General Chemistry I: Core Concepts and Study Guide

General Chemistry Fundamentals

Types of Matter

Matter is anything that has mass and occupies space. It can be classified based on its composition and properties.

• Pure Substances: Consist of only one type of particle. Examples: elements (e.g., O2), compounds (e.g., H2O).

• Mixtures: Combinations of two or more substances. Can be:

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

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

Periodic Table

The periodic table organizes elements by increasing atomic number and recurring chemical properties.

• Periods: Horizontal rows.

• Groups: Vertical columns; elements in the same group have similar properties.

• Metals, Nonmetals, Metalloids: Classified by physical and chemical properties.

• Main-group vs. Transition Elements: Main-group elements are in groups 1, 2, and 13-18; transition elements are in groups 3-12.

• Element Location: Each element's position provides information about its properties.

Metric Prefixes and Conversions

Metric prefixes indicate multiples or submultiples of units (e.g., kilo-, milli-). Dimensional analysis is used for unit conversions.

• Common Prefixes: kilo- (103), centi- (10-2), milli- (10-3), etc.

• Dimensional Analysis: Converts from one unit to another using conversion factors.

Significant Figures

Significant figures reflect the precision of a measured quantity.

• Rules: All nonzero digits are significant; zeros between nonzero digits are significant; leading zeros are not significant; trailing zeros are significant if after a decimal point.

• Rounding: Round to the correct number of significant figures based on the operation performed.

Types of Measurements and Units

Common measurements include mass, volume, density, and temperature.

• SI Units: Mass (kg), volume (L), temperature (K), etc.

• Density: $d = \frac{m}{V}$

• Percent: $\text{Percent} = \frac{\text{part}}{\text{whole}} \times 100$

• Temperature Conversions:

  • $T_F = 1.8T_C + 32$

  • $T_C = \frac{T_F - 32}{1.8}$

  • $T_K = T_C + 273$

Energy

Energy is the capacity to do work. It exists in various forms, such as potential and kinetic energy.

• Potential Energy: Stored energy due to position.

• Kinetic Energy: Energy of motion.

States of Matter

Matter exists as solids, liquids, or gases, each with distinct properties.

• Solids: Definite shape and volume.

• Liquids: Definite volume, indefinite shape.

• Gases: Indefinite shape and volume.

Chemical and Physical Changes

Chemical changes result in new substances; physical changes do not alter the chemical identity.

• Chemical Change: Rusting, burning.

• Physical Change: Melting, boiling.

Balancing Chemical Equations

Balancing ensures the same number of each atom on both sides of a chemical equation, reflecting the law of conservation of mass.

Atomic Structure and Nuclear Chemistry

Structure of the Atom

Atoms consist of protons, neutrons, and electrons.

• Protons: Positively charged, in the nucleus.

• Neutrons: Neutral, in the nucleus.

• Electrons: Negatively charged, orbit the nucleus.

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons.

• Atomic Number (Z): Number of protons.

• Mass Number (A): Number of protons + neutrons.

• Isotope Notation: $^{A}_{Z}\text{X}$

Nuclear Equations and Decay

Nuclear reactions involve changes in the nucleus, such as alpha, beta, and gamma decay.

• Alpha Decay: Emission of a helium nucleus ($^4_2\text{He}$).

• Beta Decay: Emission of an electron ($^0_{-1}\text{e}$).

• Gamma Emission: Emission of high-energy photons ($\gamma$).

• Half-life: Time required for half of a radioactive sample to decay.

Chemical Bonding and Compounds

Electron Arrangements and the Octet Rule

Electron configuration describes the arrangement of electrons in an atom. The octet rule states that atoms tend to gain, lose, or share electrons to achieve eight valence electrons.

• Valence Electrons: Electrons in the outermost shell, determine chemical reactivity.

• Group Number: Indicates the number of valence electrons for main-group elements.

Ionic Compounds

Ionic compounds form from the transfer of electrons between metals and nonmetals, resulting in cations and anions.

• Cation: Positively charged ion (metal loses electrons).

• Anion: Negatively charged ion (nonmetal gains electrons).

• Polyatomic Ions: Ions composed of multiple atoms.

• Naming: Name the cation first, then the anion.

Common Polyatomic Ions:

Main Element	Formula of Ion*	Name of Ion
Hydrogen	OH-	Hydroxide
Hydrogen	H3O+	Hydronium
Nitrogen	NH4+	Ammonium
Nitrogen	NO3-	Nitrate
Nitrogen	NO2-	Nitrite
Chlorine	ClO3-	Chlorate
Carbon	CO32-	Carbonate
Carbon	HCO3-	Hydrogen carbonate (bicarbonate)
Carbon	CH3COO-	Acetate
Sulfur	SO42-	Sulfate
Sulfur	SO32-	Sulfite
Phosphorus	PO43-	Phosphate
Phosphorus	HPO42-	Hydrogen phosphate
Phosphorus	H2PO4-	Dihydrogen phosphate
Manganese	MnO4-	Permanaganate

Covalent (Molecular) Compounds

Covalent compounds form when nonmetals share electrons. They are named using prefixes to indicate the number of atoms.

• Covalent Bond: Shared pair of electrons between atoms.

• Lewis Structures: Diagrams showing bonding and lone pairs.

• Molecular Shapes (VSEPR Theory): Predicts 3D shape based on electron group repulsions.

Electron Groups	Electron Group Geometry	Bonded Atoms	Lone Pairs	Angle	Molecular Shape	Example
2	Linear	2	0	180°	Linear	CO2
3	Trigonal planar	3	0	120°	Trigonal planar	H2CO
3	Trigonal planar	2	1	120°	Bent	SO2
4	Tetrahedral	4	0	109°	Tetrahedral	CH4
4	Tetrahedral	3	1	109°	Trigonal pyramidal	NH3
4	Tetrahedral	2	2	109°	Bent	H2O

Calculations in Chemistry

Moles, Molar Mass, and Avogadro's Number

The mole is a counting unit in chemistry, relating mass to number of particles.

• Mole: $1\ \text{mol} = 6.02 \times 10^{23}$ particles (atoms, molecules, etc.)

• Molar Mass (MM): Mass of one mole of a substance (g/mol).

• Conversions:

  • From mass to moles: $\text{moles} = \frac{\text{mass}}{\text{MM}}$

  • From moles to mass: $\text{mass} = \text{moles} \times \text{MM}$

  • From moles to particles: $\text{particles} = \text{moles} \times 6.02 \times 10^{23}$

  • From particles to moles: $\text{moles} = \frac{\text{particles}}{6.02 \times 10^{23}}$

Example Problem

How many molecules are in 18 g of water (H2O)?

• Molar mass of H2O = 18 g/mol

• Moles = $\frac{18\ \text{g}}{18\ \text{g/mol}} = 1\ \text{mol}$

• Molecules = $1\ \text{mol} \times 6.02 \times 10^{23} = 6.02 \times 10^{23}$ molecules

Additional info:

• Some context and explanations have been expanded for clarity and completeness.

• Tables have been recreated and summarized for the most common ions and molecular shapes.