• Questions assessing knowledge on balanced equations, reaction types, and net ionic equations.

• Examples include recognizing precipitate formation from given reactions and predicting products of acid-base neutralization.