CHEM 10003 Chemical Principles: Chapters 2–9 Study Guide

Chapter 2: Chemistry and Measurements

Section 2.1: Units of Measurement

Understanding units is fundamental in chemistry for accurate measurement and communication of scientific data.

• Volume: liter (L), milliliter (mL)

• Length: meter (m), centimeter (cm)

• Mass: gram (g), kilogram (kg)

• Temperature: Celsius (°C), Kelvin (K)

• Time: second (s)

Example: The SI unit for mass is the kilogram (kg).

Section 2.2: Measured Numbers and Significant Figures

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

• Measured Number: Obtained by measurement (e.g., 2.54 cm).

• Exact Number: Defined or counted (e.g., 12 eggs in a dozen).

• Significant Figures: All the digits in a measurement, including the uncertain last digit.

Example: 0.00450 has three significant figures.

Section 2.3: Significant Figures in Calculation

Calculated answers must reflect the precision of the measurements used.

• Multiplication/Division: Result has the same number of significant figures as the measurement with the fewest significant figures.

• Addition/Subtraction: Result has the same number of decimal places as the measurement with the fewest decimal places.

• Rounding: If the digit to be dropped is 5 or greater, round up.

Section 2.4: Prefixes and Equalities

Prefixes indicate multiples or fractions of units.

• Kilo- (k): $1\ \text{km} = 1000\ \text{m}$

• Centi- (c): $1\ \text{cm} = 0.01\ \text{m}$

• Milli- (m): $1\ \text{mL} = 0.001\ \text{L}$

Section 2.5: Writing Conversion Factors

Conversion factors are ratios that express how many of one unit are equal to another unit.

• Example: $1\ \text{in} = 2.54\ \text{cm}$ gives two conversion factors: $\frac{1\ \text{in}}{2.54\ \text{cm}}$ and $\frac{2.54\ \text{cm}}{1\ \text{in}}$

Section 2.6: Problem Solving Using Unit Conversion

Dimensional analysis uses conversion factors to change units.

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

Section 2.7: Density

Density relates the mass and volume of a substance.

• Formula: $\text{Density} = \frac{\text{mass}}{\text{volume}}$

• Units: g/mL or g/cm3

• Applications: Used to identify substances and convert between mass and volume.

Example: If a liquid has a mass of 25 g and a volume of 20 mL, its density is $\frac{25\ \text{g}}{20\ \text{mL}} = 1.25\ \text{g/mL}$

Chapter 3: Matter and Energy

Section 3.1: Classification of Matter

Matter can be classified based on its composition and uniformity.

• Pure Substances: Elements (e.g., O2) and compounds (e.g., H2O).

• Mixtures: Homogeneous (uniform, e.g., salt water) or heterogeneous (not uniform, e.g., salad).

Section 3.2: States and Properties of Matter

Matter exists in different states and has physical and chemical properties.

• States: Solid, liquid, gas.

• Physical Properties: Observed without changing composition (e.g., melting point).

• Chemical Properties: Observed when a substance changes composition (e.g., flammability).

Section 3.3: Temperature

Temperature can be measured in Celsius, Kelvin, or Fahrenheit.

• Conversion: $T_{\text{K}} = T_{\text{°C}} + 273.15$

• Conversion: $T_{\text{°F}} = 1.8 \times T_{\text{°C}} + 32$

Section 3.4: Energy

Energy is the capacity to do work and exists in different forms.

• Kinetic Energy: Energy of motion.

• Potential Energy: Stored energy.

• Units: Joule (J), calorie (cal)

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

Section 3.5: Energy and Nutrition

Food energy is measured in Calories (nutritional calories), kilocalories, or kilojoules.

• 1 Calorie (Cal) = 1 kilocalorie (kcal) = 1000 calories (cal)

• 1 kcal = 4.184 kJ

Section 3.6: Specific Heat

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

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

• q: heat (J or cal), m: mass (g), c: specific heat, ΔT: temperature change (°C)

Section 3.7: Changes of State

Matter changes state through physical processes.

• Melting: Solid to liquid

• Freezing: Liquid to solid

• Sublimation: Solid to gas

• Deposition: Gas to solid

• Vaporization: Liquid to gas

• Condensation: Gas to liquid

Heating and Cooling Curves: Graphs that show temperature changes as heat is added or removed.

