General Chemistry Fundamentals: Matter, Measurement, Compounds, Solutions, and Redox

Chapter 1: Matter and Measurement

States of Matter

Matter exists in different physical forms called states or phases. Each state has distinct properties regarding shape and volume.

• Solids: Definite shape and volume.

• Liquids: Definite volume, but no definite shape.

• Gases: No definite shape or volume.

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: Any single element or compound.

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

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

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

Properties of Matter

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

• Physical Property/Change: No new substances 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: Standard units for scientific measurement.

  • Mass: kilogram (kg)

  • Length: meter (m)

  • Time: second (s)

  • Temperature: kelvin (K)

  • Amount: mole (mol)

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

• Metric Prefixes: Used to express multiples or fractions of units.

  Prefix	Symbol	Factor
  Giga	G	109
  Mega	M	106
  Kilo	k	103
  Centi	c	10-2
  Milli	m	10-3
  Micro	μ	10-6
  Nano	n	10-9
  Pico	p	10-12
  Femto	f	10-15
  Atto	a	10-18

Temperature Conversions

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

• Celsius to Fahrenheit: $^\circ F = (9/5)\times ^\circ C + 32$

• Fahrenheit to Celsius: $^\circ C = (5/9)\times (^\circ F - 32)$

Density Calculations

• Density Formula: $d = \frac{m}{V}$

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

Chapter 2: Atoms, Molecules, and Ions

Types of Compounds

• Ionic Compounds: Composed of metals and nonmetals (e.g., NaCl).

• Molecular Compounds: Composed of nonmetals (e.g., CO2).

• Acids: Compounds that release H+ ions in water (e.g., HCl, H2SO4).

Naming Ionic Compounds

1. Name the cation (metal or polyatomic cation).

2. State the metal's oxidation state as a Roman numeral in parentheses (if applicable).

3. Name the non-metal with an -ide ending (or name the polyatomic anion).

• Examples: NaCl (sodium chloride), CuO (copper(II) oxide), FeCl3 (iron(III) chloride).

Naming Molecular Compounds

1. Give the numerical prefix for the first element (omit if only one).

2. Name the first element.

3. Give the numerical prefix for the second element.

4. Name the second element with the -ide suffix.

• Prefixes: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-

• Examples: CO (carbon monoxide), CO2 (carbon dioxide), N2O4 (dinitrogen tetroxide).

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).

Empirical and Molecular Formulas

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

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

• Example: C6H12O6 (molecular), CH2O (empirical).

Chapter 3: Stoichiometry

Percent Composition

• Percent by Mass: $\%\ \text{element} = \frac{\text{mass of element in 1 mol compound}}{\text{molar mass of compound}} \times 100$

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

Empirical Formula from Percent Composition

• Convert percent to grams (assume 100 g sample).

• Convert grams to moles for each element.

• Divide by smallest number of moles to get ratio.

• Example: Compound with 80.0% C and 20.0% H by mass: empirical formula is CH3.

Chapter 4: Aqueous Reactions and Solution Stoichiometry

Solutions and Solubility

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

• Solvent: Substance present in greater amount (usually liquid).

• Solute: Substance dissolved in the solvent.

Electrolytes

• Electrolytes: Substances that dissociate into ions in water, conducting electricity.

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

• Weak Electrolytes: Dissociate partially (e.g., CH3COOH, NH3).

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

Solubility Rules for Ionic Compounds in Water

Soluble Compounds	Exceptions	Insoluble Compounds	Exceptions
Compounds containing NO3-, CH3COO-, Cl-, Br-, I-, SO42-	Ag+, Hg22+, Pb2+ (for Cl-, Br-, I-); Sr2+, Ba2+, Pb2+, Ca2+, Hg22+ (for SO42-)	Compounds containing CO32-, PO43-, OH-	NH4+, alkali metal cations (Li+, Na+, K+, etc.)

Chapter 5: Thermochemistry

Endothermic and Exothermic Processes

• Endothermic: Absorbs heat from surroundings (e.g., melting ice).

• Exothermic: Releases heat to surroundings (e.g., combustion).

Enthalpy ($\Delta H$)

• Definition: The heat content of a system at constant pressure.

• Formula: $\Delta H = H_{\text{products}} - H_{\text{reactants}}$

Chapter 4/20: Redox Reactions and Electrochemistry

Oxidation and Reduction

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

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

Determining Oxidation States

1. Atoms in their elemental form: 0

2. Monatomic ions: charge of the ion

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

4. Hydrogen: +1 (with nonmetals), -1 (with metals)

5. Fluorine: always -1

6. Other halogens: -1 (unless bonded to oxygen or more electronegative elements)

7. The sum of oxidation numbers equals the charge of the molecule or ion.

Activity Series

The activity series ranks metals by their ability to be oxidized. A metal higher in the series will displace a metal lower in the series from solution.

Single Displacement Reactions

• General Form: $A + BC \rightarrow AC + B$

• Example: $Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s)$

Example Redox Reaction

• $2Mg(s) + O_2(g) \rightarrow 2MgO(s)$

• Mg is oxidized (0 to +2), O is reduced (0 to -2).

Additional info:

• Some content inferred and expanded for clarity and completeness, such as the full solubility rules and more detailed examples.

• Not all chapters from the table of contents are covered in the provided images; notes focus on the first several chapters and key foundational concepts.