Ch 7: Chemical Reactions: Types, Equations, and Aqueous Chemistry

Chemical Reactions and Equations

Introduction to Chemical Reactions

Chemical reactions involve the transformation of substances into new substances with different properties. These changes are fundamental to chemistry and are observed in everyday phenomena such as combustion, rusting, and the functioning of batteries.

• Evidence of a chemical reaction includes color change, temperature change, gas production, precipitate formation, emission of light, and emission or absorption of heat.

• Chemical equations represent these reactions, showing reactants (starting materials) and products (new substances formed).

Evidence and Types of Chemical Reactions

Observing Chemical Change

• Color Change: Indicates new substances are formed (e.g., copper turning green).

• Formation of a Solid (Precipitate): A solid appears when two solutions are mixed (e.g., milky white precipitate).

• Formation of a Gas: Bubbles or fizzing (e.g., baking soda and vinegar reaction).

• Emission of Light: Light produced (e.g., glowsticks).

• Emission or Absorption of Heat: Temperature changes (e.g., cold packs, burning fuels).

Not all changes are chemical. For example, melting ice or evaporating butane are physical changes because the molecules remain the same.

Writing and Balancing Chemical Equations

Structure of Chemical Equations

• Reactants are written on the left, products on the right.

• States of matter are indicated: (s) solid, (l) liquid, (g) gas, (aq) aqueous.

• Law of Conservation of Mass: Atoms are neither created nor destroyed; equations must be balanced.

Steps to Balance Equations

1. Write the unbalanced (skeletal) equation with correct formulas.

2. Balance elements that appear in only one compound on each side first (usually metals).

3. Balance free elements last.

4. Clear any fractional coefficients by multiplying the entire equation.

5. Verify that the number of each atom is the same on both sides.

Example: Combustion of Methane

• Unbalanced: $\mathrm{CH_4(g) + O_2(g) \rightarrow CO_2(g) + H_2O(g)}$

• Balanced: $\mathrm{CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(g)}$

Example: Photosynthesis

• Balanced: $\mathrm{6CO_2 + 6H_2O \rightarrow C_6H_{12}O_6 + 6O_2}$

Aqueous Solutions and Solubility

Properties of Aqueous Solutions

• Aqueous solution: A homogeneous mixture of a substance dissolved in water.

• Strong electrolytes: Substances that completely dissociate into ions in water (e.g., NaCl, AgNO3).

• Conductivity: Solutions with ions conduct electricity; pure water does not.

• Polyatomic ions (e.g., $\mathrm{NO_3^-}$) remain intact when dissolved.

Solubility Rules

Solubility rules help predict whether an ionic compound will dissolve in water.

Compound Contains	Solubility	Exceptions
Li+, Na+, K+, NH4+, NO3-, C2H3O2-	Soluble	None
Cl-, Br-, I-	Soluble	Ag+, Hg22+, Pb2+
SO42-	Soluble	Ca2+, Sr2+, Ba2+, Pb2+
OH-, S2-	Insoluble	Li+, Na+, K+, NH4+; S2- is soluble with Ca2+, Sr2+, Ba2+; OH- is slightly soluble with Ca2+, Sr2+, Ba2+
CO32-, PO43-	Insoluble	Li+, Na+, K+, NH4+

Examples:

• AgBr: Insoluble (Br- with Ag+ is an exception).

• CaCl2: Soluble (Cl- is generally soluble, Ca2+ is not an exception).

• Pb(NO3)2: Soluble (NO3- is always soluble).

• PbSO4: Insoluble (SO42- with Pb2+ is an exception).

Precipitation Reactions

Formation of a Precipitate

Precipitation reactions occur when two aqueous solutions combine to form an insoluble solid (precipitate).

• Example: $\mathrm{2KI(aq) + Pb(NO_3)_2(aq) \rightarrow PbI_2(s) + 2KNO_3(aq)}$ (yellow precipitate forms).

• If all products are soluble, write "NO REACTION".

Steps to Predict Precipitation Reactions

1. Write formulas of reactants.

2. Write formulas of possible products by exchanging ions.

3. Use solubility rules to determine if any product is insoluble.

4. If a product is insoluble, write it as a solid; if not, write "NO REACTION".

5. Balance the equation.

Example:

• $\mathrm{Na_2CO_3(aq) + CuCl_2(aq) \rightarrow CuCO_3(s) + 2NaCl(aq)}$ (CuCO3 is insoluble).

• $\mathrm{LiNO_3(aq) + Na_2SO_4(aq) \rightarrow NO\ REACTION}$ (all products soluble).

Chemical Equations in Solution: Molecular, Complete Ionic, and Net Ionic

Types of Equations

• Molecular Equation: Shows all reactants and products as compounds.

• Complete Ionic Equation: Shows all strong electrolytes as ions.

• Net Ionic Equation: Shows only the species that participate in the reaction (removes spectator ions).

