chapter 7

7.1: Chemical Equations

Introduction to Chemical Reactions

A chemical reaction is when one or more substances are converted into new substances, involving breaking and forming of chemical bonds among atoms. These transformations alter the structure and properties of the original substances, leading to the formation of products distinct from the reactants.

Key Terms

• Reactants: Substances that undergo the transformation.

• Products: New substances formed as a result of the reaction.

Writing Chemical Equations

Chemical equations are symbolic representations of chemical reactions using chemical formulas. The equations convey the identities of the reactants and products, as well as their quantities and states.

Common Symbols:

• (s): solid

• (l): liquid

• (g): gas

• (aq): aqueous (dissolved in water)

• →: yields or produces

Example:

Sulfur dioxide (SO2) reacts with oxygen (O2) to form sulfur trioxide (SO3).Equation: 2 SO2(g) + O2(g) → 2 SO3(g)

Steps to Convert Word Equations to Chemical Equations

1. Identify reactants and products from the word equation.

2. Write correct chemical formulas for all reactants and products.

3. Ensure the equation is balanced, meaning there are equal numbers of each type of atom on both sides of the equation.

Symbols in Chemical Equations

Important symbols used in chemical equations often include:

• +: separates multiple reactants or products.

• (s), (l), (g), (aq): indicate the physical state of the substance.

• Δ: indicates heating is required for the reaction.

Learning Objectives

• Identify reactants and products in a chemical reaction.

• Convert word equations to chemical equations with precision.

7.2: Writing Balanced Chemical Equations

Concept of Balancing Equations

The Law of Conservation of Mass states that mass cannot be created or destroyed. Therefore, during a chemical reaction, the number of atoms of each element must remain equal on both sides of the equation. This principle is crucial in ensuring accurate chemical calculations and predictions.

Coefficients and Subscripts

• Coefficients: Numbers in front of compounds that indicate how many molecules or moles of compounds are involved in the reaction.

• Subscripts: Part of the chemical formula that indicates the number of atoms of each element in a molecule. These cannot be changed to balance the equation.

Steps to Balance Equations

1. Identify the most complex molecule in the equation.

2. Balance elements one at a time, starting with those that appear in only one reactant and one product.

3. Balance polyatomic ions as a unit if they appear unchanged on both sides of the equation.

4. Use whole number ratios for coefficients to avoid fractional values when balancing the equation.

Balancing Example:

Combustion of HeptaneEquation: C7H16 + O2 → CO2 + H2OFollow the balancing steps to ensure all atoms are accounted for, providing a clearer understanding of stoichiometry in chemical reactions.

7.3: Classifying Chemical Reactions

Types of Chemical Reactions

1. Synthesis (Combination):

   • Reaction Format: A + B → AB

   • Example: 2H2 + O2 → 2H2O

2. Decomposition:

   • Reaction Format: AB → A + B

   • Example: 2HgO → 2Hg + O2

3. Single Replacement:

   • Reaction Format: AB + C → AC + B

   • Example: Zn + CuSO4 → ZnSO4 + Cu

4. Double Replacement:

   • Reaction Format: AB + CD → AD + CB

   • Example: BaCl2 + Na2SO4 → BaSO4 + 2NaCl

5. Combustion:

   • Reaction Format: Hydrocarbon + O2 → CO2 + H2O

   • Example: C3H8 + 5O2 → 3CO2 + 4H2O

Predicting Products

Utilize the properties of the reactants, such as their reactivity and physical states, to predict the likely products of the reactions and understand reaction mechanisms.

7.4: Evidence of Chemical Reactions

Indicators of a Chemical Change

1. Color change

2. Production of gas (bubbles)

3. Formation of a precipitate (solid) from liquids

4. Temperature change (heat release or absorption)

Examples of Chemical Changes

• Burning wood

• Rusting iron

• Cooking food

• Photosynthesis

7.5: Aqueous Solutions and Solubility

Definition of Electrolytes

• Electrolytes: Substances that release ions into solution when dissolved, allowing for the conduction of electricity.

• Strong Electrolytes: Completely dissociate in solution (most ionic compounds).

• Weak Electrolytes: Partially dissociate in solution (weak acids/bases).

Solubility Rules

Rules to predict the solubility of compounds in water include:

1. All salts of Group IA (alkali metals) and ammonium are soluble.

2. All nitrates and acetates are soluble.

3. Sulfates are soluble except for those of barium, lead, and calcium.

4. Most carbonates and phosphates are insoluble except those associated with alkali metals.

7.6: Precipitation Reactions

Definition

A precipitation reaction is a double-replacement reaction that results in the formation of an insoluble product, often referred to as a precipitate. This is a vital concept in analytical chemistry and environmental studies.

Example of Precipitation

Mixing silver nitrate and potassium chloride yields silver chloride precipitate, illustrating practical applications of solubility rules.

7.7: Writing Chemical Equations in Aqueous Solutions

Types of Equations

1. Molecular Equation: Presents all reactants and products in their molecular form.

2. Complete Ionic Equation: Displays all soluble ionic compounds dissociated into their constituent ions.

3. Net Ionic Equation: Only includes the species that participate in the reaction, omitting spectator ions. This distinction is important for understanding chemical processes on a molecular level.

7.8: Acid-Base Reactions

Neutralization Reactions

Definition

The reaction between an acid and a base produces salt and water, often demonstrating the principles of reaction stoichiometry.

Example

HCl + NaOH → NaCl + H2O

7.9: Gas Evolution Reactions

Definition and Types

Gas evolution reactions release gas as a product, showcasing transformations that occur during the interactions of chemicals. Common examples include acid-carbonate reactions producing CO2, emphasizing the principles of gas solubility and reaction equilibrium in aqueous solutions.

7.10: Oxidation-Reduction Reactions

Definitions

• Oxidation: Loss of electrons; increase in oxidation state.

• Reduction: Gain of electrons; decrease in oxidation state.

Balancing Redox Reactions

Use half-reaction methods to identify changes in oxidation states, which is essential in studying electrochemistry and modeling energy changes in reactions.

Example

The combustion of hydrocarbon fuels releases water and carbon dioxide, demonstrating practical implications for energy production and chemical kinetics.