General Chemistry Exam 1 Review: Matter, Measurement, Atoms, Molecules, and Ions

Chapter 1: Introduction – Matter and Measurement

Basic Definitions

This section introduces the foundational vocabulary of chemistry, essential for understanding the nature and classification of substances.

• Matter: Anything that has mass and occupies space.

• Substance: A form of matter with a definite composition and distinct properties (e.g., water, sodium chloride).

• Mixture: A combination of two or more substances in which each retains its own identity.

• Physical Property: A characteristic that can be observed without changing the substance’s identity (e.g., color, melting point).

• Chemical Property: Describes how a substance reacts with other substances (e.g., flammability, acidity).

• Intensive Property: Independent of the amount of substance (e.g., density, boiling point).

• Extensive Property: Dependent on the amount of substance (e.g., mass, volume).

• Physical Change: A change that does not alter the chemical composition (e.g., melting, freezing).

• Chemical Change: A process that alters the chemical composition (e.g., rusting, combustion).

• Element: A substance that cannot be broken down into simpler substances by chemical means.

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

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

• Molecule: Two or more atoms bonded together.

• Ion: An atom or molecule with a net electric charge due to the loss or gain of electrons.

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

Elements and the Periodic Table

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

• Know the names and symbols of the first twenty elements and commonly encountered elements.

• Diatomic Elements: H2, N2, O2, F2, Cl2, Br2, I2.

• Group Names:

  • Group 1A: Alkali Metals

  • Group 2A: Alkaline Earth Metals

  • Group 7A: Halogens

  • Group 8A: Noble Gases

• Transition metals are located in the center block (Groups 3-12).

• Semimetals (metalloids) border the stair-step line between metals and nonmetals.

Units of Measure and Conversions

Chemistry uses the International System of Units (SI) for consistency in measurements.

• SI Base Units: meter (m), kilogram (kg), second (s), kelvin (K).

• SI Prefixes: From $10^{15}$ (peta-) to $10^{-15}$ (femto-).

• Conversions: Be able to convert between English and metric units, and between compound units (e.g., miles/hour to feet/second).

Scientific Notation

Scientific notation expresses numbers as $N \times 10^n$, where $1 \leq N < 10$ and $n$ is an integer.

• Used for very large or very small numbers to simplify calculations and representation.

Significant Figures and Calculations

Significant figures reflect the precision of a measured quantity.

• Rules for Zeroes:

  • Leading zeroes: Never significant (e.g., 0.0025 has two significant figures).

  • Captive zeroes (between nonzero digits): Always significant (e.g., 205 has three significant figures).

  • Trailing zeroes: Significant only if there is a decimal point (e.g., 2.00 has three significant figures).

• Multiplying and Dividing: The result has as many significant figures as the measurement with the fewest significant figures.

• Adding and Subtracting: The result has as many decimal places as the measurement with the fewest decimal places.

• Accuracy vs. Precision:

  • Accuracy: Agreement with the true value.

  • Precision: Agreement among repeated measurements.

  • Systematic error affects accuracy; random error affects both accuracy and precision.

Solving Problems by Dimensional Analysis

Dimensional analysis uses units to guide the setup and solution of problems.

• Include all units in calculations and cancel appropriately.

• Think through each step to ensure logical consistency.

Temperature

Temperature can be measured in Celsius, Fahrenheit, or Kelvin. Know how to convert between these scales.

• Conversion Formulas:

  • $T_{\mathrm{K}} = T_{\mathrm{^{\circ}C}} + 273.15$

  • $T_{\mathrm{^{\circ}C}} = \frac{5}{9}(T_{\mathrm{^{\circ}F}} - 32)$

  • $T_{\mathrm{^{\circ}F}} = \frac{9}{5}T_{\mathrm{^{\circ}C}} + 32$

Density

Density is a measure of how much mass is contained in a given volume.

• Formula: $\mathrm{Density} = \frac{\mathrm{Mass}}{\mathrm{Volume}}$

• Common units: g/mL (solids and liquids), g/L (gases).

Chapter 2: Atoms, Molecules, and Ions

Fundamental Laws of Chemistry

These laws describe the basic principles governing chemical reactions and the composition of matter.

• Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction.

• Law of Constant Composition: A given compound always contains the same proportion of elements by mass.

• Law of Multiple Proportions: When two elements form more than one compound, the ratios of the masses of the second element that combine with a fixed mass of the first element are small whole numbers.

