Exam 1 Study Guide: Chapters 1, 2, and 3

Overview

This study guide summarizes the main topics and concepts that may be covered on Exam 1 for a GOB Chemistry course, focusing on foundational principles in general, organic, and biological chemistry. Each topic is expanded with definitions, examples, and relevant equations to support effective exam preparation.

Definitions of Terms

Key Chemical Terms

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

• Compound: A substance formed from two or more elements chemically bonded together.

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

• Atom: The smallest unit of an element that retains its chemical properties.

• Molecule: Two or more atoms bonded together.

Example: Water (H2O) is a compound made of hydrogen and oxygen atoms.

Classifying Matter

Types and Classification

• Pure Substances: Elements and compounds with uniform composition.

• Mixtures: Can be homogeneous (uniform throughout) or heterogeneous (distinct phases).

Example: Air is a homogeneous mixture; salad is a heterogeneous mixture.

Properties of Matter

Physical vs. Chemical Properties

• Physical Properties: Characteristics observed without changing the substance (e.g., melting point, density).

• Chemical Properties: Characteristics observed during a chemical change (e.g., reactivity, flammability).

Example: Boiling water is a physical change; rusting iron is a chemical change.

Types of Mixtures

Mixture Classification

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

• Heterogeneous Mixture: Non-uniform composition (e.g., sand and iron filings).

Chemical Reactions vs. Physical Changes

Distinguishing Changes

• Chemical Reaction: Substances are transformed into new substances.

• Physical Change: Change in state or appearance without altering composition.

Example: Burning wood (chemical); melting ice (physical).

Conversions

Unit Conversions and Prefixes

• Use conversion factors to change units (e.g., grams to kilograms).

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

Example:

Scientific Notation

Expressing Large and Small Numbers

• Numbers are written as where and is an integer.

Example:

Formulas and Equations

Chemical Formulas and Calculations

• Chemical Formula: Shows the elements and their ratios in a compound (e.g., CO2).

• Equation Example: Temperature conversion from Fahrenheit to Kelvin:

Percents

Calculating Percent Composition

• Percent by Mass:

Example:

Balancing Chemical Equations

Law of Conservation of Mass

• Reactants and products must have equal numbers of each atom.

• Adjust coefficients to balance equations.

Example:

Periodic Table

Metals, Nonmetals, Periods, and Groups

• Metals: Shiny, conductive, malleable (left side of table).

• Nonmetals: Dull, poor conductors (right side of table).

• Periods: Horizontal rows.

• Groups: Vertical columns; elements in a group share properties.

Subatomic Particles

Structure of the Atom

• Proton: Positively charged particle in nucleus.

• Neutron: Neutral particle in nucleus.

• Electron: Negatively charged particle in electron cloud.

Example: Carbon atom: 6 protons, 6 neutrons, 6 electrons.

Atomic Symbols

Notation for Elements

• Symbol: One or two letters (e.g., Na for sodium).

• Atomic number: Number of protons.

• Mass number: Protons + neutrons.

Example:

Nuclear Decay Reactions

Types of Radioactive Decay

• Alpha Decay: Emission of particle ().

• Beta Decay: Emission of particle ().

• Gamma Decay: Emission of rays (energy).

Half-Life Problems

Radioactive Decay Calculations

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

• Equation:

Valence Electrons and Octet

Electron Configuration

• Valence Electrons: Electrons in the outermost shell.

• Octet Rule: Atoms tend to have 8 electrons in their valence shell.

Example: Neon has a full octet (8 valence electrons).

Ions and Ionic Compounds

Formation and Properties

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

• Ionic Compound: Formed from cations and anions (e.g., NaCl).

Covalent Compounds

Bonding and Structure

• Covalent Bond: Sharing of electron pairs between atoms.

• Molecule: Discrete group of atoms held together by covalent bonds.

Example: is a covalent compound.

Lewis Structures

Drawing Electron Dot Structures

• Show valence electrons as dots around element symbols.

• Represent bonds as shared pairs of dots or lines.

Example: Lewis structure for :

O with two pairs of dots and two single bonds to H.

VSEPR Theory

Predicting Molecular Shapes

• VSEPR: Valence Shell Electron Pair Repulsion theory predicts 3D shapes of molecules.

• Electron pairs repel, determining geometry (e.g., linear, bent, tetrahedral).

Example: is tetrahedral; is bent.

Dipole Moment Arrows and Polarity

Molecular Polarity

• Dipole Moment: Measure of charge separation in a molecule.

• Arrows point from positive to negative end.

• Polar Molecule: Has an uneven distribution of electrons (e.g., ).

Example: is nonpolar; is polar.

Summary Table: Classification of Matter

Additional info: Some explanations and examples have been expanded for clarity and completeness.