Introductory Chemistry Study Notes: Measurements, Metric System, Matter, Atomic Models, Periodic Table, Chemical Language, Reactions, Mole Concept, and Equation Calculations

CHAPTER 1 – Measurements

Measurement Fundamentals

Measurements are essential in chemistry and are made using instruments. They are expressed in metric units, and every measurement carries some degree of uncertainty.

• Length: Centimeters (cm)

• Mass: Grams (g)

• Volume: Milliliters (mL)

• Uncertainty: All measurements have uncertainty, typically indicated by the last digit.

• Example: 5.0 cm (±0.1) means the true value is between 4.9 and 5.1 cm.

Significant Digits (Sig Figs)

Significant digits reflect the precision of a measurement. The uncertainty lies in the last reported digit.

• Rules for Determining Sig Figs:

  • Count digits from left to right, starting with the first nonzero digit.

  • Do not count placeholder zeros (e.g., 0.11g = 2 sigfigs).

  • Exact numbers (counted items, defined relationships) do not use sig figs.

• Expressing Placeholders: Use scientific notation to clarify significant zeros (e.g., 1.00 x 102 cm has 3 sigfigs).

• Rounding Rules:

  • If the first non-sigfig is less than 5, drop all non-sigfigs.

  • If the first non-sigfig is 5 or greater, increase the last sigfig by 1.

• Mathematical Operations:

  • Add/Subtract: Answer limited by the value with the most uncertainty (fewest decimal places).

  • Multiply/Divide: Answer limited by the measurement with the least sig figs.

Scientific Notation & Exponential Numbers

Scientific notation is used to express very large or small numbers and clarify significant digits.

• Example: 555,000 g = $5.55 \times 10^5$ g (3 sigfigs)

CHAPTER 2 – The Metric System

Basic Units and Metric Prefixes

The metric system uses base units and prefixes to indicate magnitude.

• Length: meter (m)

• Mass: gram (g)

• Volume: liter (L)

• Time: second (s)

• Prefixes:

  • Kilo (k): $1 \times 10^3$

  • Centi (c): $1 \times 10^{-2}$

  • Milli (m): $1 \times 10^{-3}$

Metric Conversion Factors

Conversions require knowledge of relationships between units.

• Unit Equation: 1 kilometer = 1000 meters

• Example: 325 mg x (1 g / 1000 mg) = 0.325 g

Metric–English Conversions

• 1 inch = 2.54 cm

• 1 pound = 454 g

• 1 quart = 946 mL

• Example: 2.0 oz x (1 lb / 16 oz) x (454 g / 1 lb) = 57 g

The Percent Concept

Percent compares a part to the whole.

• Formula: $\text{Percent} = \frac{\text{Part}}{\text{Whole}} \times 100\%$

• Example: % copper = $\frac{3.80}{3.80 + 1.27} \times 100\% = 75.1\%$

Volume Calculations

• Formula: $\text{Volume} = \text{length} \times \text{width} \times \text{height}$

• Example: 3 cm x 2 cm x 1 cm = 6 cm3

Volume by Displacement

Used for irregular objects; volume equals the amount of liquid displaced.

Density Concept

• Definition: Density is mass per unit volume.

• Formula: $d = \frac{m}{v}$

• Units: g/mL, g/cm3, g/L

• Example: $d = \frac{39.75}{25.0} = 1.59$ g/mL

Temperature Conversions

• Fahrenheit to Celsius: $\left(\text{F} - 32\right) \times \frac{100}{180} = \text{C}$

• Celsius to Kelvin: $\text{C} + 273 = \text{K}$

• Example: -196°C + 273 = 77 K

CHAPTER 3 – Matter & Energy

Physical States of Matter

Matter is anything with mass and volume. It exists in three physical states: solid, liquid, and gas.

