Chemistry 1770 Exam 3 Study Guide: Stoichiometry, Solution Chemistry, and Energy

Chapter 7 – Stoichiometry

Greenhouse Gases and Atmospheric Chemistry

Greenhouse gases are molecules in Earth's atmosphere that absorb infrared radiation, contributing to the greenhouse effect and global warming.

• Greenhouse Gas: A gas that absorbs and emits infrared radiation, e.g., CO2, CH4, H2O.

• Concentration Influence: Higher concentrations increase absorption; described by Beer's Law: $A = \epsilon b c$ where A is absorbance, \epsilon is molar absorptivity, b is path length, and c is concentration.

• Wavelength Regime: Greenhouse gases interact with infrared wavelengths, causing molecular vibrations.

• Symmetry Breaking: For a molecule to absorb IR, its vibrational mode must break symmetry and change dipole moment. O2 and N2 do not have symmetry breaking, so they do not absorb IR.

• Resonance: Energy transfer occurs when the frequency of incoming radiation matches the natural frequency of molecular vibration (e.g., pushing a swing at its natural frequency).

Balancing Chemical Equations

Balancing equations ensures the conservation of mass and atoms in a chemical reaction.

• Skeletal Equation: An unbalanced equation showing reactants and products.

• Balanced Equation: Adjust coefficients to ensure equal numbers of each atom on both sides.

• PNOM Diagrams: Particle Number Operation Model diagrams visually represent balanced equations.

Limiting Reactant and Stoichiometry Calculations

Stoichiometry involves quantitative relationships between reactants and products.

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

• Combustion Reactions: Reactions where a substance reacts with O2 to produce CO2 and H2O (for hydrocarbons).

• Stoichiometry Calculations:

  • Mole-to-mole: Use balanced equation to convert between moles of reactants and products.

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

• Limiting Reactant Problems: Identify which reactant limits the reaction, perform calculations accordingly.

• Percentage Yield: Measures efficiency of a reaction: $\text{Percentage Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$ Adjust starting material based on expected yield.

Chapter 8 – Solution Chemistry

Water and Dissolution

Water is a universal solvent due to its polarity and ability to interact with ions and molecules.

• Interaction with Ions: Water molecules surround ions, stabilizing them via solvation.

• Ionic vs. Molecular Substances: Ionic substances dissociate into ions; molecular substances may not.

• Electrolytes:

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

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

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

• Water Footprint: Water's properties (high heat capacity, polarity) make it central to many applications.

Solution Terminology and Concentration Units

Understanding solution terminology and concentration units is essential for quantitative chemistry.

• Solute: Substance dissolved in a solvent.

• Solvent: Substance in which solute dissolves (usually present in greater amount).

• Dissolve: Process of solute dispersing in solvent.

• Solvation: Interaction of solvent molecules with solute.

• Concentration Unit Structure: Ratio of amount of solute to volume of solution.

• Molarity (M): $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Millimolar (mM): $1\,\text{mM} = 10^{-3}\,\text{M}$

• Moles from Molarity: $\text{moles} = M \times \text{volume (L)}$

• Dilution Equation: $M_1 V_1 = M_2 V_2$ (use only when no reaction occurs during dilution)

Reactions in Aqueous Solution

Many important reactions occur in water, including precipitation, acid-base, and redox reactions.

• Neutralization Reactions: Acid reacts with base to form water and salt.

• Solubility Rules: Used to predict whether a precipitate will form.

  Salt Type	Solubility	Exceptions
  Most nitrate (NO3–) salts	Soluble	None
  Group 1 metals or NH4+ salts	Soluble	None
  Chloride, bromide, iodide salts	Soluble	Ag+, Pb+, Hg22+
  Sulfide (S2–) & Carbonate (CO32–)	Insoluble	Except Group 1 metals

• Precipitation Reaction: Formation of an insoluble product (precipitate) from soluble reactants.

• Spectator Ions: Ions that do not participate in the reaction; remain unchanged.

• Writing Equations:

  • Molecular Equation: Shows all reactants and products.

  • Ionic Equation: Shows ions separately.

  • Net Ionic Equation: Shows only species that change during the reaction.

• Total Number of Ions: Calculate based on formula and dissociation.

Acids, Bases, and Redox Reactions

Acids and bases are defined by their behavior in water; redox reactions involve electron transfer.