Chapter 4: Atoms and Elements

Section 4.1: Elements and Symbols

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

• Example: Sodium = Na, Potassium = K

Section 4.2: The Periodic Table

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

• Groups: Vertical columns (e.g., Group 1: Alkali metals)

• Periods: Horizontal rows

• Metals, Nonmetals, Metalloids: Classified by properties

Section 4.3: The Atom

Atoms consist of protons, neutrons, and electrons.

• Proton: Positive charge, mass ≈ 1 amu, in nucleus

• Neutron: No charge, mass ≈ 1 amu, in nucleus

• Electron: Negative charge, mass ≈ 0 amu, outside nucleus

Section 4.4: Atomic Number and Mass Number

The atomic number is the number of protons; the mass number is the sum of protons and neutrons.

• Number of electrons: Equal to protons in a neutral atom

• Formula: $\text{Mass number} = \text{protons} + \text{neutrons}$

Section 4.5: Isotopes and Atomic Mass

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

• Atomic Symbol: $^{A}_{Z}\text{X}$, where A = mass number, Z = atomic number, X = element symbol

Section 4.6: Electron Energy Levels

Electrons occupy energy levels (shells) around the nucleus.

• Example: Sodium (Na): 2, 8, 1

Section 4.7: Trend in Periodic Properties

Valence electrons determine chemical properties and are shown in Lewis symbols.

• Lewis Symbol: Element symbol with dots for valence electrons

Chapter 5: Nuclear Chemistry

Section 5.1: Natural Radioactivity

Some nuclei are unstable and emit radiation.

• Alpha (α): $^4_2\text{He}$, low penetration

• Beta (β): $^0_{-1}\text{e}$, moderate penetration

• Positron (β+): $^0_{+1}\text{e}$

• Gamma (γ): High energy, no mass or charge, high penetration

• Protection: Paper (α), clothing (β), lead/concrete (γ)

• Biological Effects: Can damage cells and DNA

Section 5.2: Nuclear Reactions

Nuclear equations show changes in atomic number and mass number.

• Alpha Decay Example: $^{238}_{92}\text{U} \rightarrow ^{234}_{90}\text{Th} + ^4_2\text{He}$

• Beta Decay Example: $^{14}_6\text{C} \rightarrow ^{14}_7\text{N} + ^0_{-1}\text{e}$

Section 5.3: Half-Life of a Radioisotope

Half-life is the time for half of a radioactive sample to decay.

• Formula: $\text{Amount remaining} = \text{initial amount} \times \left(\frac{1}{2}\right)^n$

• n: Number of half-lives elapsed

Section 5.4: Nuclear Fission and Fusion

Fission splits heavy nuclei; fusion combines light nuclei.

• Fission: Used in nuclear power plants

• Fusion: Powers the sun and stars

Chapter 6: Ionic and Molecular Compounds

Section 6.1: Ions: Transfer of Electrons

Atoms gain or lose electrons to form ions.

• Cations: Positive ions (e.g., Na+)

• Anions: Negative ions (e.g., Cl-)

Section 6.2: Ionic Compounds

Ionic compounds are formed from cations and anions in ratios that balance charge.

• Example: Na+ and Cl- form NaCl

Section 6.3: Naming and Writing Ionic Formulas

Ionic compounds are named by cation then anion; transition metals may use Roman numerals.

• Example: FeCl2 is iron(II) chloride

Section 6.4: Polyatomic Ions

Polyatomic ions are groups of atoms with a charge.

• Example: NO3- (nitrate), SO42- (sulfate)

• Example Compound: NaNO3 (sodium nitrate)

Section 6.5: Molecular Compounds: Sharing Electrons

Molecular compounds are formed by sharing electrons between nonmetals.

• Example: CO2 is carbon dioxide

Section 6.6: Lewis Structures for Molecules

Lewis structures show how atoms share electrons to achieve octets.

• Single Bond: One pair of shared electrons

• Double/Triple Bonds: Two/three pairs of shared electrons

Section 6.7: Electronegativity and Bond Polarity

Electronegativity differences determine bond polarity.

• Nonpolar Covalent: Electrons shared equally

• Polar Covalent: Electrons shared unequally

Section 6.8: Shapes of Molecules

The shape of a molecule is predicted by VSEPR theory.

• Linear: 180° (e.g., CO2)

• Bent: <120° or 104.5° (e.g., H2O)

• Trigonal Planar: 120° (e.g., BF3)

• Trigonal Pyramidal: 107° (e.g., NH3)

• Tetrahedral: 109.5° (e.g., CH4)

Section 6.9: Polarity of Molecules

The overall shape and bond polarities determine if a molecule is polar or nonpolar.