• Solid: Definite shape and volume

• Liquid: Definite volume, variable shape

• Gas: Variable shape and volume

Elements, Compounds, and Mixtures

• Element: Cannot be broken down by chemical reaction

• Compound: Combination of elements with predictable properties

• Mixture: Combination of substances; can be homogeneous (uniform) or heterogeneous (non-uniform)

• Alloy: Homogeneous mixture of metals (e.g., 14k gold)

Names & Symbols of Elements

• 81 stable elements in nature

• Each element has a one- or two-letter symbol

Metals, Non-metals, & Semi-metals

• Metals: Left side of periodic table

• Non-metals: Right side of periodic table

• Semi-metals: Intermediate properties

Compounds & Chemical Formulas

• Law of Definite Composition: Elements in a compound are in constant proportions

• Molecule: Collection of atoms

• Chemical Formula: Indicates elements and their ratios (e.g., H2SO4 has 7 atoms)

Physical & Chemical Properties

• Physical Properties: Observed without changing composition (e.g., boiling point, density)

• Chemical Properties: Describe reactivity with other substances

Physical & Chemical Changes

• Physical Change: No change in composition (e.g., melting ice)

• Chemical Change: New substances formed (e.g., burning paper)

Law of Conservation of Mass

• Mass of reactants equals mass of products

• Example: 2.430g Mg + 1.600g O2 → 4.030g MgO

Potential & Kinetic Energy

• Potential Energy (PE): Stored energy

• Kinetic Energy (KE): Energy of motion

Property	Solid	Liquid	Gas
Kinetic Energy	Very low	High	Very high
Movement	Insignificant	Restricted	Unrestricted

Conservation of Energy

• Energy cannot be created or destroyed, only converted

• Exothermic: Reactants have more PE than products; heat released

• Endothermic: Reactants have less PE than products; heat absorbed

• Formula: $E = mc^2$

CHAPTER 4 – Models of the Atom

Dalton's Model

• Atoms of different elements combine in whole number ratios to form compounds

• Atoms can combine in more than one ratio to form different compounds

Thomson's Model

• Introduced subatomic particles: electrons (negative), protons (positive)

Energy Levels and Sublevels

• Energy levels are numbered; sublevels are lettered (s, p, d, f)

• Electron limits:

  • s: 2 electrons

  • p: 6 electrons

  • d: 10 electrons

  • f: 14 electrons

• Example: Fluorine (F): 1s2 2s2 2p5

Electron Configurations

• Electrons fill orbitals from lowest to highest energy

• Example: Sodium (Na): 1s2 2s2 2p6 3s1

CHAPTER 5 – The Periodic Table

Classification and Period Law

• Elements arranged by increasing atomic number

• Properties repeat periodically

Groups & Periods

• Group: Vertical column

• Period: Horizontal row

• Blocks: s-block, p-block (main group), d-block (transition), f-block (rare earth)

Periodic Trends

• Atomic size: Decreases left to right; increases top to bottom

Electron Configurations (Core Notation)

• Example: Sodium (Na): [Ne] 3s1

• Example: Phosphorus (P): [Ne] 3s2 3p3

CHAPTER 6 – Language of Chemistry

Classification of Compounds

• Inorganic Compounds: Do not contain carbon (exceptions: CO2, CO32-, HCO3-)

• Types:

  • Binary Ionic: Metal + non-metal (e.g., NaCl)

  • Ternary Ionic: Metal + non-metal + another element (e.g., KNO3)

  • Binary Molecular: Two non-metals (e.g., H2O)

  • Binary Acid: Hydrogen + non-metal (e.g., HCl (aq))

  • Ternary Oxyacid: Hydrogen + oxygen + non-metal (e.g., HNO3 (aq))

Classifying Ions

• Cation: Positive charge

• Anion: Negative charge

• Monoatomic: Single atom

• Polyatomic: Group of atoms with overall charge

Naming Ions and Compounds

• Metal ions: Named from parent metal (e.g., Na+: Sodium ion)

• Transition metals: Use Roman numerals for charge (e.g., Fe2+: Iron(II))

• Non-metal ions: Use "-ide" suffix (e.g., Cl-: Chloride)

• Formulas: Balance charges (e.g., Mg2+ + Br- → MgBr2)

Polyatomic Ions

• Group of atoms with overall charge (e.g., SO42-)

Writing Chemical Formulas

• Balance cations and anions (e.g., Ca2+ + Br- → CaBr2)

• Include polyatomic ions as needed (e.g., MgSO4)

Binary Ionic Compounds

• With transition metals, specify charge (e.g., Fe2O3: Iron(III) oxide)

• Cation named first

Ternary Ionic Compounds

• Three elements; cation first (e.g., CaCO3: Calcium carbonate)

