Chemistry Study Guide: Moles, Mass, Formulas, and Composition
Chapter 1: Foundations — Units, Measurements, & Matter
Units and Conversions
Quantity, SI Unit, Symbol
Length: meter (m)
Mass: kilogram (kg)
Time: second (s)
Temperature: kelvin (K)
Amount of substance: mole (mol)
Scientific notation is used for very large/small numbers (e.g., Avogadro’s Number = 6.022 \times 10^{23}).
Density (d) = mass ÷ volume
Example: A 63.4 g piece of metal has a density of 12.86 g/cm³. Volume = 63.4 ÷ 12.86 = 4.93 cm^3
Chapter 2: Atoms, Molecules, and Ions
Atomic Theory
All matter is made of atoms (Dalton's Theory).
Subatomic particles:
Proton: +1 charge, ~1 u
Neutron: 0 charge, ~1 u
Electron: –1 charge, ~1/1800 u
Atomic and Mass Numbers
Atomic number (Z) = number of protons
Mass number (A) = protons + neutrons
Isotopes and Atomic Mass
Average atomic mass = weighted average of all isotopes
Average \space mass = \sum(fractional \space abundance \times isotopic \space mass)
Example: Carbon: 98.89% ¹²C and 1.11% ¹³C = (0.9889)(12.00) + (0.0111)(13.00) = 12.01 u
Chapter 3: Stoichiometry — Formulas, Moles, and Mass
Avogadro’s Number and the Mole
1 mol = 6.022 × 10^{23} particles
Mass (in grams) of 1 mole of a substance = molar mass
Examples:
1 mol H₂O = 2(1.008) + 16.00 = 18.02 g
13.2 g Na → moles = 13.2 g ÷ 22.99 g/mol = 0.574 mol
Formula, Molecular, and Molar Mass
Formula Weight: Sum of atomic masses in ionic or covalent units. Example: NaCl = 22.99 + 35.45 = 58.44 u
Molecular Weight: Same as formula weight, but for molecules. Example: H₂SO₄ = 98.09 u
Molar Mass: Mass (g) of 1 mol of substance. Example: 1 mol H₂SO₄ = 98.09 g
Percent Composition
Percent \space by \space mass = \frac{n \times atomic \space mass}{molar \space mass} \times 100\%%
Example: Glucose, C₆H₁₂O₆
Total mass = 180.16 g/mol
%C = \frac{(6 × 12.01)}{180.16} × 100 = 40.00\%%
%H = \frac{(12 × 1.008)}{180.16} × 100 = 6.72\%%
%O = \frac{(6 × 16.00)}{180.16} × 100 = 53.28\%%
Empirical vs. Molecular Formulas
Empirical Formula
Convert % to grams (assume 100 g).
Convert grams to moles.
Divide all by smallest mole value.
Multiply to get integers if necessary.
Example: 40.92% C, 4.58% H, 54.50% O → Empirical formula = C₃H₄O₃
Molecular Formula
Calculate molar mass of empirical formula.
Divide molecular mass by empirical mass.
Multiply empirical subscripts by that factor.
Limiting Reactants and Stoichiometry
Limiting reagent: Used up first in a reaction.
Use mole ratios from the balanced chemical equation.
Example: 5.52 g Na + 5.10 g Al₂O₃
Balanced: 6Na + Al₂O₃ → 2Al + 3Na₂O
Determine limiting reactant by comparing mol to mol ratio.
Percent Yield
Percent \space Yield = \frac{actual \space yield}{theoretical \space yield} × 100\%%
Example: Theoretical yield = 42.8 g; actual yield = 37.6 g → % yield = \frac{37.6}{42.8} × 100 = 87.9\%%
Memory Aids
Diatomic Elements: BrINClHOF — Br₂, I₂, N₂, Cl₂, H₂, O₂, F₂
Empirical Steps: % → g → mol → divide → multiply
Polyatomic Ions: Learn common ones (e.g., NO₃⁻ , SO₄²⁻)
Okay, I can explain each section:
Chapter 1: Foundations — Units, Measurements, & Matter
Units and Conversions: This section covers the basic units of measurement used in science, particularly the SI units. It's important to understand these units and how to convert between them to solve problems.
Quantity, SI Unit, Symbol: Lists the fundamental quantities (length, mass, time, etc.) along with their corresponding SI units and symbols.
Scientific notation: Explains how to express very large or very small numbers in a compact and convenient form. For instance, Avogadro’s Number is expressed as 6.022 × 10^{23}.
Density: Defines density as mass divided by volume and provides an example calculation.
Chapter 2: Atoms, Molecules, and Ions
Atomic Theory: Introduces Dalton's atomic theory, which states that all matter is made of atoms.
Subatomic particles: Lists the three main subatomic particles (protons, neutrons, and electrons) along with their charges and approximate masses.
Atomic and Mass Numbers: Defines atomic number as the number of protons in an atom and mass number as the sum of protons and neutrons.
Isotopes and Atomic Mass: Explains that isotopes are atoms of the same element with different numbers of neutrons. It also defines average atomic mass as the weighted average of all isotopes, with an example calculation for carbon.
Chapter 3: Stoichiometry — Formulas, Moles, and Mass
Avogadro’s Number and the Mole: Defines the mole as a unit of measurement for the amount of substance, with Avogadro’s number as the number of particles in one mole. It also explains how to calculate the mass of one mole of a substance (molar mass), with examples for water and sodium.
Formula, Molecular, and Molar Mass: Defines formula weight, molecular weight, and molar mass, with examples for NaCl and H₂SO₄.
Percent Composition: Explains how to calculate the percent by mass of each element in a compound, with an example for glucose.
Empirical vs. Molecular Formulas: Describes the difference between empirical and molecular formulas and outlines the steps to determine each, with an example.
Empirical Formula:
Convert % to grams (assume 100 g).
Convert grams to moles.
Divide all by smallest mole value.
Multiply to get integers if necessary.
Example: 40.92% C, 4.58% H, 54.50% O → Empirical formula = C₃H₄O₃Molecular Formula
Calculate molar mass of empirical formula.
Divide molecular mass by empirical mass.
Multiply empirical subscripts by that factor.
Limiting Reactants and Stoichiometry: Defines the limiting reagent as the reactant that is used up first in a reaction. It also explains how to use mole ratios from the balanced chemical equation to determine the limiting reactant, with an example.
Use mole ratios from the balanced chemical equation.
Example: 5.52 g Na + 5.10 g Al₂O₃
Balanced: 6Na + Al₂O₃ → 2Al + 3Na₂O
Determine limiting reactant by comparing mol to mol ratio