• Example: H2O is polar; CO2 is nonpolar

Chapter 7: Chemical Quantities and Reactions

Section 7.1: The Mole

The mole is a counting unit for atoms, molecules, or ions.

• Avogadro's Number: $6.022 \times 10^{23}$ particles/mol

Section 7.2: Molar Mass

Molar mass is the mass of one mole of a substance (g/mol).

• Example: H2O: 2(1.01) + 16.00 = 18.02 g/mol

Section 7.3: Calculations Using Molar Mass

Molar mass is used to convert between grams and moles.

• Example: $\text{moles} = \frac{\text{grams}}{\text{molar mass}}$

Section 7.4: Equations for Chemical Reactions

Chemical equations must be balanced to obey the law of conservation of mass.

• Example: $2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$

Section 7.6: Oxidation-Reduction Reactions

Oxidation is loss of electrons; reduction is gain of electrons.

• Inorganic Example: Zn + Cu2+ → Zn2+ + Cu

• Biological Example: NAD+ + 2H → NADH + H+

Section 7.7: Mole Relationships in Chemical Equations

Balanced equations provide mole ratios for reactants and products.

• Example: From $2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$, 2 mol H2 reacts with 1 mol O2

Section 7.8: Mass Calculations for Chemical Reactions

Stoichiometry uses molar mass and mole ratios to relate masses of reactants and products.

• Example: Calculate grams of H2O produced from 4.0 g H2

Section 7.9: Energy in Chemical Reactions

Reactions can absorb (endothermic) or release (exothermic) energy.

• Exothermic: Releases heat (e.g., combustion)

• Endothermic: Absorbs heat (e.g., photosynthesis)

• Factors Affecting Rate: Temperature, concentration, catalysts

Chapter 8: Gases

Section 8.1: Property of Gases

Gases have unique properties explained by the kinetic molecular theory.

• Properties: Compressible, expand to fill container, low density

• Units: Pressure (atm, mmHg), volume (L), temperature (K), amount (mol)

Section 8.2: Pressure and Volume (Boyle’s Law)

Boyle’s Law relates pressure and volume at constant temperature and amount.

• Formula: $P_1V_1 = P_2V_2$

• Application: Breathing involves changes in lung volume and pressure.

Section 8.3: Temperature and Volume (Charles’s Law)

Charles’s Law relates temperature and volume at constant pressure and amount.

• Formula: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (T in K)

Section 8.4: Temperature and Pressure (Gay-Lussac’s Law)

Gay-Lussac’s Law relates temperature and pressure at constant volume and amount.

• Formula: $\frac{P_1}{T_1} = \frac{P_2}{T_2}$ (T in K)

Section 8.5: The Combined Gas Law

The combined gas law relates pressure, volume, and temperature for a fixed amount of gas.

• Formula: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

Chapter 9: Solutions

Section 9.1: Solutions

A solution is a homogeneous mixture of solute dissolved in solvent.

• Solute: Substance present in lesser amount

• Solvent: Substance present in greater amount

Section 9.2: Electrolytes and Nonelectrolytes

Electrolytes conduct electricity in solution; nonelectrolytes do not.

• Electrolyte: NaCl (dissociates into ions)

• Nonelectrolyte: Sugar (does not dissociate)

Section 9.3: Solubility

Solubility is the maximum amount of solute that dissolves in a given amount of solvent at a specific temperature.

• Unsaturated Solution: Less than maximum solute dissolved

• Saturated Solution: Maximum solute dissolved

• Temperature Effects: Solubility of solids increases with temperature; gases decrease

• Solubility Table: Used to predict if an ionic compound is soluble

Section 9.4: Solution Concentration

Concentration expresses the amount of solute in a given amount of solution.

• Mass Percent: $\frac{\text{g solute}}{\text{g solution}} \times 100\%$

• Volume Percent: $\frac{\text{mL solute}}{\text{mL solution}} \times 100\%$

• Mass/Volume Percent: $\frac{\text{g solute}}{\text{mL solution}} \times 100\%$

• Molarity (M): $\frac{\text{mol solute}}{\text{L solution}}$

Section 9.5: Dilution of Solutions

Dilution decreases the concentration of a solution by adding more solvent.

• Formula: $C_1V_1 = C_2V_2$

• Application: Used to prepare solutions of desired concentration