• Arrhenius Definition:

  • Acid: Produces H+ in water.

  • Base: Produces OH– in water.

• Stoichiometry in Solutions: Used for precipitation and titration calculations.

• Oxidation Numbers: Assigned to atoms to track electron transfer.

• Redox Definitions:

  • Oxidation: Loss of electrons.

  • Oxidized: Species that loses electrons.

  • Oxidizing Agent: Causes oxidation; is reduced.

  • Reduction: Gain of electrons.

  • Reduced: Species that gains electrons.

  • Reducing Agent: Causes reduction; is oxidized.

• Identifying Redox Reactions: Use oxidation numbers to determine if oxidation/reduction occurred.

Chapter 9 – Energy

Geologic Hydrogen and Energy Sources

Hydrogen can be found underground and may serve as a future energy source.

• Drilling/Mining for Hydrogen: Extraction of hydrogen from geological formations.

• Formation of Hydrogen: Hydrogen forms underground via chemical reactions involving water and minerals.

Energy Concepts and Definitions

Energy is the capacity to do work; it exists in various forms and is central to chemical reactions.

• Energy: Capacity to do work or transfer heat.

• Heat (q): Energy transferred due to temperature difference.

• Work (w): Energy transferred when an object is moved by a force.

• Kinetic Energy: Energy of motion.

• Potential Energy: Stored energy due to position.

• System: Part of the universe under study.

• Surroundings: Everything outside the system.

Electrostatic Potential Energy and Coulomb’s Law

Electrostatic potential energy arises from interactions between charged particles.

• Coulomb’s Law: Describes the energy between two charges: $E_{el} = \frac{k Q_1 Q_2}{r}$ where k is a constant, Q1 and Q2 are charges, r is distance.

• Graphical Representation: Energy decreases as distance increases.

Thermodynamics and Enthalpy

Thermodynamics studies energy changes; enthalpy is the heat content of a system.

• Exothermic: Releases energy; ΔE or ΔH is negative.

• Endothermic: Absorbs energy; ΔE or ΔH is positive.

• First Law of Thermodynamics: Energy is conserved. $\Delta E = q + w$ Sign Conventions:

  • q is positive if heat is absorbed by the system.

  • w is positive if work is done on the system.

• Pressure-Volume Work (PV Work): Work done by expanding gases: $w = -P \Delta V$ Negative sign indicates work done by the system.

Heat Capacity and Specific Heat

Heat capacity and specific heat quantify how much energy is needed to change temperature.

• Heat Capacity (C): Amount of heat required to raise temperature by 1°C.

• Specific Heat (c): Heat required to raise 1 gram by 1°C. $q = m c \Delta T$ Water has high specific heat, important for climate and biological systems.

Thermochemical Equations and Hess’s Law

Thermochemical equations show energy changes in reactions; Hess’s Law allows calculation of enthalpy changes.

• Thermochemical Equation: Chemical equation with energy change.

• State Function: Property dependent only on current state, not path (e.g., energy, enthalpy).

• Hess’s Law: Enthalpy change is sum of enthalpy changes for individual steps. $\Delta H = \sum n \Delta H_{f}(\text{prod}) - \sum n \Delta H_{f}(\text{react})$

• Formation Reaction: Formation of 1 mole of compound from elements in their standard states.

• Heat of Formation: Enthalpy change for formation reaction.

• Bond Energies: Alternative calculation: $\Delta H = \sum \text{Bond Energies (react)} - \sum \text{Bond Energies (prod)}$

Heats of Formation Table

Standard heats of formation are used in Hess’s Law calculations.

Compound	ΔHf (kJ·mol–1)
C3H8(g)	–104
C4H10(g)	–126
C8H18(g)	–250.3
CO(g)	–110.5
CO2(g)	–393.5
H2O(l)	–285.8
O2(g)	0
N2(g)	0
NO(g)	90.3
NO2(g)	33.2
N2O(g)	82.1

Example: Calculate ΔH for a reaction using heats of formation: $\Delta H = \sum n \Delta H_{f}(\text{prod}) - \sum n \Delta H_{f}(\text{react})$

Additional info: PNOM diagrams are visual tools for representing stoichiometry and limiting reactant concepts. Concept maps may be used to organize aqueous reaction types.