General Chemistry Exam 1 Study Guide: Chapters 1 & 2

1. The Scientific Method

1.1 Qualitative and Quantitative Measurements

The scientific method is a systematic approach to investigation, involving observation, hypothesis formation, experimentation, and analysis. Measurements in chemistry can be qualitative (descriptive, non-numerical) or quantitative (numerical).

• Qualitative measurements: Describe qualities or characteristics (e.g., color, odor).

• Quantitative measurements: Involve numbers and units (e.g., mass, volume).

• Application: Scientists use both types to gather comprehensive data about chemical phenomena.

1.2–1.5 Measurements

1.2 Scientific Notation and Decimal Notation

Scientific notation expresses numbers as a product of a coefficient and a power of ten, making it easier to handle very large or small values.

• Example: $6,200 = 6.2 \times 10^3$

• Conversion: Move the decimal point to create a number between 1 and 10, then count the places moved for the exponent.

1.3–1.4 Units and Conversions

Understanding and converting between different units is essential in chemistry.

• SI Units: Standard units for scientific measurements (meter, kilogram, second, mole, etc.).

• Unit conversions: Use dimensional analysis to convert between units.

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

1.5 Accuracy, Precision, and Significant Figures

1.5.1 Definitions

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

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

• Significant figures: Digits in a measurement that are known with certainty plus one estimated digit.

1.5.2 Rules for Significant Figures

• 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.

• Example: 0.00450 has three significant figures.

1.10 Significant Figures in Calculations

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

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

• Rounding: Only round the final answer, not intermediate steps.

1.11 Converting Units

Dimensional analysis is used to convert from one unit to another using conversion factors.

• Example: To convert 10 inches to centimeters: $10 \text{ in} \times \frac{2.54 \text{ cm}}{1 \text{ in}} = 25.4 \text{ cm}$

2.1 Chemistry and the Elements

• Know the names, symbols, and atomic numbers of common elements.

• Be able to identify elements based on their symbols and vice versa.

2.2–2.3 Elements and the Periodic Table

• Classification: Elements are classified as metals, nonmetals, and metalloids.

• Groups and Periods: Vertical columns are groups (families); horizontal rows are periods.

• Common Names: Some elements have traditional names (e.g., sodium for Na).

2.4–2.5 Atomic Mass and Dalton’s Atomic Theory

• Dalton’s Atomic Theory: All matter is composed of atoms; atoms of the same element are identical; atoms combine in simple ratios to form compounds.

• Atomic mass: The weighted average mass of an element’s isotopes.

• Application: Dalton’s theory explains the law of conservation of mass and definite proportions.

2.6–2.7 Atomic Structure

• Subatomic particles: Protons (positive), neutrons (neutral), electrons (negative).

• Models: Key experiments by Millikan (oil drop), Thomson (cathode ray), and Rutherford (gold foil) established the structure of the atom.

• Nucleus: Contains protons and neutrons; electrons orbit the nucleus.

2.8 Atomic Number and the Mole

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

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

• Mole: $1 \text{ mole} = 6.022 \times 10^{23}$ particles (Avogadro’s number).

• Calculations: Use molar mass to convert between grams and moles.

2.11 Mixtures and Chemical Compounds: Molecules and Covalent Bonds

• Types of matter: Elements, compounds, and mixtures.

• Covalent bonds: Atoms share electrons to form molecules.

• Identification: Be able to distinguish between types of matter and bonding.

2.12 Ions and Ionic Bonds

• Ions: Atoms or molecules with a net charge due to loss or gain of electrons.

• Ionic bonds: Electrostatic attraction between oppositely charged ions (typically metals and nonmetals).

• Electron counting: Determine the number of electrons from atomic number and charge.

2.13 Naming Chemical Compounds

• Binary ionic compounds: Name the cation (metal) first, then the anion (nonmetal) with “-ide” ending.

• Transition metals: Use Roman numerals to indicate charge (e.g., Fe2+ is iron(II)).

• Covalent compounds: Use prefixes to indicate the number of each atom (e.g., CO2 is carbon dioxide).

• Polyatomic ions: Memorize common ions (e.g., sulfate SO42−).