Classification & Balancing of Chemical Reactions: GOB Chemistry Study Guide

Classification & Balancing of Chemical Reactions

Concept: Chemical Reaction & Chemical Change

Chemical reactions involve the breaking and forming of chemical bonds, resulting in new substances. Observable evidence of a chemical reaction includes color change, formation of a precipitate, and evolution of gas.

• Chemical Reaction: A process in which substances (reactants) are transformed into new substances (products).

• Observable Evidence: Color change, formation of a solid (precipitate), gas evolution, temperature change.

• Example: $2\text{NaHCO}_3 (s) \rightarrow \text{Na}_2\text{CO}_3 (s) + \text{H}_2\text{O} (l) + \text{CO}_2 (g)$

Law of Conservation of Mass

The Law of Conservation of Mass states that matter is neither created nor destroyed in a chemical reaction. The total mass of reactants equals the total mass of products.

• Equation: $2\text{H}_2 (g) + \text{O}_2 (g) \rightarrow 2\text{H}_2\text{O} (l)$

• Application: Used to calculate the amount of products or reactants in a chemical reaction.

Balancing Chemical Equations

Balancing chemical equations ensures that the number of atoms of each element is the same on both sides of the equation, satisfying the Law of Conservation of Mass.

• Steps to Balance:

  1. List the number of atoms for each element in reactants and products.

  2. Adjust coefficients to balance atoms for each element.

  3. Check your work to ensure all elements are balanced.

• Example: $\_\text{C}_2\text{H}_6 (g) + \_\text{O}_2 (g) \rightarrow \_\text{CO}_2 (g) + \_\text{H}_2\text{O} (g)$

Solubility Rules

Solubility rules help determine whether an ionic compound will dissolve in water (soluble) or form a precipitate (insoluble).

• Soluble Compounds: Most compounds containing alkali metal ions (Li+, Na+, K+), ammonium (NH4+), nitrate (NO3-), and acetate (C2H3O2-) are soluble.

• Insoluble Compounds: Most compounds containing carbonate (CO32-), phosphate (PO43-), and sulfide (S2-) are insoluble, except when paired with alkali metals or ammonium.

• Example: Na2SO4 is soluble in water.

• Table: Solubility Rules

Ion/Compound	Solubility	Exceptions
Alkali metals, NH4+	Soluble	None
NO3-, C2H3O2-	Soluble	None
Cl-, Br-, I-	Soluble	Ag+, Pb2+, Hg22+
CO32-, PO43-	Insoluble	Alkali metals, NH4+

Molecular, Complete Ionic, and Net Ionic Equations

Chemical reactions in aqueous solution can be represented as molecular equations, complete ionic equations, and net ionic 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 actually participate in the reaction.

• Example: $\text{AgNO}_3 (aq) + \text{NaCl} (aq) \rightarrow \text{AgCl} (s) + \text{NaNO}_3 (aq)$

Types of Chemical Reactions

Chemical reactions are classified based on the changes in reactants and products.

• Combination (Synthesis): Two or more substances combine to form one product.

• Decomposition: One substance breaks down into two or more products.

• Single Displacement: One element replaces another in a compound.

• Double Displacement: Exchange of ions between two compounds.

• Combustion: A substance reacts with oxygen, releasing energy as heat and light.

• Table: Types of Chemical Reactions

Type	General Form	Example
Combination	$A + B \rightarrow AB$	$2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$
Decomposition	$AB \rightarrow A + B$	$2\text{H}_2\text{O} \rightarrow 2\text{H}_2 + \text{O}_2$
Single Displacement	$A + BC \rightarrow AC + B$	$\text{Zn} + 2\text{HCl} \rightarrow \text{ZnCl}_2 + \text{H}_2$
Double Displacement	$AB + CD \rightarrow AD + CB$	$\text{AgNO}_3 + \text{NaCl} \rightarrow \text{AgCl} + \text{NaNO}_3$
Combustion	$C_xH_y + O_2 \rightarrow CO_2 + H_2O$	$\text{CH}_4 + 2\text{O}_2 \rightarrow CO_2 + 2\text{H}_2O$

Oxidation Numbers and Redox Reactions

Oxidation numbers are used to track electron transfer in redox reactions. Redox (reduction-oxidation) reactions involve the transfer of electrons between reactants.

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

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

• Oxidizing Agent: Causes oxidation, is itself reduced.

• Reducing Agent: Causes reduction, is itself oxidized.

• Example: $\text{Zn} (s) + \text{Cu}^{2+} (aq) \rightarrow \text{Zn}^{2+} (aq) + \text{Cu} (s)$

Spontaneous Redox Reactions & Activity Series

The activity series ranks metals by their ability to be oxidized. A metal higher in the series will displace a metal lower in the series from a compound.

• Spontaneous Reaction: Occurs if the free metal is above the metal ion in the activity series.

• Example: $\text{Ca} (s) + \text{AgNO}_3 (aq) \rightarrow \text{Ca(NO}_3)_2 (aq) + \text{Ag} (s)$

Balancing Redox Reactions (Acidic and Basic Solutions)

Balancing redox reactions requires splitting the reaction into half-reactions and balancing electrons, atoms, and charges. The process differs slightly for acidic and basic solutions.

• Steps for Acidic Solution:

  1. Write half-reactions for oxidation and reduction.

  2. Balance all elements except H and O.

  3. Balance O by adding H2O.

  4. Balance H by adding H+.

  5. Balance charge by adding electrons.

  6. Combine half-reactions and cancel electrons.

• Steps for Basic Solution:

  1. Follow steps for acidic solution.

  2. Add OH- to both sides to neutralize H+ and form H2O.

• Example: $\text{MnO}_4^- (aq) + \text{Fe}^{2+} (aq) \rightarrow \text{Mn}^{2+} (aq) + \text{Fe}^{3+} (aq)$

Galvanic and Electrolytic Cells

Electrochemical cells convert chemical energy to electrical energy (galvanic cells) or use electrical energy to drive chemical reactions (electrolytic cells).

• Galvanic Cell: Spontaneous redox reaction generates electricity.

• Electrolytic Cell: Nonspontaneous reaction driven by external electrical energy.

• Anode: Site of oxidation.

• Cathode: Site of reduction.

• Salt Bridge: Maintains electrical neutrality by allowing ion flow.

• Example: In a Zn/Cu cell, electrons flow from Zn (anode) to Cu (cathode).

Summary Table: Electrochemical Cell Components

Component	Galvanic Cell	Electrolytic Cell
Anode	Oxidation (negative)	Oxidation (positive)
Cathode	Reduction (positive)	Reduction (negative)
Electron Flow	Anode → Cathode	Anode → Cathode
Energy Source	Spontaneous reaction	External power supply

Additional info: These notes expand on the original content by providing definitions, examples, and tables for clarity and completeness, suitable for GOB Chemistry students preparing for exams.