Exam 1 Study Guide: Chemical Tools, Atoms, Molecules & Ions

1. The Scientific Method

1.1 Qualitative and Quantitative Measurements

The scientific method is a systematic approach to experimentation and measurement in chemistry. It distinguishes between qualitative (descriptive) and quantitative (numerical) observations.

• Qualitative measurements: Observations that describe properties or occurrences without using numbers (e.g., color, odor).

• Quantitative measurements: Observations that involve numbers and units (e.g., mass, volume, temperature).

• Application: Scientists use both types to formulate hypotheses and design experiments.

1.2–1.5 Measurements

1.2 Scientific Notation

• Scientific notation expresses very large or small numbers as a product of a coefficient and a power of ten.

• Example: $6,200 = 6.2 \times 10^3$

1.3–1.4 Units and SI System

• The SI (International System of Units) is the standard for scientific measurements.

• Common SI units: meter (m), kilogram (kg), second (s), mole (mol), kelvin (K), ampere (A), candela (cd).

1.5 Metric Prefixes

• Metric prefixes indicate multiples or fractions of base units (e.g., kilo-, centi-, milli-).

• Example: 1 kilometer (km) = $1 \times 10^3$ meters (m)

1.6–1.8 Significant Figures and Accuracy

1.6–1.7 Significant Figures

• Significant figures reflect the precision of a measured value.

• Rules determine which digits are significant (all nonzero digits, zeros between nonzero digits, trailing zeros after a decimal point).

• Example: 0.00450 has three significant figures.

1.8 Accuracy and Precision

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

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

• Use significant figures to report the correct precision in calculations.

1.10–1.11 Calculations and Unit Conversions

1.10 Significant Figures in Calculations

• When multiplying/dividing, the result should have as many significant figures as the measurement with the fewest significant figures.

• When adding/subtracting, the result should have as many decimal places as the measurement with the fewest decimal places.

1.11 Converting Units

• Dimensional analysis (factor-label method) is used to convert between units.

• Example: To convert 5.0 cm to meters: $5.0 \text{ cm} \times \frac{1 \text{ m}}{100 \text{ cm}} = 0.050 \text{ m}$

2.1–2.3 Chemistry and the Elements

2.1 Elements

• Elements are pure substances consisting of only one type of atom.

• Each element has a unique atomic number and symbol (e.g., H for hydrogen, O for oxygen).

2.2–2.3 The Periodic Table

• The periodic table organizes elements by increasing atomic number and recurring chemical properties.

• Groups (columns) and periods (rows) classify elements as metals, nonmetals, and metalloids.

• Common names and symbols should be memorized for frequently encountered elements.

2.4–2.5 Atomic Theory

2.4 Mass Laws and Dalton’s Atomic Theory

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

• Law of Definite Proportions: A compound always contains the same elements in the same proportion by mass.

• Dalton’s Atomic Theory: Atoms are indivisible particles that make up elements and combine in fixed ratios to form compounds.

2.6–2.7 Atomic Structure

2.6 Subatomic Particles

• Atoms are composed of protons (positive), neutrons (neutral), and electrons (negative).

• Key experiments (Millikan, Thomson, Rutherford) revealed the structure of the atom.

2.7 Atomic Number and Mass Number

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

• Mass number (A): Total number of protons and neutrons.

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

2.9–2.10 Atomic Weight and the Mole

2.9 Atomic Weight

• Atomic weight is the weighted average mass of an element’s isotopes.

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

2.10 The Mole and Avogadro’s Number

• The mole is the SI unit for amount of substance; 1 mole = $6.022 \times 10^{23}$ particles (Avogadro’s number).

• Relates mass, number of particles, and molar mass.

• Example: $\text{Moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

2.11–2.13 Compounds and Bonding

2.11 Mixtures and Chemical Compounds

• Mixtures are physical combinations of substances; compounds are chemical combinations of elements in fixed ratios.

• Chemical bonds include ionic (transfer of electrons) and covalent (sharing of electrons).

2.12 Ions and Ionic Bonds

• Ions are charged particles formed when atoms gain or lose electrons.

• Cations: Positively charged (loss of electrons); Anions: Negatively charged (gain of electrons).

• Ionic bonds form between metals and nonmetals due to electrostatic attraction.

2.13 Naming Chemical Compounds

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

• Use Roman numerals for metals with variable charges (e.g., Fe2+ is iron(II)).

• Binary covalent compounds: Use prefixes (mono-, di-, tri-) to indicate the number of each atom.

• Practice with common polyatomic ions and their names.