CHM1025 Study Guide: Foundations of General Chemistry (Corwin 8th Ed, Exam #1)

Chapter Prerequisite Science Skills

Measurements

Accurate measurement is fundamental in chemistry, enabling quantitative analysis and comparison. Measurements involve determining the amount, length, mass, or volume of substances using standardized units.

• Key Point: Always record measurements with the correct unit (e.g., meters, grams, liters).

• Example: Measuring the mass of a sample as 12.5 g.

Significant Digits (Figures)

Significant digits indicate the precision of a measurement. They include all certain digits plus one estimated digit.

• Key Point: All nonzero digits are significant; zeros may or may not be significant depending on their position.

• Example: 0.00450 has three significant digits.

Rounding

Rounding is used to express numbers to the appropriate number of significant digits based on measurement precision.

• Key Point: If the digit to be dropped is less than 5, leave the preceding digit unchanged; if 5 or greater, increase the preceding digit by one.

• Example: 2.746 rounded to two significant digits is 2.7.

Significant Figures in Calculations

When performing calculations, the number of significant figures in the result depends on the operation:

• Multiplication/Division: The result should have the same number of significant figures as the measurement with the fewest significant figures.

• Addition/Subtraction: The result should have the same number of decimal places as the measurement with the fewest decimal places.

• Example: $4.56 \times 1.4 = 6.4$ (two significant figures).

Scientific Notation

Scientific notation expresses very large or small numbers in the form $a \times 10^n$, where $a$ is a number between 1 and 10 and $n$ is an integer.

• Key Point: Useful for representing measurements in chemistry, such as Avogadro's number ($6.02 \times 10^{23}$).

• Example: 0.00056 = $5.6 \times 10^{-4}$.

Proper Use of the Calculator

Calculators are essential for performing chemical calculations. Proper use includes entering numbers in scientific notation and rounding results appropriately.

• Key Point: Always check calculator settings for scientific notation and significant figures.

Chapter 2. The Metric System

Basic Metric Units

The metric system is the standard system of measurement in science. The basic units are:

• Length: meter (m)

• Mass: gram (g)

• Volume: liter (L)

• Temperature: Celsius (°C)

Prefixes

Metric prefixes indicate multiples or fractions of basic units.

• Key Prefixes: kilo- ($10^3$), centi- ($10^{-2}$), milli- ($10^{-3}$), micro- ($10^{-6}$)

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

Conversions: Metric–Metric, Metric–English

Conversions are necessary to switch between units within the metric system or between metric and English units.

• Key Point: Use conversion factors to relate units.

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

Percent

Percent expresses a ratio as parts per hundred.

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

• Example: If 25 g of salt is dissolved in 100 g of water, percent salt = $\frac{25}{125} \times 100 = 20\%$

Volume and Density Determinations

Volume is the space occupied by a substance; density is mass per unit volume.

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

• Example: If a sample has mass 10 g and volume 2 mL, density = $\frac{10}{2} = 5 \text{ g/mL}$

Temperature Conversions

Temperature can be converted between Celsius, Fahrenheit, and Kelvin.

• Formulas:

  • $\text{K} = \text{°C} + 273.15$

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

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

• Example: 25°C = $25 + 273.15 = 298.15$ K

Chapter 3. Matter and Energy

States of Matter and Transitions

Matter exists in three primary states: solid, liquid, and gas. Transitions include melting, freezing, vaporization, condensation, and sublimation.

• Key Point: Each state has distinct properties (shape, volume, compressibility).

• Example: Water transitions from solid (ice) to liquid (water) at 0°C.

Matter: Elements and Compounds

Matter is classified as elements (pure substances of one type of atom) or compounds (substances composed of two or more elements chemically combined).

• Key Point: Elements cannot be broken down by chemical means; compounds can.

• Example: Water (H2O) is a compound; oxygen (O2) is an element.

Mixtures: Homogeneous vs Heterogeneous

Mixtures are combinations of substances not chemically bonded. Homogeneous mixtures are uniform throughout; heterogeneous mixtures are not.

• Example: Saltwater is homogeneous; sand and water is heterogeneous.

Names and Symbols of the Elements

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

• Example: Sodium (Na), Potassium (K), Iron (Fe)

Periodic Table

The periodic table organizes elements by increasing atomic number and similar properties.

