Unit Exam 1 Study Guide: Introductory Chemistry Topics

Chapter 1 – Chemistry in Our Lives

Scientific Method

The scientific method is a systematic approach used by scientists to investigate phenomena, acquire new knowledge, or correct and integrate previous knowledge. It involves several steps:

• Observation: Gathering information about the world.

• Hypothesis: Formulating a testable explanation.

• Experiment: Testing the hypothesis through controlled procedures.

• Analysis: Interpreting data and drawing conclusions.

• Conclusion: Accepting, rejecting, or modifying the hypothesis based on results.

Percent Calculations

Percent calculations are used to express quantities as parts of a whole. The formula is:

• $\text{Percent} = \frac{\text{Part}}{\text{Whole}} \times 100\%$

• Example: If 20 out of 100 students passed, $\frac{20}{100} \times 100\% = 20\%$ passed.

Chapter 2 – Chemistry and Measurements

Metric Conversions

The metric system uses prefixes to indicate multiples of base units. Common prefixes include:

Prefix	Symbol	Factor
Kilo	k	1,000
Centi	c	0.01
Milli	m	0.001
Micro	μ	0.000001
Nano	n	0.000000001
Additional info: Table 2.6 may include more prefixes such as deci (d, 0.1), mega (M, 1,000,000).

Making Measurements

Measurements in chemistry require precision and accuracy. Always record all certain digits and one uncertain digit.

Significant Figures

Significant figures indicate the precision of a measurement. Rules include:

• All nonzero digits are significant.

• Zeros between nonzero digits are significant.

• Leading zeros are not significant.

• Trailing zeros in a decimal number are significant.

Dimensional Analysis

Dimensional analysis is a method for converting units using conversion factors.

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

Density

Density is the mass per unit volume of a substance.

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

• Example: If a block has a mass of 10 g and a volume of 2 cm3, its density is $\frac{10\ \text{g}}{2\ \text{cm}^3} = 5\ \text{g/cm}^3$

Chapter 3 – Energy and Matter

Classification of Matter

Matter can be classified as:

• Pure substances: Elements (e.g., oxygen, gold) and compounds (e.g., water, sodium chloride).

• Mixtures: Homogeneous (uniform composition, e.g., salt water) and heterogeneous (non-uniform, e.g., sand and water).

Temperature Conversions

Temperature can be converted between Celsius, Fahrenheit, and Kelvin using formulas:

• $\text{Celsius to Fahrenheit}: F = C \times \frac{9}{5} + 32$

• $\text{Celsius to Kelvin}: K = C + 273.15$

Potential and Kinetic Energy

Potential energy is stored energy due to position; kinetic energy is energy of motion.

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

Specific Heat

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

• Formula: $q = m \times c \times \Delta T$

• Where $q$ = heat energy, $m$ = mass, $c$ = specific heat, $\Delta T$ = change in temperature.

Energy Calculations

To calculate energy gained or lost:

• Use the specific heat equation above.

• Example: 10 g of water ($c = 4.18\ \text{J/g}^\circ\text{C}$) heated from 20°C to 30°C: $q = 10 \times 4.18 \times (30-20) = 418\ \text{J}$

Chapter 10 – Section 10.7 Change of State

Phase Transitions

Phase transitions are changes between solid, liquid, and gas states. Types include:

• Melting: Solid to liquid

• Freezing: Liquid to solid

• Vaporization: Liquid to gas

• Condensation: Gas to liquid

• Sublimation: Solid to gas

• Deposition: Gas to solid

Energy in Phase Transitions

Energy is absorbed or released during phase changes. Calculations use:

• $q = m \times \Delta H$ (where $\Delta H$ is enthalpy of fusion or vaporization)

• Combined heat calculations may involve both temperature change and phase change.

Heating and Cooling Curves

A heating/cooling curve shows temperature changes and phase transitions as a substance is heated or cooled.

• Plateaus represent phase changes (e.g., melting, boiling).

• Sloped regions represent temperature changes within a single phase.

Chapter 4 – Atoms and Elements

Periodic Table Organization

The periodic table is organized by increasing atomic number. Elements are grouped by similar properties:

• Metals: Left and center; shiny, conductive.

• Non-metals: Right; dull, poor conductors.

• Metalloids: Border metals/non-metals; intermediate properties.

• First 20 elements: Know their symbols and names (e.g., H - Hydrogen, He - Helium, Li - Lithium, etc.).

Structure of an Atom

An atom consists of:

• Protons: Positive charge, in nucleus.

• Neutrons: Neutral, in nucleus.

• Electrons: Negative charge, orbit nucleus.

Isotopes

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

• Atomic number (Z): Number of protons.

• Mass number (A): Protons + neutrons.

• Calculations: Number of neutrons = mass number - atomic number.

• Electrons: Equal to protons in a neutral atom.

Average Atomic Mass

Average atomic mass is calculated using the masses and abundances of isotopes:

• $\text{Average atomic mass} = \sum (\text{isotope mass} \times \text{fractional abundance})$

• Example: If isotope A is 50% at 10 amu, isotope B is 50% at 12 amu: $10 \times 0.5 + 12 \times 0.5 = 11$ amu

Chapter 5 – Electronic Structure of Atoms and Periodic Trends

Electronic Structure

Electrons occupy energy levels and subshells (orbitals) around the nucleus.

• Energy levels: Principal shells (n=1,2,3...)

• Subshells: s, p, d, f orbitals

Electron Configurations

Electron configuration shows the arrangement of electrons:

• Long form: Lists all orbitals (e.g., 1s2 2s2 2p6).

• Shorthand: Uses noble gas core (e.g., [Ne] 3s2).

Periodic Trends

Periodic trends describe how properties change across the periodic table:

Trend	Definition	Direction
Atomic Size	Size of atom	Increases down, decreases across
Ionization Energy	Energy to remove electron	Decreases down, increases across
Metallic Character	Metal-like properties	Increases down, decreases across

Chapter 16 – Nuclear Chemistry

Types of Radiation

Common types of radiation include:

Type	Symbol	Mass	Charge
Alpha	$\alpha$ or $^4_2\text{He}$	4	+2
Beta	$\beta$ or $^0_{-1}\text{e}$	0	-1
Gamma	$\gamma$	0	0

Radioactive Decay Equations

Radioactive decay is represented by nuclear equations:

• Example: $^\text{A}_\text{Z}\text{X} \rightarrow ^{A-4}_{Z-2}\text{Y} + ^4_2\text{He}$ (alpha decay)

• Balance mass and atomic numbers on both sides.

Additional info: Nuclear chemistry also includes concepts such as half-life and applications in medicine and energy.