Introductory Chemistry Study Guide: Measurements, Matter, Atomic Structure, and Electronic Structure

Chapter 2: Chemistry and Measurements

Units of Measurement

Understanding units is fundamental in chemistry for quantifying physical properties. The three main systems are SI (International System), metric, and English units.

• Length: SI unit is meter (m); metric uses centimeters (cm), millimeters (mm); English uses inches (in), feet (ft).

• Volume: SI unit is cubic meter (m3); metric uses liter (L), milliliter (mL); English uses gallons (gal), quarts (qt).

• Mass: SI unit is kilogram (kg); metric uses gram (g), milligram (mg); English uses pound (lb), ounce (oz).

• Temperature: SI unit is kelvin (K); metric uses Celsius (°C); English uses Fahrenheit (°F).

Example: 1 inch = 2.54 cm (exact conversion).

Unit Conversions

Converting between units requires understanding conversion factors and dimensional analysis.

• Conversion Factor: A ratio expressing how many of one unit equals another unit.

• Dimensional Analysis: Multiply by conversion factors to cancel units and obtain desired unit.

Example: To convert 5.0 cm to inches: $5.0\ \text{cm} \times \frac{1\ \text{in}}{2.54\ \text{cm}} = 1.97\ \text{in}$

Scientific Notation

Scientific notation expresses very large or small numbers in the form $a \times 10^n$.

• Standard Number: Regular decimal form.

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

Example: 0.00045 = $4.5 \times 10^{-4}$

Measured vs. Exact Numbers

Measured numbers are obtained by measurement and have uncertainty; exact numbers are counted or defined and have no uncertainty.

• Measured: 12.3 g (from a scale)

• Exact: 12 eggs (counted), 1 inch = 2.54 cm (defined)

Significant Figures

Significant figures reflect the precision of a measurement. Rules determine which digits are significant.

• 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.

Calculations: Round answers to the correct number of significant figures based on the operation.

Metric Prefixes

Metric prefixes indicate multiples or fractions of units.

• Kilo- (k): $10^3$

• Centi- (c): $10^{-2}$

• Milli- (m): $10^{-3}$

• Micro- (\mu): $10^{-6}$

Example: 1 km = 1000 m

Density Calculations

Density is a physical property defined as mass per unit volume.

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

• Can solve for mass or volume if density and one variable are known.

Example: If density = 2.70 g/mL and volume = 10.0 mL, mass = $2.70\ \text{g/mL} \times 10.0\ \text{mL} = 27.0\ \text{g}$

Chapter 3: Matter and Energy

Classification of Matter

Matter can be classified based on composition and uniformity.

• Pure Substance: Has a fixed composition; can be an element or compound.

• Mixture: Combination of two or more substances; can be homogeneous (uniform) or heterogeneous (non-uniform).

• Element: Simplest form of matter; cannot be broken down.

• Compound: Composed of two or more elements chemically combined.

Example: Salt water is a homogeneous mixture; sand and water is a heterogeneous mixture.

Physical vs. Chemical Changes and Properties

Physical changes do not alter the chemical composition; chemical changes do.

• Physical Change: Change in state or appearance (e.g., melting, boiling).

• Chemical Change: Formation of new substances (e.g., rusting, burning).

• Physical Property: Observable without changing composition (e.g., color, density).

• Chemical Property: Describes ability to undergo chemical change (e.g., flammability).

Temperature Scales

Three main temperature scales are used in chemistry.

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

• Fahrenheit (°F): Water freezes at 32°F, boils at 212°F.

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

Conversions:

• $\text{K} = \text{°C} + 273$

• $\text{°F} = \frac{9}{5}\text{°C} + 32$

• $\text{°C} = \frac{5}{9}(\text{°F} - 32)$

Types of Energy

Energy is the capacity to do work. Two main types are kinetic and potential energy.

• Kinetic Energy: Energy of motion.

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

Example: A ball at the top of a hill has potential energy; rolling down, it has kinetic energy.

Units of Energy and Conversions

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

• 1 cal = 4.184 J

• Food energy is often measured in kilocalories (kcal).

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

Food Caloric Values

Different macronutrients provide different caloric values.

