Exam 2 Study Guide: Nuclear Chemistry, Ionic & Molecular Compounds, Chemical Reactions, and Gases

Chapter 5: Nuclear Chemistry

Balanced Nuclear Equations

Nuclear equations represent changes in atomic nuclei, showing both mass number and atomic number for all reactants and products.

• Mass Number (A): Total number of protons and neutrons.

• Atomic Number (Z): Number of protons.

• Balanced Equation: The sum of mass numbers and atomic numbers must be equal on both sides.

• Example: Alpha decay of uranium-238: $\ ^{238}_{92}U \rightarrow \ ^{234}_{90}Th + \ ^{4}_{2}He$

Types of Nuclear Decay

Nuclear decay involves emission of particles or energy from unstable nuclei.

• Alpha Decay (α): Emission of an alpha particle ($\ ^{4}_{2}He$).

• Beta Decay (β): Emission of a beta particle ($\ ^{0}_{-1}e$).

• Gamma Decay (γ): Emission of gamma rays (energy, no mass or charge).

• Positron Emission: Emission of a positron ($\ ^{0}_{+1}e$).

• Electron Capture: Nucleus captures an inner electron.

Units in Nuclear Chemistry

• Becquerel (Bq): SI unit for radioactivity (1 disintegration/sec).

• Curie (Ci): Older unit (1 Ci = $3.7 \times 10^{10}$ disintegrations/sec).

• Gray (Gy): Unit for absorbed dose (1 Gy = 1 J/kg).

• Rem: Unit for biological effect.

Producing Radioactive Isotopes

• Artificial Transmutation: Bombarding stable nuclei with particles to produce radioisotopes.

• Example: $\ ^{14}_{7}N + \ ^{4}_{2}He \rightarrow \ ^{17}_{8}O + \ ^{1}_{1}H$

Half-Life

The half-life is the time required for half the atoms in a radioactive sample to decay.

• Formula: $\text{Remaining amount} = \text{Initial amount} \times (\frac{1}{2})^{\text{number of half-lives}}$

• Application: Used in dating, medical diagnostics, and nuclear waste management.

Nuclear Fission and Fusion

• Fission: Splitting of a heavy nucleus into lighter nuclei, releasing energy.

• Fusion: Combining of light nuclei to form a heavier nucleus, releasing more energy than fission.

• Example (Fission): $\ ^{235}_{92}U + \ ^{1}_{0}n \rightarrow \ ^{141}_{56}Ba + \ ^{92}_{36}Kr + 3\ ^{1}_{0}n$

• Example (Fusion): $\ ^{2}_{1}H + \ ^{3}_{1}H \rightarrow \ ^{4}_{2}He + \ ^{1}_{0}n$

Chapter 6: Ionic and Molecular Compounds

Properties of Ionic Compounds

Ionic compounds are formed from the electrostatic attraction between cations and anions.

• High melting and boiling points

• Conduct electricity when dissolved in water

• Usually crystalline solids

Writing Ionic Formulas and Charges

• Formula: Combine ions so total charge is zero.

• Variable Charge Metals: Use Roman numerals to indicate charge (e.g., Fe(III) for $Fe^{3+}$).

• Example: $Na^+ + Cl^- \rightarrow NaCl$

Polyatomic Ions and Writing Formulas

• Polyatomic Ion: A charged group of covalently bonded atoms (e.g., $NO_3^-$, $SO_4^{2-}$).

• Formula Example: $Ca^{2+} + 2NO_3^- \rightarrow Ca(NO_3)_2$

Names and Formulas of Molecular Compounds

• Molecular Compounds: Formed by sharing electrons between nonmetals.

• Naming: Use prefixes (mono-, di-, tri-, etc.) to indicate number of atoms.

• Example: $CO_2$ is carbon dioxide; $N_2O$ is dinitrogen monoxide.

Lewis Structure

Lewis structures show the arrangement of electrons in molecules.

• Steps: Count valence electrons, arrange atoms, distribute electrons to satisfy octet rule.

• Example: Lewis structure for $H_2O$ shows two lone pairs on oxygen.

Electronegativity and Bond Polarity

• Electronegativity: Ability of an atom to attract electrons in a bond.

• Bond Polarity: Difference in electronegativity creates polar bonds.

• Example: $H_2O$ is polar due to oxygen's higher electronegativity.

Shapes of Molecules (Simple)

• Linear: 2 atoms or 3 atoms with no lone pairs (e.g., $CO_2$).

• Bent: 3 atoms with lone pairs (e.g., $H_2O$).

• Trigonal Planar: 3 atoms bonded to central atom (e.g., $BF_3$).

• Tetrahedral: 4 atoms bonded to central atom (e.g., $CH_4$).

Chapter 7: Chemical Reactions and Quantities

Writing a Balanced Chemical Equation

Balanced equations show equal numbers of each atom on both sides.

• Steps: Identify reactants and products, balance atoms by adjusting coefficients.

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

Types of Chemical Reactions

• Combination: Two or more substances form one product.

• Decomposition: One substance breaks into two or more products.

• Single Replacement: One element replaces another in a compound.

• Double Replacement: Exchange of ions between two compounds.

• Combustion: Substance reacts with oxygen, producing energy.

Concept of Mole and Avogadro’s Number

• Mole: Amount of substance containing $6.022 \times 10^{23}$ particles (Avogadro’s number).

• Application: Used to relate mass, particles, and volume in chemical calculations.

Molar Mass

• Molar Mass: Mass of one mole of a substance (g/mol).

• Calculation: Sum atomic masses from periodic table.

• Example: Molar mass of $H_2O$ is $2 \times 1.01 + 16.00 = 18.02$ g/mol.

Mole-Mole Factor Concept

• Mole Ratio: Ratio of moles of reactants and products from balanced equation.

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

Mass Calculation for Chemical Reaction

• Steps: Convert mass to moles, use mole ratio, convert moles to mass.

• Formula: $\text{Mass} = \text{Moles} \times \text{Molar Mass}$

Limiting Reactants and Percent Yield

• Limiting Reactant: Reactant that is completely consumed, limits product formed.

• Percent Yield: $\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$

• Application: Used to assess efficiency of reactions.

Exothermic and Endothermic Reactions

• Exothermic: Releases heat ($\Delta H < 0$).

• Endothermic: Absorbs heat ($\Delta H > 0$).

• Example: Combustion is exothermic; photosynthesis is endothermic.

Chapter 8: Gases

Four Properties of Gases

• Pressure (P): Force exerted by gas particles.

• Volume (V): Space occupied by gas.

• Temperature (T): Measure of kinetic energy.

• Amount (n): Number of moles.

Gas Laws

Gas laws relate the properties of gases mathematically.

• Boyle’s Law: $P_1V_1 = P_2V_2$ (at constant T and n)

• Charles’s Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (at constant P and n)

• Gay-Lussac’s Law: $\frac{P_1}{T_1} = \frac{P_2}{T_2}$ (at constant V and n)

• Combined Gas Law: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

• Ideal Gas Law: $PV = nRT$

Important Equalities (Except R Value)

• Standard Pressure: 1 atm = 760 mmHg = 101.3 kPa

• Standard Temperature: 0°C = 273 K

• Volume of 1 mole at STP: 22.4 L

Application of Blood Gases to Health

Blood gases are crucial for respiratory and metabolic health.

• Oxygen (O2): Essential for cellular respiration.

• Carbon Dioxide (CO2): Waste product, affects blood pH.

• Clinical Application: Measurement of blood gases helps diagnose respiratory and metabolic disorders.

• Example: Low O2 or high CO2 can indicate hypoxia or respiratory failure.