CHE 1011 Exam 1 Review: Foundations of Chemistry

The Metric System (SI Units)

Prefixes and Unit Conversions

The metric system is a standardized system of measurement used in science. It is based on units of ten and uses prefixes to indicate multiples or fractions of base units.

• Common Prefixes: kilo- (k, 103), centi- (c, 10-2), milli- (m, 10-3), micro- (μ, 10-6), nano- (n, 10-9).

• Conversions: To convert between units, multiply or divide by powers of ten according to the prefixes.

• Example: 1 kilometer (km) = 1000 meters (m); 1 milligram (mg) = 0.001 grams (g).

Matter

Classification of Matter

Matter is anything that has mass and occupies space. It can be classified based on composition and physical state.

• Elements: Pure substances consisting of only one type of atom (e.g., O2).

• Compounds: Substances composed of two or more elements chemically combined (e.g., H2O).

• Mixtures: Physical combinations of substances.

  • Solutions (Homogeneous Mixtures): Uniform composition throughout (e.g., saltwater).

  • Heterogeneous Mixtures: Non-uniform composition (e.g., salad, sand in water).

Phases of Matter

• Solid: Definite shape and volume.

• Liquid: Definite volume, takes shape of container.

• Gas: No definite shape or volume; fills container.

Chemical vs. Physical Changes

• Physical Change: Does not alter the chemical composition (e.g., melting ice).

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

Energy

Heat Energy and Particle Motion

Energy is the capacity to do work or produce heat. Heat energy is related to the motion of particles in matter.

• Kinetic Energy: Energy of motion.

• Potential Energy: Stored energy due to position or composition.

• Heat Capacity (C): The amount of heat required to change the temperature of a substance by 1°C.

• Formula: $q = mC\Delta T$ where q = heat (J), m = mass (g), C = specific heat (J/g·°C), ΔT = temperature change (°C).

• Example: Calculating the heat required to raise the temperature of 100 g of water by 10°C using water's specific heat (4.18 J/g·°C).

Scientific Method

Key Components

The scientific method is a systematic approach to investigation.

• Observation: Gathering data.

• Hypothesis: Proposed explanation.

• Experiment: Testing the hypothesis.

• Analysis: Interpreting results.

• Conclusion: Accepting or rejecting the hypothesis.

Scientific Law vs. Theory

• Law: Describes a consistent natural phenomenon (e.g., Law of Conservation of Mass).

• Theory: Explains why phenomena occur (e.g., Atomic Theory).

Measurement

Standard and Scientific Notation

• Standard Notation: Regular decimal numbers (e.g., 0.00056).

• Scientific Notation: Expresses numbers as a product of a coefficient and a power of ten (e.g., 5.6 × 10-4).

Accuracy and Precision

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

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

Significant Figures

• Identifying Significant Digits: All nonzero digits are significant; zeros between nonzero digits are significant; leading zeros are not significant; trailing zeros are significant if there is a decimal point.

• Reporting Measurements: Use the correct number of significant digits based on the measuring instrument.

• Calculations: For multiplication/division, the result should have as many significant figures as the measurement with the fewest significant figures. For addition/subtraction, the result should have as many decimal places as the measurement with the fewest decimal places.

Dimensional Analysis

Unit Conversions (Factor-Label Method)

Dimensional analysis uses conversion factors to change units.

• Steps:

  1. Identify the starting and desired units.

  2. Set up conversion factors so units cancel appropriately.

  3. Multiply through to obtain the answer in the desired units.

• Example: Convert 5.0 km to meters: $5.0\ \text{km} \times \frac{1000\ \text{m}}{1\ \text{km}} = 5000\ \text{m}$

Density

Definition and Formula

• Density (D): The mass per unit volume of a substance.

• Formula: $D = \frac{M}{V}$ where M = mass, V = volume.

• Example: If a block has a mass of 50 g and a volume of 10 mL, its density is $D = \frac{50\ \text{g}}{10\ \text{mL}} = 5\ \text{g/mL}$.

Temperature

Celsius and Kelvin Scales

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

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

• Conversion: $T_{\text{K}} = T_{\text{°C}} + 273.15$

• Example: 25°C = 298.15 K

Atoms

Development of Atomic Theory

• Democritus: Proposed that matter is made of indivisible particles called atoms.

• John Dalton: Formulated the first modern atomic theory; atoms are indivisible and combine in fixed ratios.

• J.J. Thomson: Discovered the electron using the cathode ray tube experiment.

• Ernest Rutherford: Discovered the nucleus via the gold foil experiment.

• Niels Bohr: Proposed that electrons orbit the nucleus in quantized energy levels.

Law of Conservation of Matter

• Matter is neither created nor destroyed in a chemical reaction.

• Application: The total mass of reactants equals the total mass of products.

Counting Atoms in Compounds

• Example: Zinc phosphate, Zn3(PO4)2, contains 3 Zn, 2 × 1 = 2 P, and 2 × 4 = 8 O atoms per formula unit.

The Periodic Table

Organization and Classification

• Metals: Left and center of the table; good conductors, malleable, ductile.

• Nonmetals: Right side; poor conductors, brittle.

• Metalloids: Border metals and nonmetals; intermediate properties.

• Main Group Elements (Representative Elements): Groups 1, 2, and 13–18.

• Transition Metals: Groups 3–12.

• Alkali Metals: Group 1 (except H); very reactive.

• Alkaline Earth Metals: Group 2; reactive, but less so than alkali metals.

• Halogens: Group 17; very reactive nonmetals.

• Noble Gases: Group 18; inert gases.

Atomic Structure

Subatomic Particles

• Protons: Positive charge (+1), mass ≈ 1 amu, located in nucleus.

• Neutrons: No charge, mass ≈ 1 amu, located in nucleus.

• Electrons: Negative charge (–1), mass ≈ 1/1836 amu, located outside nucleus.

Atomic Number and Mass Number

• Atomic Number (Z): Number of protons; defines the element.

• Mass Number (A): Number of protons + neutrons.

• Difference: Atomic number is always whole; mass number varies due to isotopes.

Calculating Atomic Mass

• Atomic mass is the weighted average of all naturally occurring isotopes.

• Formula: $\text{Atomic mass} = \sum (\text{fractional abundance} \times \text{isotope mass})$

Bohr Model vs. Quantum Model

• Bohr Model: Electrons orbit nucleus in fixed paths (energy levels).

• Quantum Model: Electrons exist in probability clouds (orbitals) around the nucleus.

Charge of Atoms and Ions

• Neutral Atom: Number of protons = number of electrons.

• Ion: Atom with unequal protons and electrons; cations (+), anions (–).

• Determining Charge: For main group elements, charge = group number (for metals) or group number – 8 (for nonmetals).

Isotopes

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

• Example: Carbon-12 and Carbon-14.

Chemical Nomenclature

Ionic Compounds

• Binary Ionic Compounds: Composed of two elements (metal + nonmetal). Name: cation first, then anion with -ide ending (e.g., NaCl = sodium chloride).

• Transition Metals: May have multiple charges; indicate charge with Roman numerals (e.g., FeCl2 = iron(II) chloride).

• Polyatomic Ions: Ions composed of multiple atoms (e.g., NO3– = nitrate).

Molecular Compounds

• Composed of nonmetals; use prefixes to indicate number of atoms (e.g., CO2 = carbon dioxide).

Acids

• Binary Acids: H + nonmetal; named as hydro- + base name + -ic acid (e.g., HCl = hydrochloric acid).

• Oxyacids: H + polyatomic ion; if ion ends in -ate, acid ends in -ic; if -ite, ends in -ous (e.g., H2SO4 = sulfuric acid).