CHM1032 Study Notes: Chemical Reactions, Stoichiometry, and Atomic Structure (Chapters 7, 8, and 9)

Chemical Reactions

Evidence of Chemical Reactions

Chemical reactions involve the transformation of substances into new products. Evidence that a chemical reaction has occurred includes:

• Color change

• Formation of a precipitate (solid from solution)

• Gas evolution (bubbling, odor change)

• Temperature change (exothermic or endothermic)

• Emission of light

Combustion reactions are a type of chemical reaction where a substance reacts with oxygen, releasing energy, often used in laundry detergents to break down stains.

Chemical Equations

A chemical equation represents a chemical reaction using symbols and formulas. It shows reactants (starting materials) and products (new substances formed).

• Balanced chemical equations have equal numbers of each type of atom on both sides, obeying the Law of Conservation of Mass.

Example: The combustion of methane:

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

Aqueous Solutions and Solubility

Many reactions occur in water (aqueous solutions). Solubility is the ability of a substance to dissolve in water.

• Electrolytes: Substances that dissolve in water to produce ions, conducting electricity (e.g., NaCl).

• Nonelectrolytes: Substances that dissolve but do not produce ions (e.g., sugar).

Precipitation Reactions

When two aqueous solutions are mixed, an insoluble product (precipitate) may form.

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

Writing Chemical Equations for Reactions in Solution

• 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 change during the reaction.

Acid-Base and Gas Evolution Reactions

• Acid-base reactions: Involve transfer of protons (H+), typically forming water and a salt.

• Gas evolution reactions: Produce a gas as a product (e.g., CO2, H2).

Oxidation-Reduction (Redox) Reactions

Redox reactions involve the transfer of electrons between substances.

• Oxidation: Loss of electrons

• Reduction: Gain of electrons

Oxidation States

• Assigned to atoms to track electron transfer.

• Rules for assigning oxidation states help identify redox changes.

Rules for Assigning Oxidation States

1. Elemental form: 0

2. Monatomic ion: charge of the ion

3. Oxygen: usually -2

4. Hydrogen: +1 with nonmetals, -1 with metals

5. Fluorine: always -1

6. Sum of oxidation states equals the charge of the molecule/ion

Identifying Redox Reactions

• Look for changes in oxidation states of elements between reactants and products.

Classifying Chemical Reactions by Atom Movement

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

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

• Displacement (Single Replacement): An element replaces another in a compound.

Stoichiometry and Reaction Quantities

The Greenhouse Effect

The greenhouse effect is the warming of Earth's atmosphere due to the trapping of heat by greenhouse gases like CO2. Chemical reactions, such as combustion, contribute to greenhouse gas emissions.

Reaction Stoichiometry

Stoichiometry is the calculation of reactants and products in chemical reactions using balanced equations.

• Mole-to-mole conversions: Use coefficients from the balanced equation to relate moles of reactants and products.

• Mass-to-mass conversions: Convert mass to moles, use stoichiometry, then convert back to mass.

Example: For $\mathrm{2H_2 + O_2 \rightarrow 2H_2O}$, 2 moles of H2 produce 2 moles of 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: $\mathrm{\%\ yield = \frac{actual\ yield}{theoretical\ yield} \times 100\%}$

To determine these, start with initial masses, convert to moles, identify the limiting reactant, calculate theoretical yield, and compare to actual yield.

Enthalpy: Heat in Chemical Reactions

Enthalpy (ΔH) measures the heat evolved or absorbed in a reaction at constant pressure.

• Sign of ΔHrxn: Negative for exothermic (releases heat), positive for endothermic (absorbs heat).

• Stoichiometry of ΔHrxn: The enthalpy change corresponds to the amounts in the balanced equation.

Example: $\mathrm{CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O}$, $\Delta H = -890\ \mathrm{kJ}$ (exothermic)

Atomic Structure and Periodicity

Models Explaining Element Reactivity

• Bohr model: Electrons orbit the nucleus in fixed paths (energy levels).

• Quantum-mechanical model: Electrons exist in orbitals, regions of probability.

Light and Electromagnetic Radiation

• Electromagnetic radiation: Energy transmitted as waves (includes visible light, UV, IR, etc.).

• Photon: A particle of light energy.

• Electromagnetic spectrum: Range of all types of electromagnetic radiation.

The Bohr Model and Atomic Spectra

• Excitation: Electron absorbs energy and moves to a higher energy level.

• Emission: Electron returns to lower energy level, emitting a photon.

• Emission spectra: Unique set of wavelengths emitted by an element.

The Quantum-Mechanical Model: Orbitals

• Principal quantum number (n): Indicates main energy level.

• Orbitals: Regions where electrons are likely to be found (s, p, d, f types).

• Ground state: Lowest energy arrangement of electrons.

• Excited state: Any arrangement with higher energy than ground state.

Electron Configurations

• Electron configuration: Distribution of electrons among orbitals.

• Orbital diagrams: Visual representation using arrows for electrons.

• Electron spin: Each orbital holds two electrons with opposite spins (Pauli Exclusion Principle).

• Ground-state configuration: Fill lowest energy orbitals first (Aufbau principle).

• Energy levels and sublevels: s < p < d < f in energy.

• Noble gas core notation: Use noble gas symbol to abbreviate inner electrons.

Electron Configurations and the Periodic Table

• Valence electrons: Electrons in the outermost shell, determine chemical properties.

• Irregular configurations: Some elements (especially transition metals) have exceptions.

• Transition metal configurations: d sublevel fills after s sublevel.

• Properties and electron configuration: Reactivity, magnetism, and color depend on electron arrangement.

• Noble gas configuration: Stable, full outer shell.

Periodic Trends

• Atomic radius (main group): Increases down a group, decreases across a period.

• Transition metals: Atomic radius changes less across the period.

• Ionization energy: Energy required to remove an electron; increases across a period, decreases down a group.

• Metallic character: Increases down a group, decreases across a period.

Electron Configurations of Ions

• Main-group anions: Gain electrons to achieve noble gas configuration.

• Main-group cations: Lose electrons to achieve noble gas configuration.

Type	Electron Change	Example
Main-group cation	Lose electrons	Na: $1s^2 2s^2 2p^6 3s^1$ → Na+: $1s^2 2s^2 2p^6$
Main-group anion	Gain electrons	Cl: $1s^2 2s^2 2p^6 3s^2 3p^5$ → Cl-: $1s^2 2s^2 2p^6 3s^2 3p^6$

Additional info: Some transition metals lose s electrons before d electrons when forming cations.