• Protein: 4 kcal/g

• Fat: 9 kcal/g

• Carbohydrates: 4 kcal/g

Example: A food with 10 g protein, 5 g fat, 20 g carbs: $10 \times 4 + 5 \times 9 + 20 \times 4 = 40 + 45 + 80 = 165$ kcal

Chapter 4: Atoms and Elements

Names and Symbols of Elements

Each element has a unique name and symbol, often derived from Latin or English.

• Example: Hydrogen (H), Carbon (C), Sodium (Na), Iron (Fe)

Periodic Table Organization

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

• Groups: Vertical columns; examples include alkali metals (Group 1), alkaline earth metals (Group 2), halogens (Group 17), noble gases (Group 18).

• Periods: Horizontal rows.

Classification of Elements

Elements are classified as metals, nonmetals, or metalloids based on their properties.

• Metals: Shiny, conductive, malleable.

• Nonmetals: Dull, poor conductors, brittle.

• Metalloids: Properties intermediate between metals and nonmetals.

Dalton’s Atomic Theory

Dalton proposed four principles to explain the nature of atoms.

• All matter is composed of atoms.

• Atoms of the same element are identical; atoms of different elements are different.

• Atoms combine in simple whole-number ratios to form compounds.

• Atoms cannot be created or destroyed in chemical reactions.

Modern View: Some aspects have been modified (e.g., atoms can be divided in nuclear reactions).

Subatomic Particles

Atoms consist of protons, neutrons, and electrons.

Particle	Charge	Position	Relative Mass
Proton	+1	Nucleus	1
Neutron	0	Nucleus	1
Electron	-1	Outside nucleus	~0.0005

Atomic Number and Mass Number

Atomic number (Z) is the number of protons; mass number (A) is the sum of protons and neutrons.

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

• Finding Subatomic Particles: Protons = atomic number; neutrons = mass number - atomic number; electrons = protons (for neutral atom).

Example: Carbon-14: 6 protons, 8 neutrons, 6 electrons.

Atomic Mass vs. Mass Number

Mass number is a whole number; atomic mass is the weighted average of isotopes.

• Mass Number: Integer sum of protons and neutrons.

• Atomic Mass: Decimal value from periodic table.

Chapter 5: Electronic Structure of Atoms and Periodic Trends

Electromagnetic Spectrum

The electromagnetic spectrum includes all types of electromagnetic radiation, differing in wavelength and energy.

• Regions: Gamma rays, X-rays, ultraviolet, visible, infrared, microwaves, radio waves.

• Wavelength and Energy: Shorter wavelength = higher energy.

Example: Visible light has longer wavelength and lower energy than X-rays.

Energy Level Changes

Electrons absorb energy to move to higher levels; emit energy when moving to lower levels.

• Absorption: Electron moves up; energy is absorbed.

• Emission: Electron moves down; energy is released as light.

Energy Levels, Sublevels, and Orbitals

Electrons occupy energy levels, which contain sublevels and orbitals.

• Energy Level (n): Principal quantum number; n = 1, 2, 3, ...

• Sublevel: s, p, d, f; each has a specific shape and energy.

• Orbital: Region where electrons are likely found; each orbital holds up to 2 electrons.

Example: n = 2 has 2 sublevels (2s, 2p); 2s has 1 orbital, 2p has 3 orbitals.

Modern Atomic Model

Electrons fill orbitals in order of increasing energy, following the Aufbau principle.

• Orbital Names: s (sphere), p (dumbbell), d (clover), f (complex).

• Maximum Capacities: s = 2, p = 6, d = 10, f = 14 electrons.

• Filling Order: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p, etc.

Electron Configuration

Electron configuration shows the arrangement of electrons in an atom.

• Full Configuration: Lists all occupied sublevels.

• Abbreviated Configuration: Uses noble gas core to shorten notation.

Example: Sodium (Na): Full: $1s^2\ 2s^2\ 2p^6\ 3s^1$; Abbreviated: [Ne] $3s^1$

Valence Electrons

Valence electrons are the outermost electrons involved in chemical bonding.

• Main Group Elements: Number of valence electrons equals group number (for Groups 1-8).

Example: Oxygen (Group 6A) has 6 valence electrons.