CHEM 10003 Chemical Principles: Chapters 2–9 Study Guide

Chapter 2: Chemistry and Measurements

Section 2.1: Units of Measurement

Understanding the metric and SI units is fundamental for all chemical measurements.

• Volume: liter (L), milliliter (mL)

• Length: meter (m), centimeter (cm)

• Mass: gram (g), kilogram (kg)

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

• Time: second (s)

Example: 1 L = 1000 mL

Section 2.2: Measured Numbers and Significant Figures

Measured numbers are obtained by measurement, while exact numbers are counted or defined.

• Significant Figures: Digits in a measured number that include all certain digits plus one estimated digit.

• Determining Significant Figures: 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.

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: The answer has as many significant figures as the measurement with the fewest significant figures.

• Addition/Subtraction: The answer has as many decimal places as the measurement with the fewest decimal places.

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

Section 2.4: Prefixes and Equalities

Metric prefixes indicate multiples or fractions of units.

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

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

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

Example: 1 km = 1000 m

Section 2.5: Writing Conversion Factors

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

• Metric: $1\ \text{kg} = 1000\ \text{g}$

• US System: $1\ \text{in} = 2.54\ \text{cm}$ (exact)

Section 2.6: Problem Solving Using Unit Conversion

Dimensional analysis uses conversion factors to convert between units.

• Set up the problem so units cancel, leaving the desired unit.

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

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

• Can be used to find mass or volume if the other is known.

Example: If a sample has a mass of 10 g and a volume of 2 mL, its density is $5\ \text{g/mL}$.

Chapter 3: Matter and Energy

Section 3.1: Classification of Matter

Matter can be classified by 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 three 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 into another (e.g., flammability)

Section 3.3: Temperature

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

• Conversion formulas:

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

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

Section 3.4: Energy

Energy is the capacity to do work and exists as potential or kinetic energy.

• Potential Energy: Stored energy (e.g., chemical bonds)

• Kinetic Energy: Energy of motion

• 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 (1 Cal = 1 kcal = 1000 cal).

• Calculate energy from food using nutritional values (carbohydrates, fats, proteins).

Section 3.6: Specific Heat

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

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

• Where $q$ = heat (J), $m$ = mass (g), $c$ = specific heat (J/g·°C), $\Delta T$ = 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 show temperature changes during state changes.

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 position and 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, negligible mass, outside nucleus

Section 4.4: Atomic Number and Mass Number

Atomic number is the number of protons; mass number is protons plus neutrons.

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

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.

• First 20 elements: Electron arrangements follow the 2-8-8 rule.

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): Helium nucleus, low penetration

• Beta (β): Electron, moderate penetration

• Positron (β+): Positive electron

• Gamma (γ): High-energy photon, high penetration

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

• Biological effects: Can damage living tissue

Section 5.2: Nuclear Reactions

Nuclear equations show changes in atomic and mass numbers during decay.

• Alpha decay: $^{A}_{Z}\text{X} \rightarrow ^{A-4}_{Z-2}\text{Y} + ^{4}_{2}\text{He}$

• Beta decay: $^{A}_{Z}\text{X} \rightarrow ^{A}_{Z+1}\text{Y} + ^{0}_{-1}\text{e}$

• Positron emission: $^{A}_{Z}\text{X} \rightarrow ^{A}_{Z-1}\text{Y} + ^{0}_{+1}\text{e}$

• Gamma emission: $^{A}_{Z}\text{X}^* \rightarrow ^{A}_{Z}\text{X} + \gamma$

Section 5.3: Half-Life of a Radioisotope

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

• Formula: $N = N_0 \times (\frac{1}{2})^n$ where n = number of half-lives

Section 5.4: Nuclear Fission and Fusion

Fission splits heavy nuclei; fusion combines light nuclei.

• Fission: Used in nuclear reactors

• Fusion: Powers the sun

Chapter 6: Ionic and Molecular Compounds

Section 6.1: Ions: Transfer of Electrons

Atoms gain or lose electrons to form ions.

• Cations: Positive ions (loss of electrons)

• Anions: Negative ions (gain of electrons)

Section 6.2: Ionic Compounds

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

• Formula: Subscripts indicate the ratio of ions

Section 6.3: Naming and Writing Ionic Formulas

Names and formulas reflect the ions present.

• Transition metals: Use Roman numerals for charge (e.g., iron(III) chloride)

Section 6.4: Polyatomic Ions

Polyatomic ions are groups of atoms with a charge.

• Example: Nitrate, NO3-

Section 6.5: Molecular Compounds: Sharing Electrons

Molecular compounds form when atoms share electrons.

• Naming: Use prefixes (mono-, di-, tri-, etc.)

Section 6.6: Lewis Structures for Molecules

Lewis structures show bonding and lone pairs for molecules obeying the octet rule.

Section 6.7: Electronegativity and Bond Polarity

Electronegativity differences determine bond polarity.

• Nonpolar: Electrons shared equally

• Polar: Electrons shared unequally

Section 6.8: Shapes of Molecules

Molecular geometry is predicted by VSEPR theory.

• Shapes: Linear, bent, trigonal planar, trigonal pyramidal, tetrahedral

Section 6.9: Polarity of Molecules and Intermolecular Forces

The shape and bond polarity determine if a molecule is polar or 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).

Section 7.3: Calculations Using Molar Mass

Use molar mass to convert between grams and moles.

• Formula: $\text{Moles} = \frac{\text{grams}}{\text{molar mass}}$

Section 7.4: Equations for Chemical Reactions

Chemical equations show reactants and products in a reaction.

• Must be balanced to conserve mass and atoms.

Section 7.6: Oxidation-Reduction Reactions

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

• Inorganic: Identify oxidized/reduced species

• Biological: Often involves transfer of hydrogen or oxygen

Section 7.7: Mole Relationships in Chemical Equations

Balanced equations provide mole ratios for reactants and products.

Section 7.8: Mass Calculations for Chemical Reactions

Use stoichiometry to relate masses of substances in a reaction.

Section 7.9: Energy in Chemical Reactions

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

• Factors affecting rate: Concentration, temperature, 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: At constant temperature and amount, pressure and volume are inversely related.

$P_1V_1 = P_2V_2$

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

Charles’s Law: At constant pressure and amount, volume and temperature are directly related.

$\frac{V_1}{T_1} = \frac{V_2}{T_2}$

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

Gay-Lussac’s Law: At constant volume and amount, pressure and temperature are directly related.

$\frac{P_1}{T_1} = \frac{P_2}{T_2}$

Section 8.5: The Combined Gas Law

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

$\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

Chapter 9: Solutions

Section 9.1: Solutions

A solution is a homogeneous mixture of solute and 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.

Section 9.3: Solubility

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

• Unsaturated: Less than maximum solute dissolved

• Saturated: Maximum solute dissolved

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

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 reduces concentration by adding solvent.

$C_1V_1 = C_2V_2$

Where $C$ = concentration, $V$ = volume