General Chemistry I: Fundamental Concepts and Applications

Chapter 1: Matter and Measurement

States of Matter

Matter exists in different physical forms called states or phases. The three primary states are:

• Solid: Definite shape and volume.

• Liquid: Definite volume but no definite shape.

• Gas: 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: Consist of a single element or compound.

• Mixtures: Physical 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 of 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
  tera	T	1012
  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: $^\circ F = \frac{9}{5}(^\circ C) + 32$

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

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

Elements, Compounds, and Mixtures

• Element: A pure substance consisting of only one type of atom.

• Compound: A substance composed of two or more elements in fixed proportions.

• Mixture: A physical blend of two or more substances.

Naming Ionic Compounds

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

2. State the metal's oxidation state as a Roman numeral in parentheses (except Group 1/2 metals, Ag+, Zn2+, Cd2+).

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

4. Examples: NaCl (sodium chloride), CuCl (copper(I) chloride), FeCl2 (iron(II) 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.

5. Examples: CO (carbon monoxide), CO2 (carbon 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
HNO2 (nitrous acid)	HClO (hypochlorous acid)
HNO3 (nitric acid)	HClO2 (chlorous acid)
	HClO3 (chloric acid)
	HClO4 (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: The percentage by mass of each element in a compound.

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

• Example: For CaCO3, the percent of C 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. Find the empirical formula.

  • 80.0 g C / 12.01 g/mol = 6.66 mol C

  • 20.0 g H / 1.008 g/mol = 19.8 mol H

  • Ratio: 6.66:19.8 ≈ 1:3 → CH3

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 liquid).

• Solute: The component dissolved in the solvent.

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

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

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

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

Solubility Rules for Ionic Compounds in Water

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

Chapter 5: Thermochemistry

Additional info: Not directly visible in the images, but typically includes energy, work, heat, and the first law of thermodynamics.

Chapter 6: Electronic Structure

Additional info: Not directly visible, but would cover the electromagnetic spectrum, quantization of energy, and atomic orbitals.

Chapter 7: Periodic Properties of the Elements

Additional info: Not directly visible, but would include trends such as atomic radius, ionization energy, and electronegativity.

Chapter 8: Bonding and Lewis Structures

Formal Charge

• Formal Charge Formula: $\text{Formal Charge} = \text{Valence electrons} - [\text{nonbonding electrons} + \frac{1}{2}\text{bonding electrons}]$

Chapter 9: Molecular Geometry and Bonding

Additional info: Would include VSEPR theory and molecular shapes.

Chapter 10: Gases

Additional info: Would include gas laws such as Boyle's, Charles's, and the ideal gas law.

Chapter 11: Intermolecular Forces, Liquids, and Solids

Additional info: Would include types of intermolecular forces and properties of liquids and solids.

Chapter 12: Solutions and Colligative Properties

Additional info: Would include concentration units and colligative properties such as boiling point elevation.

Chapter 13: Chemical Kinetics

Additional info: Would include rate laws and factors affecting reaction rates.

Chapter 14: Chemical Equilibrium

Additional info: Would include the equilibrium constant and Le Chatelier's principle.

Chapter 15: Acid-Base Equilibria

Additional info: Would include definitions of acids and bases, pH, and titrations.

Chapter 16: Buffers, Titrations, and Solubility Equilibria

Additional info: Would include buffer solutions and solubility product constants.

Chapter 17: Thermodynamics

Additional info: Would include entropy, enthalpy, and Gibbs free energy.

Chapter 18: 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: oxidation state = 0.

2. Monatomic ions: oxidation state = ion charge.

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

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

5. Fluorine: always -1.

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

7. The sum of oxidation numbers in a compound or polyatomic ion equals the overall charge.

Activity Series

• Ranks metals by their tendency to be oxidized (lose electrons).

• More active metals can displace less active metals from solutions.

Single Displacement Reactions

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

• Occurs if A is more active than B in the activity series.

Additional Info

• Some chapters (e.g., Thermochemistry, Electronic Structure, Periodic Properties, etc.) are referenced in the table of contents but not detailed in the visible notes. For a complete study guide, consult the full textbook or lecture materials.