Study Guide: Chemistry and Measurements, Matter and Energy, Atoms and Elements (Chapters 2–4)

Chapter 2: Chemistry and Measurements

Metric Units and Prefixes

The metric system is a standardized system of measurement used in science. It is based on units such as the meter (length), gram (mass), and liter (volume), with prefixes to indicate multiples or fractions.

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

• Example: 1 kilometer (km) = 1,000 meters (m).

Measured Numbers, Significant Figures, and Exact Numbers

Measured numbers are obtained by measurement and have a degree of uncertainty. Significant figures (sig figs) reflect the precision of a measurement. Exact numbers are counted or defined and have infinite significant figures.

• 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.

• Rounding Off: If the digit to be dropped is less than 5, leave the preceding digit unchanged; if 5 or greater, increase the preceding digit by 1.

• Adding Zeroes: Used to indicate the correct number of significant figures.

• Exact Numbers: Numbers from counting objects or defined relationships (e.g., 1 dozen = 12).

Density and Specific Gravity

Density is the mass of a substance per unit volume. Specific gravity is the ratio of the density of a substance to the density of water.

• Density Formula:

$\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Specific Gravity Formula:

$\text{Specific Gravity} = \frac{\text{Density of substance}}{\text{Density of water}}$

• Example: If a liquid has a density of 1.2 g/mL, its specific gravity is 1.2 (since water's density is 1.0 g/mL).

Conversion Factors and Equalities

Conversion factors are ratios derived from equalities that relate different units. They are used to convert from one unit to another.

• Example: 1 inch = 2.54 cm, so $\frac{1\ \text{in}}{2.54\ \text{cm}}$ and $\frac{2.54\ \text{cm}}{1\ \text{in}}$ are both valid conversion factors.

Problem Solving in Chemistry

Problem solving often involves unit conversions using dimensional analysis (factor-label method).

• Steps:

  1. Identify the given and required units.

  2. Write the conversion factors.

  3. Set up the calculation so units cancel appropriately.

• Example: Convert 5.0 inches to centimeters: $5.0\ \text{in} \times \frac{2.54\ \text{cm}}{1\ \text{in}} = 13\ \text{cm}$

Chapter 3: Matter and Energy

Pure Substances and Mixtures

Matter can be classified as pure substances or mixtures.

• Pure Substance: Has a fixed composition (elements or compounds).

• Mixture: Physical blend of two or more substances; can be homogeneous (uniform) or heterogeneous (not uniform).

• Example: Salt water is a homogeneous mixture; sand and iron filings are a heterogeneous mixture.

States of Matter

Matter exists in three primary states: solid, liquid, and gas.

• Solid: Definite shape and volume.

• Liquid: Definite volume, takes shape of container.

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

Physical and Chemical Changes

Physical changes do not alter the chemical composition (e.g., melting, boiling). Chemical changes result in new substances (e.g., burning, rusting).

• Example: Ice melting is a physical change; iron rusting is a chemical change.

Heat and Phase Changes

Phase changes involve energy transfer as heat.

• Heat of Fusion: Energy required to melt 1 g of a substance at its melting point.

• Heat of Vaporization: Energy required to vaporize 1 g of a substance at its boiling point.

• Melting/Freezing: Solid to liquid/liquid to solid.

• Boiling/Condensation: Liquid to gas/gas to liquid.

• Sublimation/Deposition: Solid to gas/gas to solid.

• Steam Burns: Steam contains more energy than boiling water, causing more severe burns.

Energy Units and Conversions

Energy is measured in calories (cal) or joules (J).

• Conversion: $1\ \text{cal} = 4.184\ \text{J}$

• Example: 100 cal = $100 \times 4.184 = 418.4$ J

Specific Heat and Heat Calculations

Specific heat is the amount of heat required to raise the temperature of 1 g of a substance by 1°C.

• Specific Heat Formula:

$q = m \times c \times \Delta T$

• Where q = heat (J), m = mass (g), c = specific heat (J/g°C), ΔT = change in temperature (°C).

• Example: How much heat is needed to raise 50 g of water from 20°C to 30°C? $q = 50 \times 4.18 \times (30-20) = 2,090$ J

Heating and Cooling Curves

Heating and cooling curves show temperature changes as a substance is heated or cooled, including plateaus where phase changes occur.

• Plateau: Temperature remains constant during phase change (e.g., melting, boiling).

• Interpretation: The length of the plateau indicates the amount of energy required for the phase change.

Chapter 4: Atoms and Elements

Periodic Table: Groups and Periods

The periodic table organizes elements by increasing atomic number. Groups are vertical columns; periods are horizontal rows.

• Groups: Elements in the same group have similar chemical properties.

• Periods: Elements in the same period have the same number of electron shells.

Metals, Nonmetals, and Metalloids

Elements are classified based on their properties.

• Metals: Shiny, good conductors, malleable (left side of table).

• Nonmetals: Dull, poor conductors, brittle (right side of table).

• Metalloids: Properties intermediate between metals and nonmetals (along the staircase line).

Atomic Number and Atomic Mass

Atomic number is the number of protons in an atom. Atomic mass is the weighted average mass of all isotopes of an element.

• Calculation: Atomic number = number of protons = number of electrons (in a neutral atom).

• Atomic mass: Sum of protons and neutrons.

Valence Electrons

Valence electrons are the outermost electrons involved in chemical bonding.

• Group Number: For main group elements, the group number indicates the number of valence electrons.

• Example: Carbon (Group 4A) has 4 valence electrons.

Orbital Diagrams and Electron Configuration

Electrons occupy orbitals in a specific order. Orbital diagrams use arrows to represent electrons in boxes (orbitals). Electron configuration shows the arrangement of electrons by energy level and sublevel.

• Example: For oxygen (atomic number 8): $1s^2\ 2s^2\ 2p^4$

Periodic Trends

Properties of elements change predictably across periods and groups.

• Atomic Size: Increases down a group, decreases across a period.

• Ionization Energy: Energy required to remove an electron; decreases down a group, increases across a period.

• Metallic Character: Increases down a group, decreases across a period.

Lewis Symbols

Lewis symbols represent valence electrons as dots around the element symbol.

• Example: Sodium (Na) has one dot; oxygen (O) has six dots.

Element	Lewis Symbol	Valence Electrons
Na	Na·	1
O	O······	6
Cl	Cl·······	7

Additional info: Some explanations and examples have been expanded for clarity and completeness.