GOB Chemistry Study Guide: Matter, Measurements, Atomic Structure, and Chemical Bonding (Chapters 1-4)

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This section introduces the language of chemistry, focusing on how matter is classified and how measurements are made using numbers and units.

States of Matter: Matter exists in three common states: solid, liquid, and gas. Each state has distinct properties such as shape and volume.
Temperature and Pressure: These factors affect the state of matter. For example, water can exist as ice, liquid water, or steam depending on temperature and pressure.
Classification Flowchart: Matter can be classified as a pure substance (uniform chemical composition) or a mixture (combination of substances).
Elements vs. Compounds: Pure substances are further classified as elements (single type of atom) or chemical compounds (two or more types of atoms chemically bonded).
Physical vs. Chemical Change: Physical changes do not alter chemical composition (e.g., melting ice), while chemical changes do (e.g., cooking an egg).

Accurate measurement is essential in chemistry. This section covers the basic units and how to use them.

SI Units: Standard units include mass (kg), length (m), volume (m3 or L), temperature (K), and time (s).
Metric Prefixes: Common prefixes include kilo-, centi-, milli-, micro-, and nano-. These are used to express measurements in scientific notation.
Scientific Notation: Used to convert between large and small numbers efficiently.

Significant figures indicate the precision of a measurement. Calculations must follow rules for significant figures.

Determining Significant Figures: The number of significant digits in a measurement reflects its precision.
Rules for Calculations:
- Multiplication/Division: The result has the same number of significant figures as the measurement with the fewest significant figures.
- Addition/Subtraction: The result has the same number of decimal places as the measurement with the fewest decimal places.
Factor-Label Method: Used to convert units.
Density: Used to relate mass and volume.

This chapter introduces the fundamental building blocks of matter and their organization.

Subatomic Particles: Atoms consist of protons (p), neutrons (n), and electrons (e-).
Atomic Symbol: Used to represent an atom, showing the number of protons, neutrons, and electrons.

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

Isotopic Mass: The mass of a specific isotope, measured in atomic mass units (amu).
Atomic Weight: The weighted average mass of all naturally occurring isotopes of an element.

The periodic table organizes elements by increasing atomic number and groups elements with similar properties.

Groups: Vertical columns; elements in the same group have similar chemical properties.
Periods: Horizontal rows.
Main Groups:
- Group 1: Alkali Metals
- Group 2: Alkaline Earth Metals
- Group 17: Halogens
- Group 18: Noble Gases

Electron configuration describes how electrons are arranged in an atom.

Shells and Orbitals: Electrons are organized into shells (n=1, 2, 3...), subshells (s, p, d, f), and orbitals.
Valence Shell: The outermost shell; the number of valence electrons determines chemical reactivity.
Electron-Dot Symbols: Used to represent valence electrons for main group elements.

This chapter explains how atoms gain or lose electrons to form ions and compounds.

Octet Rule: Atoms tend to gain, lose, or share electrons to achieve a full valence shell (usually 8 electrons).
Cations: Positively charged ions formed when atoms lose electrons.
Anions: Negatively charged ions formed when atoms gain electrons.

Ionic bonds are formed by the electrostatic attraction between oppositely charged ions.

Formation: Occurs between metals and nonmetals (e.g., NaCl).
Naming: Main group cations are named by the element; transition metals use Roman numerals for charge; anions are named by replacing the element’s ending with “-ide”.
Polyatomic Ions: Common examples include OH-, CO32-, SO42-, PO43-, NH4+.

Covalent bonds are formed when atoms share electrons to achieve an octet. Molecules are held together by covalent bonds.

Lewis Dot Structures: Show all atoms, bonding pairs, and lone pairs of electrons in a molecule.
Example: Carbon has 4 valence electrons and forms 4 bonds.

The Valence Shell Electron Pair Repulsion (VSEPR) model predicts molecular shape based on electron cloud repulsion.

Molecular Geometry: Determined by the number of electron clouds and lone pairs around the central atom.
Bond Angles:
- 2 clouds: Linear (180°)
- 3 clouds: Trigonal planar (120°), Bent (<120°)
- 4 clouds: Tetrahedral (109.5°), Pyramidal (<109.5°), Bent (<109.5°)

Electronegativity measures an atom’s ability to attract electrons in a bond. It determines bond polarity and molecular polarity.

Bond Polarity:
- Non-polar covalent: Electronegativity difference < 0.5
- Polar covalent: Difference 0.5 – 2.0
- Ionic: Difference > 2.0
Molecular Polarity: Determined by bond polarity and molecular shape. Use partial charges and dipole arrows to represent direction.
Example: CO2 is linear and nonpolar; BF3 is trigonal planar and nonpolar due to symmetry.

This table summarizes the main types of chemical bonds and their properties.

Bond Type	Electronegativity Difference	Example	Properties
Non-polar Covalent	< 0.5	Cl2, H2	Equal sharing of electrons
Polar Covalent	0.5 – 2.0	H2O, NH3	Unequal sharing of electrons
Ionic	> 2.0	NaCl, KBr	Transfer of electrons