General Chemistry I: Foundations, Nomenclature, Stoichiometry, and Aqueous Reactions

Chapter 1: Matter and Measurement

States of Matter

Matter exists in different physical forms, each with distinct properties.

• Solid: Definite shape and volume.

• Liquid: Definite volume, but takes the shape of its container.

• Gas: No definite shape or volume; expands to fill its container.

Classification of Matter

• Elements: Substances that cannot be broken down into simpler substances by chemical means (e.g., Na, H2, S8).

• Compounds: Substances composed of two or more elements chemically combined (e.g., CO2, NaCl).

• Pure Substances: Consist of a single element or compound.

• Mixtures: Combinations of two or more substances (elements and/or compounds).

• Homogeneous Mixture: Uniform composition throughout (e.g., saltwater).

• Heterogeneous Mixture: Non-uniform composition (e.g., salad).

Properties of Matter

• Chemical Property/Change: Involves conversion of substances into different substances (e.g., rusting of iron).

• Physical Property/Change: No new substances are formed (e.g., melting ice).

• Intensive Property: Independent of sample size (e.g., density, boiling point).

• Extensive Property: Dependent on sample size (e.g., mass, volume).

Measurement and Units

• SI Base Units:

  • Mass: kilogram (kg)

  • Length: meter (m)

  • Time: second (s)

  • Temperature: kelvin (K)

  • Amount: mole (mol)

• Derived Units: Formed from base units (e.g., volume: m3, pressure: Pa).

• Metric Prefixes: Used to express multiples or fractions of units (e.g., kilo-, centi-, milli-).

Prefix	Symbol	Factor
kilo	k	103
centi	c	10-2
milli	m	10-3
micro	μ	10-6
nano	n	10-9

Temperature Conversions

• Kelvin to Celsius: $K = ^\circ C + 273.15$

• Celsius to Fahrenheit: $^\circ F = \frac{9}{5}(^\circ C) + 32$

Density Calculations

• Density Formula: $d = \frac{m}{V}$

• Example: A cube with 2.00 cm edges weighs 12.0 g. Density = $\frac{12.0\ g}{(2.00\ cm)^3} = 1.50\ g/cm^3$

Chapter 2: Atoms, Molecules, and Ions

Naming Ionic Compounds

• Name the cation (metal or polyatomic cation).

• State the metal's oxidation state as a Roman numeral in parentheses (for transition metals).

• Name the non-metal with an -ide ending or name the polyatomic anion.

• Examples:

  • NaCl: sodium chloride

  • CuCl: copper(I) chloride

  • CuCl2: copper(II) chloride

  • NH4Br: ammonium bromide

Naming Molecular Compounds

• Use numerical prefixes for the number of atoms (mono-, di-, tri-, etc.).

• Name the first element (omit 'mono-' if only one atom).

• Name the second element with the -ide suffix.

• Examples:

  • CO: carbon monoxide

  • CO2: carbon dioxide

  • SO2: sulfur dioxide

Prefix	Number
mono-	1
di-	2
tri-	3
tetra-	4
penta-	5
hexa-	6
hepta-	7
octa-	8
nona-	9
deca-	10

Naming Acids

• Binary Acids: Hydro- + element + -ic acid (e.g., HCl: hydrochloric acid).

• Oxoacids: Based on polyatomic ions (e.g., HNO3: nitric acid, H2SO4: sulfuric acid).

Binary Acids	Oxoacids
HCl(aq): hydrochloric acid	HClO(aq): hypochlorous acid
HF(aq): hydrofluoric acid	HClO2(aq): chlorous acid
	HClO3(aq): chloric acid
	HClO4(aq): perchloric acid

Empirical and Molecular Formulas

• Empirical Formula: Simplest whole-number ratio of atoms in a compound.

• Molecular Formula: Actual number of atoms of each element in a molecule; may be a multiple of the empirical formula.

Chapter 3: Stoichiometry

Percent Composition

• Percent composition of an element in a compound is calculated as: $\%\ \text{Element} = \frac{\text{mass of element in 1 mol compound}}{\text{molar mass of compound}} \times 100$

• Example: In 300 g of CaCO3, the percent of carbon is: $\frac{12.01}{100.09} \times 100 = 12.0\%$

Empirical Formula Determination

• Given percent composition, convert to grams (assume 100 g sample), then to moles, and find the simplest ratio.

• Example: A compound is 80.0% C and 20.0% H by mass. Empirical formula is CH3.

• If the molecular weight is given, divide by empirical formula mass to find the molecular formula.

Chapter 4: Aqueous Reactions and Solution Stoichiometry

Solutions and Electrolytes

• Solution: Homogeneous mixture of solute(s) dissolved in a solvent.

• Solvent: The component present in greater amount (usually a liquid).

• Solute: The component dissolved in the solvent.

Electrolytes

• Strong Electrolytes: Dissociate completely into ions (e.g., NaCl, HCl, NaOH).

• Weak Electrolytes: Partially dissociate into ions (e.g., CH3COOH, NH3).

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

Type	Examples
Strong Electrolytes	Soluble ionic compounds, strong acids (HCl, HNO3), strong bases (NaOH, KOH)
Weak Electrolytes	Weak acids (CH3COOH), weak bases (NH3)
Nonelectrolytes	Molecular compounds (except acids/bases)

Solubility Rules for Ionic Compounds in Water

Soluble Compounds	Exceptions	Insoluble Compounds	Exceptions
Compounds containing NO3-, CH3COO-	None	CO32-, PO43-	NH4+, alkali metals
Cl-, Br-, I-	Ag+, Hg22+, Pb2+	S2-, OH-	NH4+, alkali metals, Ca2+, Sr2+, Ba2+
SO42-	Sr2+, Ba2+, Hg22+, Pb2+

Chapter 5: Thermochemistry (Selected Concepts)

• Oxidation: Loss of electrons (OIL: Oxidation Is Loss).

• Reduction: Gain of electrons (RIG: Reduction Is Gain).

Determining Oxidation States

• Atoms in their elemental form: oxidation state = 0.

• Monatomic ions: oxidation state = ionic charge.

• Oxygen: usually -2 (except in peroxides: -1).

• Hydrogen: +1 (when bonded to nonmetals), -1 (when bonded to metals).

• Fluorine: always -1.

• Halogens: usually -1, except when bonded to oxygen or other halogens.

• The sum of oxidation numbers in a compound is zero; in a polyatomic ion, it equals the ion's charge.

Activity Series and Displacement Reactions

• The activity series ranks metals by their ability to be oxidized (lose electrons).

• A more active metal will displace a less active metal from solution.

• Example: Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)

Additional info:

• These notes cover foundational topics in General Chemistry I, including matter classification, measurement, nomenclature, stoichiometry, solution chemistry, and introductory redox concepts.

• For more advanced topics (e.g., thermodynamics, equilibrium, kinetics), refer to later chapters in the course.