General Chemistry Exam 1 Comprehensive Study Guide

Chapter 1: Introduction to Chemistry and the Periodic Table

Elements, Compounds, and Mixtures

Understanding the basic types of matter is foundational in chemistry. Elements, compounds, and mixtures differ in composition and properties.

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

• Compound: A substance formed from two or more elements chemically bonded in fixed proportions. Example: Water (H2O).

• Mixture: A combination of two or more substances not chemically bonded. Example: Air.

Main Groups of the Periodic Table

The periodic table organizes elements into groups with similar properties. Main groups include:

• Group 1: Alkali metals

• Group 2: Alkaline earth metals

• Groups 13-18: Main group elements (including halogens and noble gases)

Transition metals are found in groups 3-12.

Physical vs. Chemical Changes

Distinguishing between physical and chemical changes is essential:

• Physical Change: Alters the form but not the chemical identity (e.g., melting ice).

• Chemical Change: Produces new substances (e.g., rusting iron).

Periodic Table Structure

The periodic table is arranged in vertical columns (groups) and horizontal rows (periods). Elements in the same group share similar chemical properties.

• Groups: Vertical columns (e.g., Group 1: Alkali metals)

• Periods: Horizontal rows (e.g., Period 2: Li, Be, B, C, N, O, F, Ne)

Groups of the periodic table have similar physical and chemical properties.

Identifying Metals, Nonmetals, and Metalloids

Given an element, you should be able to classify it as a metal, nonmetal, or metalloid based on its position and properties.

• Metals: Good conductors, malleable, shiny (e.g., Fe, Cu)

• Nonmetals: Poor conductors, brittle, dull (e.g., O, N)

• Metalloids: Properties intermediate between metals and nonmetals (e.g., Si, As)

Chapter 2: Scientific Notation and Metric System

Scientific Notation

Scientific notation expresses numbers as a product of a coefficient and a power of ten.

• Example: $6,500 = 6.5 \times 10^3$

Metric Units and Conversions

The metric system is used for scientific measurements. Key units include:

• Length: meter (m)

• Mass: kilogram (kg)

• Volume: liter (L)

Conversion example: $1 \text{ kg} = 1000 \text{ g}$

Significant Figures

Significant figures reflect the precision of a measurement.

• Rules for determining significant figures:

  • All nonzero digits are significant.

  • Zeros between nonzero digits are significant.

  • Leading zeros are not significant.

  • Trailing zeros in a decimal number are significant.

• When multiplying/dividing, the result should have the same number of significant figures as the measurement with the fewest significant figures.

English-Metric Conversions

Be able to convert between English and metric units using conversion factors.

• Example: $1 \text{ inch} = 2.54 \text{ cm}$

Chapter 3: Atomic Structure and Isotopes

Structure of the Atom

Atoms consist of subatomic particles: protons, neutrons, and electrons.

• Proton: Positively charged, found in nucleus

• Neutron: Neutral, found in nucleus

• Electron: Negatively charged, found in electron cloud

Mass Number and Isotopes

The mass number is the sum of protons and neutrons in an atom.

• Formula: $\text{Mass Number} = \text{Number of Protons} + \text{Number of Neutrons}$

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

Electron Configuration and Orbital Diagrams

Electron configuration describes the arrangement of electrons in an atom.

• Example: $\text{Oxygen: } 1s^2\ 2s^2\ 2p^4$

• Orbital diagrams show the distribution of electrons among orbitals.

Recognize the electron dot structures (Lewis structures) for main group elements.

Chapter 5: Ionic and Covalent Compounds

Ionic Compounds

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

• Cation: Positively charged ion (metal loses electrons)

• Anion: Negatively charged ion (nonmetal gains electrons)

• Valence electrons determine the charge of ions.

Polyatomic Ions

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

Nonmetal	Formula	Name
Carbon	CO32-	Carbonate
Carbon	HCO3-	Hydrogen carbonate (bicarbonate)
Nitrogen	NO3-	Nitrate
Oxygen	OH-	Hydroxide
Phosphorus	PO43-	Phosphate
Sulfur	SO42-	Sulfate

Covalent Compounds

Covalent compounds are formed by the sharing of electrons between nonmetals.

• Recognize the names and formulas of common covalent compounds.

Chapter 6: Molecular Structure and Bonding

Diatomic Molecules

Certain elements naturally exist as diatomic molecules (two atoms bonded together).

• Examples: H2, N2, O2, F2, Cl2, Br2, I2

Lewis Structures and Valence Electrons

Lewis structures represent the arrangement of electrons in molecules.

• Valence electrons determine the number of covalent bonds an element can form.

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

• Exceptions to the octet rule exist (e.g., molecules with odd numbers of electrons, expanded octets).

Molecular Geometry

The shape of a molecule is determined by the number of bonds and lone pairs around the central atom.

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

Electronegativity and Bond Polarity

Electronegativity is the tendency of an atom to attract electrons in a bond.

• Polar covalent bonds form when atoms have different electronegativities.

• Nonpolar covalent bonds form when atoms have similar electronegativities.

• Ionic bonds form when the difference in electronegativity is large.

Hydrogen bonds form between molecules containing H bonded to highly electronegative atoms (N, O, F).

Distinguishing Bond Types

• Ionic Compounds: Metal + Nonmetal, electron transfer

• Polar Covalent Compounds: Nonmetals, unequal sharing of electrons

• Nonpolar Covalent Compounds: Nonmetals, equal sharing of electrons

Appendices

Appendix I: Periodic Table

The periodic table organizes elements by atomic number and groups with similar properties. Refer to the table for element classification and properties.

Appendix III: Polyatomic Ions

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

Appendix V: Electronegativity

Electronegativity values vary across the periodic table. Fluorine is the most electronegative element.

• Electronegativity difference determines bond type:

  • Large difference: Ionic bond

  • Moderate difference: Polar covalent bond

  • Small or zero difference: Nonpolar covalent bond

Summary Table: Key Concepts

Concept	Definition	Example
Element	Pure substance of one type of atom	O2
Compound	Substance of two or more elements chemically bonded	H2O
Mixture	Physical blend of substances	Air
Isotope	Atoms of same element, different neutrons	C-12, C-14
Polyatomic Ion	Ion of multiple atoms	SO42-
Electronegativity	Atom's ability to attract electrons	F is highest

Additional info: The study guide covers foundational topics for a first exam in General Chemistry, including atomic structure, periodic table organization, chemical bonding, molecular geometry, and measurement systems. Tables have been recreated for polyatomic ions and molecular geometry. The periodic table and electronegativity chart are referenced for classification and trends.