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 warming of the planet. Their concentration directly influences the amount of energy absorbed.

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

• Absorption and Beer's Law: The absorption of light by greenhouse gases is described by Beer's Law: $A = \epsilon b c$ where A is absorbance, \epsilon is the molar absorptivity, b is path length, and c is concentration.

• Wavelength/Frequency 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 the dipole moment. O2(g) and N2(g) do not have symmetry breaking in their vibrational modes, so they do not absorb IR.

• Resonance in Energy Transfer: Resonance occurs when the frequency of an external force matches the natural frequency of a system (e.g., pushing a swing or slinky at the right 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

The limiting reactant determines the maximum amount of product formed in a reaction.

• Limiting Reactant: The reactant that is completely consumed first, limiting product formation.

• Stoichiometry Calculations: Use mole ratios from balanced equations to convert between moles and masses of reactants/products.

• Mole-to-Mole: Use coefficients to relate moles of one substance to another.

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

Combustion Reactions

Combustion reactions involve a substance reacting with oxygen to produce energy, often forming CO2 and H2O.

• Complete Combustion: All fuel reacts with O2 to form CO2 and H2O.

• O2 as Reactant: Oxygen is always a reactant in combustion.

Percentage Yield Problems

Percentage yield compares actual product obtained to theoretical yield.

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

• Adjusting Starting Material: Use expected yield to calculate required reactant amounts.

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 in solution (solvation).

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

• Electrolytes: Strong electrolytes dissociate completely; weak electrolytes partially; nonelectrolytes do not dissociate.

Key Properties and Definitions

• 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 particles.

Concentration Units and Calculations

Concentration measures the amount of solute in a given volume of solution.

• Basic Structure: Concentration = amount of solute / amount of solution.

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

• Millimolar (mM): 1 mM = 0.001 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

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

• Molecular, Ionic, Net Ionic Equations: Show reactants/products as molecules, ions, or only those involved in reaction.

• Solubility Rules: Predict whether a compound is soluble or forms a precipitate.

• 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 for Ionic Substances in Water

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

• Net Ionic Equations: Show only ions and molecules directly involved in the reaction.

Acids, Bases, and Redox Reactions

• Arrhenius Definition: Acids produce H+ in water; bases produce OH–.

• Identifying Acids/Bases: Use chemical equations to determine.

• Stoichiometry in Solutions: Apply to precipitation and titration reactions.

• Assigning Oxidation Numbers: Systematic method to track electron transfer.

• Redox Definitions:

  • Oxidation: Loss of electrons.

  • Oxidized: Substance that loses electrons.

  • Oxidizing Agent: Causes oxidation; gains electrons.

  • Reduction: Gain of electrons.

  • Reduced: Substance that gains electrons.

  • Reducing Agent: Causes reduction; loses electrons.

• Identifying Redox Reactions: Use oxidation numbers to determine what is oxidized/reduced.

Abbreviated Solubility Rules

Compound Type	Solubility
Most nitrate (NO3–) salts	Soluble
Most salts with Group 1 metals or NH4+	Soluble
Most chloride, bromide, iodide salts	Soluble, exceptions: Ag+, Pb+, Hg22+
Sulfide (S2–) & Carbonate (CO32–)	Insoluble (except Group 1 metals)

Chapter 9 – Energy

Geologic Hydrogen as an Energy Source

Hydrogen found underground may be a future energy resource, obtained by drilling or mining. It forms through geologic processes.

• Formation: Hydrogen forms underground via reactions between water and minerals.

Definitions Related to Energy

• 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 or arrangement.

• 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: $E_{el} = \frac{k Q_1 Q_2}{r}$ where k is a constant, Q1 and Q2 are charges, r is distance.

• Graphical Representation: Energy decreases as distance increases.

Exothermic and Endothermic Processes

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

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

First Law of Thermodynamics

Energy is conserved; it cannot be created or destroyed.

• Equation: $\Delta E = q + w$

• Sign Conventions: q and w are positive if energy flows into the system, negative if out.

• Calculating ΔE: Add heat and work values.

Pressure-Volume Work (PV Work)

• PV Work: Work done by expanding or compressing gases.

• Sign Convention: Work done by the system is negative; work done on the system is positive.

• Applications: Internal combustion engines, breathing.

Heat Capacity and Specific Heat

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

• Specific Heat (c): Heat required to raise 1 g of substance by 1°C.

• Calculation: $q = m c \Delta T$ where m is mass, c is specific heat, ΔT is temperature change.

• Water’s High Heat Capacity: Important for climate and biological systems.

Thermochemical Equations and Energy Calculations

• Thermochemical Equation: Shows energy change with chemical reaction.

• Energy Implications: Use stoichiometry to relate energy to reactant/product amounts.

State Functions

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

Hess’s Law and Heats of Formation

• Hess’s Law: Total enthalpy change is sum of enthalpy changes for individual steps.

• Equation: $\Delta H = \sum n \Delta H_{f}^{\circ}(\text{prod}) - \sum n \Delta H_{f}^{\circ}(\text{react})$

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

• Heat of Formation: Enthalpy change for formation reaction.

• Using Heats of Formation: Apply in Hess’s Law calculations.

Heats of Formation Table

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

Bond Energy Equation

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