Example: Silver Nitrate and Sodium Chloride

• Molecular: $\mathrm{AgNO_3(aq) + NaCl(aq) \rightarrow AgCl(s) + NaNO_3(aq)}$

• Complete Ionic: $\mathrm{Ag^+(aq) + NO_3^-(aq) + Na^+(aq) + Cl^-(aq) \rightarrow AgCl(s) + Na^+(aq) + NO_3^-(aq)}$

• Net Ionic: $\mathrm{Ag^+(aq) + Cl^-(aq) \rightarrow AgCl(s)}$

Acid–Base and Gas-Evolution Reactions

Acid–Base (Neutralization) Reactions

• Acids produce H+ ions in solution; bases produce OH- ions.

• Acid–base reactions form water and a salt.

• General equation: $\mathrm{Acid + Base \rightarrow Water + Salt}$

• Net ionic equation: $\mathrm{H^+(aq) + OH^-(aq) \rightarrow H_2O(l)}$

Example:

• $\mathrm{HCl(aq) + NaOH(aq) \rightarrow H_2O(l) + NaCl(aq)}$

• Net ionic: $\mathrm{H^+(aq) + OH^-(aq) \rightarrow H_2O(l)}$

Gas-Evolution Reactions

• Produce a gas as a product, often via acid–base reactions.

• Some reactions form an unstable intermediate that decomposes into a gas.

Common Gas-Evolution Patterns:

Reactant Type	Intermediate	Gas Formed
Sulfides	None	H2S
Carbonates/Bicarbonates	H2CO3	CO2
Sulfites/Bisulfites	H2SO3	SO2
Ammonium Compounds	NH4OH	NH3

Example:

• $\mathrm{2HNO_3(aq) + Na_2CO_3(aq) \rightarrow H_2O(l) + CO_2(g) + 2NaNO_3(aq)}$

Oxidation–Reduction (Redox) and Combustion Reactions

Redox Reactions

• Involve the transfer of electrons between substances.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Mnemonic: OIL RIG (Oxidation Is Loss, Reduction Is Gain).

• Common in rusting, combustion, and battery operation.

Examples:

• $\mathrm{2H_2(g) + O_2(g) \rightarrow 2H_2O(g)}$

• $\mathrm{4Fe(s) + 3O_2(g) \rightarrow 2Fe_2O_3(s)}$

• $\mathrm{CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(g)}$

Combustion Reactions

• Type of redox reaction where a substance reacts with O2, releasing energy as heat and light.

• Products are usually CO2 and H2O (for hydrocarbons).

Example:

• $\mathrm{2C_8H_{18}(l) + 25O_2(g) \rightarrow 16CO_2(g) + 18H_2O(g)}$

Classifying Chemical Reactions

Main Types of Reactions

• Synthesis (Combination): Two or more substances combine to form one product. General: $\mathrm{A + B \rightarrow AB}$

• Decomposition: One compound breaks into two or more simpler substances. General: $\mathrm{AB \rightarrow A + B}$

• Single-Displacement: One element replaces another in a compound. General: $\mathrm{A + BC \rightarrow AC + B}$

• Double-Displacement: Ions in two compounds exchange places. General: $\mathrm{AB + CD \rightarrow AD + CB}$

Examples:

• Synthesis: $\mathrm{2Na(s) + Cl_2(g) \rightarrow 2NaCl(s)}$

• Decomposition: $\mathrm{2H_2O_2(aq) \rightarrow 2H_2O(l) + O_2(g)}$

• Single-Displacement: $\mathrm{Zn(s) + CuCl_2(aq) \rightarrow ZnCl_2(aq) + Cu(s)}$

• Double-Displacement: $\mathrm{AgNO_3(aq) + NaCl(aq) \rightarrow AgCl(s) + NaNO_3(aq)}$

Special Types

• Precipitation reactions and acid–base reactions are double-displacement reactions.

• Combustion reactions are a type of redox reaction.

Summary Table: Reaction Types and Examples

Type	General Form	Example
Synthesis	A + B → AB	2Na(s) + Cl2(g) → 2NaCl(s)
Decomposition	AB → A + B	2H2O2(aq) → 2H2O(l) + O2(g)
Single-Displacement	A + BC → AC + B	Zn(s) + CuCl2(aq) → ZnCl2(aq) + Cu(s)
Double-Displacement	AB + CD → AD + CB	AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)
Combustion	Fuel + O2 → CO2 + H2O	CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)

Applications and Relevance

• Acid–base reactions are essential for biological systems (e.g., maintaining blood pH).

• Combustion reactions provide energy for transportation and electricity.

• Precipitation reactions are used in water treatment and analytical chemistry.

• Redox reactions are fundamental to batteries and metabolic processes.

Key Definitions

• Reactant: A starting substance in a chemical reaction.

• Product: A substance formed as a result of a chemical reaction.

• Precipitate: An insoluble solid formed from a reaction in solution.

• Electrolyte: A substance that conducts electricity when dissolved in water.

• Redox Reaction: A reaction involving electron transfer.

• Acid: Substance that produces H+ in water.

• Base: Substance that produces OH- in water.

Additional info: This guide expands on the provided notes with definitions, examples, and tables for clarity and completeness, as would be expected in a mini-textbook summary for introductory chemistry students.