CHEM 10003 Chemical Principles: Chapters 2–9 Mini-Study Guide

Chapter 2: Chemistry and Measurements

Section 2.1: Units of Measurement

Measurements in chemistry use standardized units to ensure consistency and accuracy. The metric system and SI (International System of Units) are commonly used.

• Volume: Liter (L), milliliter (mL)

• Length: Meter (m), centimeter (cm)

• Mass: Gram (g), kilogram (kg)

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

• Time: Second (s)

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 indicate precision.

• 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.

Section 2.3: Significant Figures in Calculation

Calculated answers must reflect the correct number of significant figures.

• 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 less than 5, leave the preceding digit unchanged; if 5 or greater, increase the preceding digit by 1.

Section 2.4: Prefixes and Equalities

Prefixes indicate multiples or fractions of base units.

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

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

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

• Equality Example: $1\ \text{m} = 100\ \text{cm}$

Section 2.5: Writing Conversion Factors

Conversion factors relate two units describing the same quantity.

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

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

Section 2.6: Problem Solving Using Unit Conversion

Dimensional analysis uses conversion factors to change units.

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

Section 2.7: Density

Density is the ratio of mass to volume.

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

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

• Applications: Use density to find mass or volume if the other is known.

Chapter 3: Matter and Energy

Section 3.1: Classification of Matter

Matter can be classified based on composition and uniformity.

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

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

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: Color, melting point, density

• Chemical Properties: Reactivity, flammability

Section 3.3: Temperature

Temperature can be converted between scales.

• Celsius to Kelvin: $K = ^\circ C + 273$

• Celsius to Fahrenheit: $^\circ F = (^\circ C \times \frac{9}{5}) + 32$

Section 3.4: Energy

Energy is the capacity to do work; it can be potential or kinetic.

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

• Kinetic Energy: Energy of motion

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

Section 3.5: Energy and Nutrition

Food energy is measured in calories (Cal), kilocalories (kcal), or kilojoules (kJ).

• Example: If a food contains 100 kcal, it is equivalent to $100 \times 4.184 = 418.4\ \text{kJ}$

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: $q = m \times c \times \Delta T$ where $q$ = heat, $m$ = mass, $c$ = specific heat, $\Delta T$ = change in temperature

• Prediction: Substances with higher specific heat require more energy to change temperature.

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/Cooling Curves: Show temperature changes during state transitions

Chapter 4: Atoms and Elements

Section 4.1: Elements and Symbols

Each element has a unique name and symbol.

• Example: Sodium (Na), Carbon (C), Oxygen (O)

• Group Examples: Group 1: H, Li, Na, K; Group 17: F, Cl, Br, I

Section 4.2: The Periodic Table

The periodic table organizes elements by groups and periods.

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

• Periods: Horizontal rows

• Classification: Metals, nonmetals, metalloids

Section 4.3: The Atom

Atoms consist of protons, neutrons, and electrons.

• Proton: Charge +1, mass ~1 amu, located in nucleus

• Neutron: Charge 0, mass ~1 amu, located in nucleus

• Electron: Charge -1, mass ~0.0005 amu, located 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

• Example: Carbon-12: 6 protons, 6 neutrons, 6 electrons

Section 4.5: Isotopes and Atomic Mass

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

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

• Example: $\text{^{14}_{6}C}$

Section 4.6: Electron Energy Levels

Electrons occupy energy levels around the nucleus.

• Example: Sodium (Na): 2, 8, 1

Section 4.7: Trend in Periodic Properties

Periodic properties depend on electron arrangement.

• Valence Electrons: Electrons in the outermost energy level

• Lewis Symbols: Dots around element symbol representing valence electrons

Chapter 5: Nuclear Chemistry

Section 5.1: Natural Radioactivity

Radioactive decay emits particles or energy.

• Alpha ($\alpha$): $^4_2He$; low penetration, shielded by paper

• Beta ($\beta$): $^0_{-1}e$; moderate penetration, shielded by plastic

• Positron ($\beta^+$): $^0_{+1}e$

• Gamma ($\gamma$): Energy; high penetration, shielded by lead

• Biological Effects: Can damage cells and DNA

Section 5.2: Nuclear Reactions

Nuclear equations show changes in atomic and mass numbers.

• Alpha Decay Example: $^A_ZX \rightarrow ^{A-4}_{Z-2}Y + ^4_2He$

• Beta Decay Example: $^A_ZX \rightarrow ^A_{Z+1}Y + ^0_{-1}e$

• Positron Emission Example: $^A_ZX \rightarrow ^A_{Z-1}Y + ^0_{+1}e$

• Gamma Emission: $^A_ZX^* \rightarrow ^A_ZX + \gamma$

Section 5.3: Half-Life of a Radioisotope

Half-life is the time for half the radioactive atoms to decay.