Binary Molecular Compounds

• Two non-metals; use Greek prefixes (e.g., CO2: Carbon dioxide)

• Second element gets "-ide" suffix

Binary Acids

• Aqueous hydrogen + non-metal; use "hydro-" prefix and "-ic acid" suffix (e.g., HCl (aq): Hydrochloric acid)

CHAPTER 7 – Chemical Reactions

Evidence of Chemical Reaction

• Gas produced

• Insoluble solid formed

• Color change

• Energy change

Writing Chemical Equations

• Reactants → Products

• Physical states indicated: (g), (l), (s), (aq)

• Catalyst and heat shown as needed

• Diatomic elements: H2, O2, N2, F2, Cl2, Br2, I2

Balancing Chemical Equations

• Conservation of mass: atoms on both sides must be equal

• Use coefficients to balance

• Example: 4Al + 3O2 → 2Al2O3

Classification of Chemical Reactions

• Combination: A + Z → AZ

• Decomposition: AZ → A + Z

• Single Replacement: A + BZ → AZ + B

• Double Replacement: AX + BZ → AZ + BX

• Neutralization: HX + BOH → H2O + BX

Activity Series & Single Replacement

• More reactive metals replace less reactive metals

• Example: Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)

Solubility Rules & Double Replacement

• Solubility rules determine if a precipitate forms

• Example: 2AgNO3(aq) + Na2CO3(aq) → Ag2CO3(s) + 2NaNO3(aq)

Neutralization Reactions

• Acid + base → salt + water

• Example: HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq)

CHAPTER 8 – The Mole Concept

Avogadro’s Number

• 1 mole = $6.02 \times 10^{23}$ particles

• Allows conversion between mass and number of particles

Mole Calculations

• Example: 1.50 moles x $6.02 \times 10^{23}$ = $9.03 \times 10^{23}$ atoms

• Convert atoms/molecules to moles and vice versa

Molar Mass

• Definition: Mass of 1 mole of a substance

• Formula: $\text{Molar mass} = \text{sum of atomic masses}$

• Example: NH3: 17.04 g/mol

Mole Calculation II

• Example: Mass of 2.01 x 1022 atoms S: $2.01 \times 10^{22} \times \frac{1}{6.02 \times 10^{23}} \times 32.07 = 1.07$ g

Molar Volumes

• 1 mole of any gas at STP occupies 22.4 L

• Gas Density Formula: $\text{Density} = \frac{\text{Molar mass}}{\text{Molar volume}}$

• Example: NH3: $\frac{17.04}{22.4} = 0.761$ g/L

Mole Calculation III

• Central unit: 1 mole = $6.02 \times 10^{23}$ particles = molar mass (g/mol) = 22.4 L (gas at STP)

• Example: Mass of 3.36 L O3: $3.36 \times \frac{1}{22.4} \times 48.0 = 7.20$ g

Percent Composition

• Mass % of each element in a compound

• Formula: $\text{Mass \%} = \frac{\text{mass of element}}{\text{total mass}} \times 100\%$

Empirical Formula

• Simplest whole-number ratio of elements

• Convert masses to moles, then find ratio

• Example: Benzene: 92.2 g C / 12.01 = 7.68 mol; 7.83 g H / 1.01 = 7.75 mol → ratio ≈ 1:1 → CH

Molecular Formula

• Actual number of atoms in a molecule

• Formula: $\text{Molecular formula} = n \times \text{Empirical formula}$

• Example: Benzene: $78 \div 13 = 6$ → C6H6

CHAPTER 9 – Chemical Equation Calculations

Interpreting Chemical Equations

• Coefficients represent moles, liters (for gases), or grams

• Example: 2NO + O2 → 2NO2

Mole-Mole Relationship

• Use mole ratios as conversion factors

• Example: 2.25 mol O2 x (1 mol N2 / 1 mol O2) = 2.25 mol N2

Types of Stoichiometry Problems

• Mass-mass: mass of reactant → mass of product

• Mass-volume: mass of reactant → volume of gas product

• Volume-volume: volume of gas reactant → volume of gas product

Mass-Mass Problems

• Calculate unknown mass from known mass using molar mass and mole ratios

• Example: 1.25g HgO x (1 mol HgO / 216.59g) x (2 mol Hg / 2 mol HgO) x (200.59g Hg / 1 mol Hg) = 1.16g Hg