Chapter 4: Chemical Reactions and Stoichiometry – Study Notes

Chapter 4: Chemical Reactions and Stoichiometry

Writing and Balancing Chemical Equations

Identifying and Balancing Chemical Equations

Chemical equations represent the transformation of reactants into products. Balancing these equations ensures the law of conservation of mass is obeyed, meaning the number of atoms of each element is the same on both sides of the equation.

• Reactants: Substances present before the reaction (left side of the arrow).

• Products: Substances formed by the reaction (right side of the arrow).

• Physical States: Indicated in parentheses: (s) solid, (l) liquid, (g) gas, (aq) aqueous.

• Stoichiometric Coefficients: Numbers placed before compounds to balance the number of atoms.

Example: Combustion of methane:

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

• 4 H atoms on both sides, 4 O atoms on both sides, 1 C atom on both sides.

• Coefficients must be the lowest possible integers.

Nuclear Reactions (Brief Introduction)

Nuclear reactions involve changes in the nucleus and may emit particles such as alpha ($\alpha$), beta ($\beta$), or gamma ($\gamma$) radiation.

• Alpha Decay ($\alpha$): Emission of a $^4_2\mathrm{He}$ nucleus (2 protons, 2 neutrons).

• Beta Decay ($\beta$): Emission of an electron ($^0_{-1}\mathrm{e}$) when a neutron converts to a proton.

• Gamma Emission ($\gamma$): Emission of high-energy electromagnetic radiation, often accompanying other decay types.

Example: Alpha decay of uranium-238:

$^{238}_{92}\mathrm{U} \rightarrow ^{234}_{90}\mathrm{Th} + ^4_2\mathrm{He}$

• Sum of atomic and mass numbers is conserved.

Solutions and Solubility

Solubility and Dissolution

Solubility describes how well a solute dissolves in a solvent. For ionic compounds, dissolution in water involves the separation of ions due to electrostatic interactions with water molecules.

• Solute: Substance being dissolved.

• Solvent: Substance doing the dissolving (water in most cases).

• Electrolytes: Substances that produce ions in solution and conduct electricity.

• Strong Electrolytes: Completely dissociate (e.g., NaCl).

• Weak Electrolytes: Partially dissociate (e.g., acetic acid).

• Nonelectrolytes: Do not produce ions (e.g., sugar).

Solubility Rules for Ionic Compounds

Solubility rules help predict whether an ionic compound will dissolve in water. The following table summarizes key rules:

Compound Type	Solubility
Group 1 cations, NH4+	Soluble
Nitrates (NO3-), acetates (CH3COO-), chlorates, perchlorates	Soluble
Chlorides, bromides, iodides	Soluble (except Ag+, Pb2+, Hg22+)
Sulfates (SO42-)	Soluble (except Ca2+, Sr2+, Ba2+, Pb2+)
Carbonates, hydroxides, oxides, phosphates, sulfides	Insoluble (except as above)

Additional info: These rules are hierarchical; earlier rules take precedence over later ones.

Precipitation Reactions

Formation of Precipitates

When two aqueous solutions are mixed, an insoluble product (precipitate) may form if the combination of ions produces an insoluble compound.

• 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 change during the reaction (removes spectator ions).

Example:

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

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

Acid-Base Reactions

Arrhenius Definition

Acids are substances that produce $\mathrm{H^+}$ ions in aqueous solution; bases produce $\mathrm{OH^-}$ ions.

• Strong Acids: Completely dissociate (e.g., HCl, HNO3, H2SO4).

• Weak Acids: Partially dissociate (e.g., HF, acetic acid).

• Strong Bases: Group 1 and 2 metal hydroxides (e.g., NaOH, KOH, Ca(OH)2).

• Diprotic/Polyprotic Acids: Can donate more than one proton (e.g., H2SO4 is diprotic).

Example: Neutralization reaction:

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

Oxidation-Reduction (Redox) Reactions

Oxidation States and Redox Concepts

Redox reactions involve the transfer of electrons, resulting in changes in oxidation states of elements.

• Oxidation: Loss of electrons (increase in oxidation state).

• Reduction: Gain of electrons (decrease in oxidation state).

• Oxidizing Agent: Causes oxidation, is itself reduced.

• Reducing Agent: Causes reduction, is itself oxidized.

Rules for Assigning Oxidation States:

• Elemental form: 0

• Monoatomic ion: Equal to its charge

• Sum in neutral molecule: 0; in polyatomic ion: equals ion charge

• Group 1 metals: +1; Group 2 metals: +2

• Hydrogen: +1 (usually)

• Fluorine: -1; Other group 17: usually -1

• Oxygen: usually -2; Group 16: usually -2

• Group 15: usually -3

Balancing Redox Equations (Half-Reaction Method)

1. Write separate half-reactions for oxidation and reduction.

2. Balance atoms other than O and H.

3. Balance O by adding $\mathrm{H_2O}$; balance H by adding $\mathrm{H^+}$ (or $\mathrm{OH^-}$ in basic media).

4. Balance charge by adding electrons ($\mathrm{e^-}$).

5. Multiply half-reactions to equalize electrons, then add together.

6. Check that atoms and charges are balanced.

Example: $\mathrm{Fe^{2+} \rightarrow Fe^{3+} + e^-}$ (oxidation half-reaction)

Disproportionation Reactions

In disproportionation, a single species is both oxidized and reduced. For example:

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

• Oxygen in $\mathrm{H_2O_2}$ is both reduced (to $\mathrm{H_2O}$) and oxidized (to $\mathrm{O_2}$).

Reaction Stoichiometry

Mole and Mass Relationships

Stoichiometry uses balanced chemical equations to relate the amounts (in moles or mass) of reactants and products.

• Coefficients in the equation indicate mole ratios.

• Use molar mass to convert between mass and moles.

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

• 2 moles H2 react with 1 mole O2 to produce 2 moles H2O.

Limiting Reactant, Theoretical Yield, and Percent Yield

• Limiting Reactant: The reactant that is completely consumed first, limiting the amount of product formed.

• Theoretical Yield: Maximum amount of product possible from given reactants.

• Percent Yield: $\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$

Example: If 3.00 moles H2 and 2.35 moles O2 are mixed, H2 is limiting and only 3.00 moles H2O can form.

Solution Concentration and Solution Stoichiometry

Molarity and Dilution

Molarity (M): Number of moles of solute per liter of solution.

$M = \frac{n}{V}$

• n = moles of solute

• V = volume of solution in liters

Dilution Equation: $M_1V_1 = M_2V_2$

• Used to calculate the concentration after dilution.

Example: To dilute 1.0 mol/L stock to 0.01 mol/L in a 100 mL flask:

$M_1V_1 = M_2V_2 \implies (1.0)(V_1) = (0.01)(0.100)$ $V_1 = 0.001$ L = 1.0 mL

Solution Stoichiometry

• Use molarity, volume, and balanced equations to calculate amounts of reactants or products in solution reactions.

Additional info: For more complex solution stoichiometry, combine the concepts of molarity, mole ratios, and limiting reactant calculations.