• Formula: $\text{Remaining} = \text{Initial} \times (\frac{1}{2})^n$ 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 (e.g., Na+)

• Anions: Negative ions (e.g., Cl-)

Section 6.2: Ionic Compounds

Ionic compounds are formed by charge balance between cations and anions.

• Example: Na+ + Cl- → NaCl

Section 6.3: Naming and Writing Ionic Formulas

Ionic compounds are named based on their ions.

• Representative Elements: Name cation, then anion (e.g., sodium chloride)

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

Section 6.4: Polyatomic Ions

Polyatomic ions are groups of atoms with a charge.

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

• Compound Example: NaNO3 (sodium nitrate)

Section 6.5: Molecular Compounds: Sharing Electrons

Molecular compounds are formed by sharing electrons.

• Example: H2O (water), CO2 (carbon dioxide)

Section 6.6: Lewis Structures for Molecules

Lewis structures show how atoms share electrons.

• Octet Rule: Atoms tend to have eight electrons in their valence shell

• Example: Water: H–O–H with two lone pairs on O

Section 6.7: Electronegativity and Bond Polarity

Electronegativity determines bond polarity.

• Polar Bond: Unequal sharing of electrons (e.g., H–Cl)

• Nonpolar Bond: Equal sharing (e.g., H–H)

Section 6.8: Shapes of Molecules

Molecular shape is predicted by VSEPR theory.

• Linear: CO2

• Bent: H2O

• Trigonal Planar: BF3

• Trigonal Pyramidal: NH3

• Tetrahedral: CH4

Section 6.9: Polarity of Molecules

Polarity depends on shape and bond polarity.

• Polar Molecule: H2O

• Nonpolar Molecule: CO2

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 per mole

• Example: 2 moles of Na = $2 \times 6.022 \times 10^{23}$ atoms

Section 7.2: Molar Mass

Molar mass is the mass of one mole of a substance.

• Example: H2O: $2 \times 1.01 + 16.00 = 18.02$ g/mol

Section 7.3: Calculations Using Molar Mass

Molar mass is used to convert between grams and moles.

• Formula: $\text{Moles} = \frac{\text{Mass (g)}}{\text{Molar Mass (g/mol)}}$

Section 7.4: Equations for Chemical Reactions

Chemical equations show reactants and products.

• Balanced Equation: Number of atoms is equal on both sides

• Example: $2H_2 + O_2 \rightarrow 2H_2O$

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 calculations.

• Example: $2H_2 + O_2 \rightarrow 2H_2O$ (2:1:2 ratio)

Section 7.8: Mass Calculations for Chemical Reactions

Mass of reactants and products can be calculated using mole ratios and molar mass.

• Example: Calculate mass of H2O produced from 4 g H2

Section 7.9: Energy in Chemical Reactions

Reactions can absorb or release energy.

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

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

• Factors Affecting Rate: Concentration, temperature, catalysts

Chapter 8: Gases

Section 8.1: Property of Gases

Gases are described by kinetic molecular theory.

• Properties: Pressure, volume, temperature, amount (moles)

• 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: Inspiration and expiration in lungs

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

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

• Formula: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$

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

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

• Formula: $\frac{P_1}{T_1} = \frac{P_2}{T_2}$

Section 8.5: The Combined Gas Law

The combined gas law relates pressure, volume, and temperature.

• 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 and solvent.

• Solute: Substance dissolved

• Solvent: Substance doing the dissolving (usually water)

• Formation: Solute particles disperse uniformly in solvent

Section 9.2: Electrolytes and Nonelectrolytes

Electrolytes conduct electricity; nonelectrolytes do not.

• Electrolytes: Ionic compounds (e.g., NaCl)

• Nonelectrolytes: Molecular compounds (e.g., sugar)

Section 9.3: Solubility

Solubility is the maximum amount of solute that can dissolve in a solvent.

• Unsaturated Solution: Can dissolve more solute

• Saturated Solution: Cannot dissolve more solute

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

• Solubility Table: Used to identify soluble/insoluble ionic compounds

Section 9.4: Solution Concentration

Concentration expresses the amount of solute in a solution.

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

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

• Mass/Volume Percent: $\frac{\text{mass of solute (g)}}{\text{volume of solution (mL)}} \times 100\%$

• Molarity: $\frac{\text{moles of solute}}{\text{liters of solution}}$

Section 9.5: Dilution of Solutions

Dilution reduces concentration by adding solvent.

• Formula: $C_1V_1 = C_2V_2$ where $C$ = concentration, $V$ = volume

• Example: To dilute 10 mL of 2 M solution to 1 M: $2 \times 10 = 1 \times V_2 \rightarrow V_2 = 20$ mL