General Chemistry Exam 1 Review: Key Concepts and Takeaways

Classification of Matter

Elements, Compounds, Homogeneous Mixtures, and Heterogeneous Mixtures

Matter can be classified based on its composition and uniformity. Understanding these categories is fundamental in chemistry.

• Element: A pure substance that cannot be broken down into simpler substances by chemical means. Example: O2, Fe.

• Compound: A substance composed of two or more elements chemically combined in fixed proportions. Example: H2O, NaCl.

• Homogeneous Mixture (Solution): A mixture with uniform composition throughout. Example: Saltwater, air.

• Heterogeneous Mixture: A mixture with non-uniform composition; different parts are visibly distinct. Example: Salad, sand in water.

Key identification tips:

• If you can write a chemical formula, it is a compound.

• If it is on the periodic table, it is an element.

• If it is visibly uniform but not a pure substance, it is a homogeneous mixture.

• If it is not uniform, it is a heterogeneous mixture.

Chemical and Physical Changes

Definitions and Examples

Chemical and physical changes describe how matter transforms or rearranges.

• Chemical Change: Alters the composition of a substance, forming new substances. Example: Burning wood, rusting iron.

• Physical Change: Alters the form or appearance but not the composition. Example: Melting ice, dissolving sugar in water.

Temperature Scales

Celsius, Fahrenheit, Kelvin

Temperature can be measured in different units. Conversions are essential for calculations.

• Celsius (°C): Water freezes at 0°C and boils at 100°C.

• Fahrenheit (°F): Water freezes at 32°F and boils at 212°F.

• Kelvin (K): Absolute temperature scale; 0 K is absolute zero.

Key Equations:

• $K = °C + 273.15$

• $°F = (9/5) × °C + 32$

• $°C = (°F - 32) × 5/9$

Density and Volume Calculations

Density, Volume, and Mass Relationships

Density relates mass and volume and is a key property for identifying substances.

• Density Formula: $d = \frac{m}{V}$

• Volume Formula: $V = \frac{m}{d}$

• Mass Formula: $m = d \times V$

Units: Commonly g/mL or g/cm3 for liquids and solids; kg/L for gases.

Accuracy vs. Precision

Definitions and Importance

• Accuracy: How close a measurement is to the true value.

• Precision: How close repeated measurements are to each other.

High accuracy means measurements are correct; high precision means measurements are consistent.

Significant Figures

Rules and Calculations

• All nonzero digits are significant.

• Zeros between nonzero digits are significant.

• Leading zeros are not significant.

• Trailing zeros are significant only if there is a decimal point.

Multiplication/Division: Result has the same number of significant figures as the measurement with the fewest significant figures.

Addition/Subtraction: Result has the same number of decimal places as the measurement with the fewest decimal places.

Metric System and Unit Conversions

Common Prefixes and Conversion Factors

• 1 kilogram (kg) = 1000 grams (g)

• 1 meter (m) = 100 centimeters (cm)

• 1 liter (L) = 1000 milliliters (mL)

Use dimensional analysis to convert between units.

Volume Calculations

Geometric Formulas

• For a cube: $V = \text{side}^3$

• For a cylinder: $V = \pi r^2 h$

Dalton’s Atomic Theory

Postulates and Modern Modifications

• All matter is composed of atoms.

• Atoms of the same element are identical (modern science recognizes isotopes).

• Atoms combine in simple whole-number ratios to form compounds.

• Chemical reactions involve rearrangement of atoms; atoms are not created or destroyed.

Subatomic Particles

Protons, Neutrons, and Electrons

• Electron: Negative charge, negligible mass.

• Proton: Positive charge, mass ≈ 1 amu.

• Neutron: No charge, mass ≈ 1 amu.

Isotopes

Definition and Notation

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

• Notation: $^{A}_{Z}\text{X}$, where A = mass number, Z = atomic number, X = element symbol.

Example: $^{24}_{12}\text{Mg}$

Fractional Abundance of Isotopes

Calculating Average Atomic Mass

• Average atomic mass = Σ (fractional abundance × isotope mass)

Example: If an element has two isotopes, calculate the weighted average using their masses and relative abundances.

Mass Spectrometry

Determining Isotopic Composition

• Mass spectrometry separates isotopes based on mass-to-charge ratio.

• Used to determine atomic and molecular masses and isotopic abundances.

Periodic Table Organization

Groups, Periods, and Properties

• Groups (columns) contain elements with similar chemical properties.

• Periods (rows) show trends in properties across the table.

• Metals, nonmetals, and metalloids are classified based on properties.

Empirical and Molecular Formulas

Definitions and Calculations

• Empirical Formula: Simplest whole-number ratio of elements in a compound.

• Molecular Formula: Actual number of atoms of each element in a molecule.

To determine the empirical formula, divide the subscripts in the molecular formula by their greatest common factor.

Nomenclature

Naming Ionic and Molecular Compounds

• Ionic compounds: Name cation first, then anion (e.g., NaCl = sodium chloride).

• Molecular compounds: Use prefixes to indicate number of atoms (e.g., CO2 = carbon dioxide).

Functional Groups

Common Organic Functional Groups

• Carboxylic acid: –COOH group

• Amines: –NH2 group

Functional groups determine the chemical reactivity of organic molecules.

Additional info: Some explanations and examples were expanded for clarity and completeness based on standard General Chemistry curriculum.