General Chemistry: Stoichiometry, Aqueous Reactions, and Thermochemistry Study Notes

Chapter 3: Mass, Moles, Formulas, and Stoichiometry

Percentage Composition

Percentage composition describes the mass percent of each element in a compound. This is essential for determining empirical formulas and analyzing chemical samples.

• Fraction = part / whole

• Percent = (part / whole) × 100

• For compounds, the "whole" is the molar mass (formula weight), and the "part" is the mass contribution from each element.

• Steps:

  1. Find the molar mass of the compound.

  2. Calculate the mass from each element (subscript × atomic mass).

  3. Compute % element = (element mass / total molar mass) × 100.

  4. Check units: amu/amu or g/mol over g/mol cancels, leaving percent.

• Example: Sucrose, C12H22O11 (C = 12, H = 1, O = 16)

  • Carbon: 12 × 12 = 144

  • Hydrogen: 22 × 1 = 22

  • Oxygen: 11 × 16 = 176

  • Total = 342 g/mol

  • %C = (144 / 342) × 100 ≈ 42.1%

  • %H = (22 / 342) × 100 ≈ 6.43%

  • %O = (176 / 342) × 100 ≈ 51.5%

The Mole and Avogadro’s Number

The mole is a counting unit in chemistry, analogous to a dozen, but representing 6.02 × 1023 particles (Avogadro’s number).

• 1 mole = 6.02 × 1023 particles (atoms, molecules, or ions).

• Dimensional analysis is used to convert between moles, molecules, and atoms.

• Example: Number of hydrogen atoms in 0.35 mol C6H12O6:

  1. 0.35 mol × (6.02 × 1023 molecules / 1 mol) = 2.11 × 1023 molecules

  2. 2.11 × 1023 molecules × (12 H atoms / 1 molecule) = 2.53 × 1024 H atoms

Molar Mass

Molar mass is the mass of one mole of a substance, numerically equal to the formula weight in g/mol.

• Example: Glucose, C6H12O6:

  • C: 6 × 12 = 72

  • H: 12 × 1 = 12

  • O: 6 × 16 = 96

  • Total = 180 g/mol

Grams ↔ Moles Conversions

Conversions between mass and moles are fundamental in stoichiometry.

• Core formula: $\text{moles} = \frac{\text{grams}}{\text{molar mass}}$

• Rearrangements:

  • grams = moles × molar mass

  • molar mass = grams / moles

• Example: Moles of glucose in 5.38 g:

  • mol = 5.38 g ÷ 180 g/mol = 0.0299 mol

Stoichiometry: Amounts of Reactants and Products

Stoichiometry uses balanced chemical equations to relate amounts of reactants and products.

• Balance the equation first.

• Coefficients represent mole ratios.

• Convert between grams and moles using molar mass.

• Units must cancel throughout calculations.

• Example: Oxidation of Glucose

The balanced equation for the oxidation of glucose is:

$\mathrm{C_6H_{12}O_6 + 6O_2 \rightarrow 6CO_2 + 6H_2O}$

• This equation shows that 1 mole of glucose reacts with 6 moles of oxygen to produce 6 moles of carbon dioxide and 6 moles of water.

• Mass relationships can be established using molar masses and stoichiometric coefficients.

Chapter 4: Reactions in Aqueous Solution

What “Aqueous” Means

An aqueous solution is one in which a substance is dissolved in water. Water is a common solvent due to its polarity and ability to dissolve many ionic and molecular compounds.

• Solvent: The component in greater amount (water in aqueous solutions).

• Solute: The component in lesser amount (the dissolved substance).

Electrolytes vs Nonelectrolytes

Electrolytes are substances that dissolve in water to produce ions and conduct electricity. Nonelectrolytes dissolve but do not produce ions.

• Strong electrolytes: Dissociate completely (e.g., NaCl).

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

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

Solubility Rules

Solubility rules help predict whether an ionic compound will dissolve in water.

• All nitrates (NO3−) are soluble.

• Acetates are soluble.

• Chlorides, bromides, and iodides are soluble except with Ag+, Hg, Pb2+.

• Sulfates are generally soluble except with Ba2+, Sr2+, Pb2+.

• Sulfides, carbonates, phosphates, and hydroxides are generally insoluble except with Group 1 and NH4+ (and some exceptions for hydroxides).

Precipitation Reactions

Precipitation reactions occur when two aqueous solutions combine to form an insoluble solid (precipitate).

• Example: KI + Pb(NO3)2 → PbI2 (yellow precipitate) + KNO3 (remains dissolved)

Acids and Bases

Acids and bases can be defined by their behavior in water (Arrhenius) or by their ability to donate or accept protons (Brønsted–Lowry).

• Arrhenius acid: Produces H+ in water.

• Arrhenius base: Produces OH− in water.

• Brønsted–Lowry acid: Proton donor.

• Brønsted–Lowry base: Proton acceptor.

• Strong acids: HCl, HBr, HI, HNO3, HClO3, HClO4, H2SO4 (first dissociation).

• Strong bases: Group 1A hydroxides, Ca(OH)2, Sr(OH)2, Ba(OH)2.

Neutralization Reactions

Neutralization occurs when an acid reacts with a base to produce a salt and water.

• Molecular equation: $\mathrm{HCl}(aq) + \mathrm{NaOH}(aq) \rightarrow \mathrm{NaCl}(aq) + \mathrm{H_2O}(l)$

• Total ionic equation: $\mathrm{H^+} + \mathrm{Cl^-} + \mathrm{Na^+} + \mathrm{OH^-} \rightarrow \mathrm{Na^+} + \mathrm{Cl^-} + \mathrm{H_2O}$

• Net ionic equation: $\mathrm{H^+} + \mathrm{OH^-} \rightarrow \mathrm{H_2O}(l)$

Chapter 5: Thermochemistry and Thermodynamics

System vs Surroundings

The system is the part of the universe being studied, while the surroundings are everything else. Energy transfer occurs between the system and surroundings.

Exothermic vs Endothermic Processes

• Exothermic: System releases heat to surroundings (ΔH < 0).

• Endothermic: System absorbs heat from surroundings (ΔH > 0).

State Functions vs Path Functions

• State function: Depends only on initial and final states (e.g., internal energy E, enthalpy H).

• Path function: Depends on the path taken (e.g., heat q, work w).

Work in Thermodynamics

• Pressure–volume work: $w = -P\Delta V$

• If gas expands (ΔV > 0): w is negative (system does work on surroundings).

• If gas is compressed (ΔV < 0): w is positive (surroundings do work on system).

Internal Energy and Enthalpy

• Internal energy change: $\Delta E = q + w$

• Enthalpy: $H = E + PV$

• At constant pressure: $\Delta H = q_p$

• ΔH > 0: endothermic; ΔH < 0: exothermic

Calorimetry

• Measures heat flow via temperature change.

• Coffee-cup calorimeter (constant pressure): $q = m c \Delta T$

• Bomb calorimeter (constant volume): $q = C_{cal} \Delta T$

Hess’s Law and Standard Enthalpy of Formation

• Hess’s Law: $\Delta H_{overall} = \sum \Delta H_{steps}$

• Standard enthalpy of formation (ΔHf°): Enthalpy change for forming 1 mole of a compound from its elements in standard states at 25°C and 1 atm.

Additional info:

• These notes cover key concepts from Chapters 3, 4, and 5 of a general chemistry course, including stoichiometry, aqueous reactions, and thermochemistry.

• All equations are provided in LaTeX format for clarity and academic rigor.