• Key Point: Groups (columns) contain elements with similar chemical properties.

• Example: Group 1: Alkali metals; Group 17: Halogens

Classification of Elements

Elements are classified as metals, non-metals, and semi-metals (metalloids).

Type	Properties	Examples
Metals	Conduct electricity, malleable, shiny	Iron, Copper
Non-metals	Poor conductors, brittle, dull	Oxygen, Sulfur
Metalloids	Intermediate properties	Silicon, Boron

Physical vs Chemical Changes

Physical changes alter the form but not the composition of a substance; chemical changes produce new substances.

• Example: Melting ice (physical); burning wood (chemical)

Law of Conservation of Mass and Energy

The law of conservation of mass states that mass is neither created nor destroyed in a chemical reaction. The law of conservation of energy states that energy cannot be created or destroyed, only transformed.

• Example: In combustion, the mass of reactants equals the mass of products.

Kinetic vs Potential Energy

Kinetic energy is energy of motion; potential energy is stored energy due to position or composition.

• Example: A moving car has kinetic energy; a stretched spring has potential energy.

Chapter 4. Models of the Atom

Dalton’s Theory – Four Postulates

Dalton’s atomic theory laid the foundation for modern chemistry:

1. All matter is composed of atoms.

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

3. Atoms cannot be created, divided, or destroyed.

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

Model of the Atom (Bohr): Protons, Neutrons, and Electrons

The Bohr model describes the atom as a nucleus (protons and neutrons) surrounded by electrons in defined orbits.

• Proton: Positive charge, mass ≈ 1 amu

• Neutron: Neutral, mass ≈ 1 amu

• Electron: Negative charge, mass ≈ 1/1836 amu

Atomic Notation and Isotopes

Atomic notation represents the number of protons, neutrons, and electrons. Isotopes are atoms of the same element with different numbers of neutrons.

• Notation: $\text{Element}^{\text{Mass Number}}_{\text{Atomic Number}}$

• Example: $^{14}_{6}\text{C}$ (Carbon-14)

Determining Number of Protons, Neutrons, Electrons

• Protons: Equal to atomic number

• Neutrons: Mass number minus atomic number

• Electrons: Equal to protons in a neutral atom

• Example: $^{23}_{11}\text{Na}$: 11 protons, 12 neutrons, 11 electrons

Atomic Mass – Determining Average Mass

Average atomic mass is calculated based on the abundance and mass of each isotope.

• Formula: $\text{Average Atomic Mass} = \sum (\text{Isotope Mass} \times \text{Fractional Abundance})$

• Example: If $^{35}\text{Cl}$ is 75% and $^{37}\text{Cl}$ is 25%, average mass = $(35 \times 0.75) + (37 \times 0.25) = 35.5$

Properties of Light

Light exhibits both wave and particle properties. Key properties include speed, frequency, and wavelength.

• Speed of Light (c): $c = 3.00 \times 10^8 \text{ m/s}$

• Frequency (\nu): Number of cycles per second (Hz)

• Wavelength (\lambda): Distance between peaks (m)

• Formula: $c = \lambda \nu$

• Example: If $\lambda = 500 \text{ nm}$, $\nu = \frac{c}{\lambda}$

Calculations to Determine Frequency or Wavelength

• Formula: $\nu = \frac{c}{\lambda}$

• Example: For $\lambda = 600 \text{ nm}$, $\nu = \frac{3.00 \times 10^8}{6.00 \times 10^{-7}} = 5.00 \times 10^{14} \text{ Hz}$

Electron Configuration

Electron configuration describes the arrangement of electrons in an atom’s orbitals.

• Key Point: Electrons fill orbitals in order of increasing energy (Aufbau principle).

• Example: Sodium: 1s2 2s2 2p6 3s1

Identifying Elements from Electron Configuration

• Key Point: The total number of electrons equals the atomic number.

• Example: 1s2 2s2 2p6 3s2 3p1 corresponds to aluminum (atomic number 13).

Diagonal Rule for Predicting Order of Levels

The diagonal rule helps predict the order in which orbitals are filled:

• Order: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s

Additional info: The diagonal rule is a mnemonic for the Aufbau principle, guiding electron filling order based on energy levels.