GOB Chemistry Exam 1 Comprehensive Study Guide

Chapter 1: Introduction to Chemistry and the Periodic Table

Elements, Compounds, and Mixtures

Chemistry distinguishes between pure substances and mixtures. Understanding these differences is foundational for all chemical studies.

• Element: A pure substance consisting of only one type of atom. Example: Oxygen (O2).

• Compound: A substance formed when two or more elements are chemically bonded. Example: Water (H2O).

• Mixture: A combination of two or more substances that are not chemically bonded. Example: Salt water.

Main Groups of the Periodic Table

The periodic table is organized into groups and periods, which help classify elements based on their properties.

• Main group elements: Groups 1-2 and 13-18 (s- and p-blocks).

• Transition metals: Groups 3-12 (d-block).

• Physical vs. Chemical Changes:

  • Physical change: Change in state or appearance without altering composition (e.g., melting ice).

  • Chemical change: Change that produces new substances (e.g., rusting iron).

• Periodic Table Numbering:

  • Groups: Vertical columns (1-18).

  • Periods: Horizontal rows (1-7).

• Similar Properties: Elements in the same group have similar chemical and physical properties.

Identifying Elements

• Given an element, determine if it is a metal, nonmetal, or metalloid.

Chapter 2: Scientific Notation and Metric System

Scientific Notation

Scientific notation is used to express very large or small numbers efficiently.

• Format: $a \times 10^n$ where $1 \leq a < 10$ and $n$ is an integer.

• Example: $3,000 = 3 \times 10^3$

Metric Units and Conversions

• Base units: Meter (length), kilogram (mass), second (time), liter (volume).

• Metric prefixes:

  • Kilo- ($10^3$), centi- ($10^{-2}$), milli- ($10^{-3}$), micro- ($10^{-6}$)

• Conversion: Use conversion factors to change between units. Example: $1 \text{ kg} = 1000 \text{ g}$

Significant Figures

• Digits that carry meaning in a measurement.

• Rules for multiplication/division: The result should have the same number of significant figures as the measurement with the fewest significant figures.

Chapter 3: Atomic Structure

Subatomic Particles

Atoms are composed of protons, neutrons, and electrons.

• Proton: Positively charged, found in nucleus.

• Neutron: Neutral, found in nucleus.

• Electron: Negatively charged, found in electron cloud.

Isotopes and Atomic Mass

• Isotope: Atoms of the same element with different numbers of neutrons.

• Mass number: $\text{Mass number} = \text{Number of protons} + \text{Number of neutrons}$

Electron Configuration and Orbital Diagrams

• Electron configuration shows the distribution of electrons among orbitals. Example: $1s^2 2s^2 2p^6$

• Orbital diagrams use arrows to represent electron spins in orbitals.

Chapter 5: Ionic Compounds and Polyatomic Ions

Ionic Compounds

Ionic compounds are formed from the transfer of electrons between metals and nonmetals.

• Ionic bond: Electrostatic attraction between cations and anions.

• Properties: High melting points, conduct electricity when dissolved in water.

• Electron configuration: Atoms gain or lose electrons to achieve a stable octet.

Polyatomic Ions

Polyatomic ions are charged species composed of two or more atoms covalently bonded.

Nonmetal	Formula	Name
Carbon	CO32-	Carbonate
Carbon	HCO3-	Hydrogen carbonate or bicarbonate
Carbon	CH3COO-	Acetate
Carbon	CN-	Cyanide
Nitrogen	NO3-	Nitrate
Nitrogen	NO2-	Nitrite
Oxygen	OH-	Hydroxide
Phosphorus	PO43-	Phosphate
Phosphorus	HPO42-	Hydrogen phosphate
Phosphorus	H2PO4-	Dihydrogen phosphate
Sulfur	SO42-	Sulfate
Sulfur	SO32-	Sulfite
Sulfur	HSO4-	Hydrogen sulfate or bisulfate

Valence Electrons and Charges

• Valence electrons are the outermost electrons involved in bonding.

• Charges of ions are determined by the loss or gain of electrons to achieve a stable configuration.

Chapter 6: Covalent Compounds and Molecular Structure

Diatomic Molecules

• Certain elements exist naturally as diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2.

Lewis Structures and Bonding

• Lewis structures represent the arrangement of electrons in molecules.

• Octet rule: Atoms tend to have eight electrons in their valence shell.

• Exceptions: Some molecules have expanded octets or fewer than eight electrons.

• Bonding and nonbonding electron pairs are shown as lines and dots, respectively.

Molecular Geometry and VSEPR Theory

The shape of molecules is determined by the number of bonding and lone pairs around the central atom.

Number of Bonds	Number of Lone Pairs	Total Electron Groups	Molecular Geometry	Example	Approximate Bond Angle
2	0	2	Linear	CO2	180°
3	0	3	Trigonal planar	BF3	120°
2	1	3	Bent	SO2	~120°
4	0	4	Tetrahedral	CH4	109.5°
3	1	4	Trigonal pyramidal	NH3	~107°
2	2	4	Bent	H2O	~104.5°

Electronegativity and Bond Polarity

• Electronegativity: The tendency of an atom to attract electrons in a bond.

• Polar covalent bonds form when there is a significant difference in electronegativity between atoms.

• Nonpolar covalent bonds form when atoms have similar electronegativities.

• Ionic bonds form when the difference is very large.

Hydrogen Bonding

• Occurs when hydrogen is bonded to highly electronegative atoms (N, O, F).

• Responsible for unique properties of water and biological molecules.

Appendices

Appendix I: Periodic Table

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

Appendix III: Polyatomic Ions

See table above for common polyatomic ions, their formulas, and names.

Appendix V: Electronegativity

Electronegativity values increase across a period and decrease down a group. Fluorine is the most electronegative element.

Additional info:

• Expanded context and definitions were added for clarity and completeness.

• Tables were reconstructed and formatted for study purposes.