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

Chapter 7 – Stoichiometry

Greenhouse Gases and Atmospheric Chemistry

Greenhouse gases are atmospheric gases that absorb and emit infrared radiation, contributing to the greenhouse effect and warming the Earth's surface.

• Greenhouse Gas: A gas that absorbs infrared radiation, such as CO2, CH4, and H2O vapor.

• Concentration Influence: Higher concentrations increase absorption, affecting atmospheric temperature.

• Beer's Law: Describes how absorbance depends on concentration and path length: $A = \epsilon b c$ where $A$ is absorbance, $\epsilon$ is 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 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, slinky demonstration).

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: Visual representations of particle numbers and organization in reactions.

Limiting Reactant Concept

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

• Identification: Compare mole ratios of reactants to the balanced equation.

• PNOM Diagrams: Used to visualize which reactant runs out first.

Combustion Reactions

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

• Complete Combustion: Oxygen is always a reactant; products are CO2 and H2O.

Stoichiometry Calculations

Stoichiometry involves quantitative relationships between reactants and products.

• Mole-to-Mole: Use balanced equations to convert between moles of substances.

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

Limiting Reactant Problems

• Recognition: If reactant amounts are not in the exact stoichiometric ratio, a limiting reactant problem exists.

• Math and PNOM: Solve using calculations and diagrams.

Percentage Yield Problems

Percentage yield measures the efficiency of a reaction.

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

• Adjusting Reactant Amounts: Use expected yield to determine required starting material.

Chapter 8 – Solution Chemistry

Water and Ionic Interactions

Water dissolves ionic substances by surrounding ions, separating them from the crystal lattice.

• 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.

Properties and Applications of Water

Water's polarity, high heat capacity, and solvent abilities make it central to chemistry and environmental science.

• Water Footprint: Measures water usage in processes and products.

Definitions: Solute, Solvent, Dissolve, Solvation

• Solute: Substance dissolved in a solvent.

• Solvent: Substance in which solute dissolves (often water).

• Dissolve: Process of solute dispersing in solvent.

• Solvation: Interaction of solvent molecules with solute particles.

Concentration Units and Molarity

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

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

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

• Calculations: Use $\text{moles} = M \times \text{volume (L)}$

Dilution Calculations

Dilution reduces concentration by adding solvent.

• Equation: $M_1 V_1 = M_2 V_2$

• When Not to Use: If a reaction occurs during dilution, use stoichiometry instead.

Reactions in Aqueous Solution

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

• Concept Maps: Used to organize reaction types and processes.

Neutralization Reactions

Acid and base react to form water and a salt.

• Molecular Equation: Shows all reactants and products.

• Ionic Equation: Shows ions involved.

• Net Ionic Equation: Shows only ions that change.

Solubility Rules

Solubility rules help predict whether a compound will dissolve in water.

Compound 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 Reactions

Precipitation occurs when two solutions form an insoluble product.

• Definition: Reaction forming a solid from two aqueous solutions.

Spectator Ions

Spectator ions do not participate in the reaction and remain unchanged.

• Identification: Present on both sides of the equation.

Writing Ionic Equations

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

• Net Ionic Equation: Remove spectator ions to show actual chemical change.

Arrhenius Acids and Bases

Arrhenius acids produce H+ in water; bases produce OH–.

• Identification: Use chemical equation to determine acid or base.

Stoichiometry in Solution

• Precipitation and Titration: Use molarity and volume to calculate reactant/product amounts.

Assigning Oxidation Numbers

Oxidation numbers help track electron transfer in reactions.

• Rules: Assign based on element, compound, and ion.

Redox Definitions

• Oxidation: Loss of electrons.

• Oxidized: Substance losing electrons.

• Oxidizing Agent: Causes oxidation, is reduced.

• Reduction: Gain of electrons.

• Reduced: Substance gaining electrons.

• Reducing Agent: Causes reduction, is oxidized.

Identifying Redox Reactions

• Use Oxidation Numbers: If numbers change, redox has occurred.

• Identify: What was oxidized and what was reduced.

Chapter 9 – Energy

Geologic Hydrogen as Energy Source

Hydrogen found underground may be extracted for energy.

• Drilling/Mining: Methods to obtain geologic hydrogen.

• Formation: Hydrogen forms underground via chemical reactions.

Energy Definitions

• Energy: Capacity to do work or produce heat.

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

• Work (w): Energy transferred when an object is moved by 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: $E = \frac{k Q_1 Q_2}{r}$ where $E$ is energy, $k$ is a constant, $Q_1$ and $Q_2$ are charges, $r$ is distance.

• Graphical Description: Energy decreases as distance increases.

Exothermic and Endothermic Reactions

Exothermic reactions release energy; endothermic reactions absorb energy.

• Sign Convention: $\Delta E$ or $\Delta H$ is negative for exothermic, positive for endothermic.

First Law of Thermodynamics

Energy cannot be created or destroyed, only transferred.

• Equation: $\Delta E = q + w$

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

• Calculation: Use values of $q$ and $w$ to find $\Delta E$.

Pressure-Volume Work (PV Work)

Work done by a system expanding or contracting against external pressure.

• Equation: $w = -P \Delta V$

• Sign Convention: Work is negative when system does work on surroundings.

• Applications: Internal combustion engines, breathing.

Heat Capacity and Specific Heat

Heat capacity is the amount of heat required to raise temperature; specific heat is per unit mass.

• Heat Capacity (C): $C = \frac{q}{\Delta T}$

• Specific Heat (c): $q = m c \Delta T$

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

Thermochemical Equations

Thermochemical equations show energy changes in reactions.

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

State Functions

State functions depend only on the current state, not the path taken.

• Examples: Energy, enthalpy, pressure, volume.

Hess’ Law

Hess’ Law allows calculation of reaction enthalpy by summing enthalpies of component reactions.

• Calculation: Add or subtract equations and their enthalpy changes.

Formation Reactions and Heats of Formation

Formation reactions produce one mole of a compound from elements in their standard states.

• Heat of Formation: $\Delta H_f$ is enthalpy change for formation reaction.

• Use in Hess’ Law: Combine heats of formation to calculate overall reaction enthalpy.