Experiments Relative to Atomic Structure and Atomic Particles

Key experiments led to the discovery of atomic structure and subatomic particles.

• J.J. Thomson’s Cathode Ray Tube Experiment: Discovered the electron and its negative charge.

• Millikan Oil Drop Experiment: Measured the charge of the electron.

• Rutherford’s Gold Foil Experiment: Discovered the atomic nucleus and proposed that atoms are mostly empty space.

Atomic Structure

Atoms are composed of protons, neutrons, and electrons.

• Proton: Positively charged particle in the nucleus; mass ≈ 1 amu.

• Neutron: Neutral particle in the nucleus; mass ≈ 1 amu.

• Electron: Negatively charged particle outside the nucleus; mass ≈ 1/1836 amu.

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

• Notation: $^{A}_{Z}\mathrm{X}^{C}$, where:

  • $A$ = mass number (protons + neutrons)

  • $Z$ = atomic number (number of protons)

  • $C$ = charge (if ion)

• Number of neutrons: $A - Z$

• All atoms of an element have the same number of protons; isotopes differ in neutrons; ions differ in electrons.

Calculating Average Atomic Mass

• Formula: $\text{Atomic Weight} = (\text{mass of isotope 1}) \times (\% \text{abundance of isotope 1}) + (\text{mass of isotope 2}) \times (\% \text{abundance of isotope 2}) + \ldots$

The Periodic Table

The periodic table arranges elements by atomic number and groups them by similar properties.

• Groups (vertical columns): Elements with similar chemical properties.

• Periods (horizontal rows): Elements with the same number of electron shells.

• Categories: Metals, nonmetals, metalloids (semimetals).

• Special Group Names:

  • Group 1A: Alkali Metals

  • Group 2A: Alkaline Earth Metals

  • Group 6A: Chalcogens

  • Group 7A: Halogens

  • Group 8A: Noble Gases

• Metalloids border the stair-step line (except aluminum, which is a metal).

Molecules and Molecular Compounds

Chemical compounds can be classified by the types of bonds and particles they contain.

• Covalent Compounds: Composed of molecules formed by sharing electrons between nonmetals.

• Ionic Compounds: Composed of ions formed by the transfer of electrons between metals and nonmetals.

• Molecular Formula: Shows the actual number and types of atoms in a molecule (e.g., H2O).

• Empirical Formula: Shows the simplest whole-number ratio of atoms (e.g., CH6 for C2H12).

Ions and Ionic Compounds

Ionic compounds are formed from the electrostatic attraction between cations and anions.

• Cations: Positively charged ions (fewer electrons than protons).

• Anions: Negatively charged ions (more electrons than protons).

• Formed typically between metals (Groups 1 or 2) and nonmetals (Groups 6A or 7A).

• Some ions are polyatomic (contain more than one atom).

• Memorize common polyatomic ions (see textbook Section 2.8).

Naming Inorganic Compounds

Systematic rules are used to name chemical compounds based on their composition.

• Binary Ionic Compounds (Metal forms one cation):

  • Metal (cation) named first, nonmetal (anion) named second with -ide suffix.

  • Example: NaCl – sodium chloride; CaCl2 – calcium chloride.

• Binary Ionic Compounds (Metal forms multiple cations):

  • Charge of cation specified by Roman numeral.

  • Example: FeCl2 – iron(II) chloride; FeCl3 – iron(III) chloride.

• Ionic Compounds with Polyatomic Ions:

  • Name the cation and the polyatomic anion.

  • Example: NaNO3 – sodium nitrate.

• Acids:

  • Without oxygen: Prefix hydro-, suffix -ic (e.g., HCl – hydrochloric acid).

  • With oxygen (oxyacids): Named based on the number of oxygens (see textbook for details).

• Binary Molecular Compounds:

  • Use prefixes to indicate the number of atoms (mono-, di-, tri-, etc.).

  • Mono- is not used for the first element.

  • Example: CO2 – carbon dioxide; PCl5 – phosphorus pentachloride.

Simple Organic Compounds

Learn the names of the alkanes and their derivatives (see textbook pages 69-70).

• Alkanes: Saturated hydrocarbons with the general formula CnH2n+2.

• Examples: Methane (CH4), Ethane (C2H6), Propane (C3H8), etc.

Additional info: For a comprehensive understanding, refer to the textbook sections and tables mentioned for detailed lists of elements, ions, and naming conventions. Practice with problems involving unit conversions, significant figures, and naming compounds to reinforce these concepts.