CHEM 1110 Exam 3 Study Guide: Nuclear Chemistry, Ionic and Molecular Compounds

General Exam Review Topics

• Definitions of new vocabulary: Ensure understanding of all key terms introduced in each chapter.

• Calculations: Practice calculations from chapters 5 and 6, focusing on problem-solving steps and correct use of units.

• General units: Know what each unit measures and how to convert between units.

• Homework and practice problems: Review assigned homework and Mastering Chemistry questions for additional practice.

• Class discussions: Revisit important points discussed in class for deeper understanding.

Chapter 5: Nuclear Chemistry

Types of Radiation

Nuclear chemistry involves the study of changes in atomic nuclei, including radioactive decay and nuclear reactions. Understanding the types of radiation is essential for interpreting nuclear processes.

• Alpha (α) radiation: Consists of helium nuclei (2 protons, 2 neutrons). Low penetration, stopped by paper.

• Beta (β) radiation: Consists of electrons or positrons. Moderate penetration, stopped by aluminum.

• Gamma (γ) radiation: High-energy electromagnetic waves. High penetration, requires thick lead or concrete for shielding.

• Positron (β+) emission: Emission of a positron (antiparticle of electron).

• Electron capture: Nucleus captures an inner electron, converting a proton to a neutron.

Sources and Effects of Radiation

Radiation can originate from natural or artificial sources, and its effects depend on type, energy, and exposure duration.

• Natural sources: Cosmic rays, radon gas, naturally occurring radioactive isotopes.

• Artificial sources: Medical imaging, nuclear reactors, industrial applications.

• Units of radiation: Common units include becquerel (Bq) for activity, gray (Gy) for absorbed dose, and sievert (Sv) for biological effect.

Balancing Nuclear Equations

Nuclear equations must be balanced for both mass number and atomic number.

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

• Atomic number (Z): Number of protons.

• Example: $^{14}_{6}\text{C} \rightarrow ^{14}_{7}\text{N} + \beta^{-}$

Half-Life Calculations

The half-life of a radioactive isotope is the time required for half of the sample to decay.

• Formula: $N = N_0 \left(\frac{1}{2}\right)^{t/t_{1/2}}$

• Application: Used to determine remaining quantity of a substance after a given time.

Fusion and Fission Reactions

Nuclear reactions can involve the splitting (fission) or combining (fusion) of nuclei.

• Fission: Large nucleus splits into smaller nuclei, releasing energy. Example: $^{235}\text{U} + n \rightarrow ^{141}\text{Ba} + ^{92}\text{Kr} + 3n$

• Fusion: Small nuclei combine to form a larger nucleus, releasing energy. Example: $^{2}\text{H} + ^{3}\text{H} \rightarrow ^{4}\text{He} + n$

Chapter 6: Ionic and Molecular Compounds

Identifying Ionic and Covalent Compounds

Chemical compounds are classified as ionic or covalent based on the nature of the bonding between atoms.

• Ionic compounds: Formed from metals and nonmetals; involve transfer of electrons.

• Covalent compounds: Formed from nonmetals; involve sharing of electrons.

• Example: NaCl (ionic), H2O (covalent)

Formulas and Naming of Compounds

Systematic naming and formula writing are essential for clear communication in chemistry.

• Binary ionic compounds: Metal + nonmetal, e.g., MgCl2

• Polyatomic ions: Charged groups of atoms, e.g., SO42-, NO3-

• Transition metals: Use Roman numerals to indicate charge, e.g., FeCl3 (iron(III) chloride)

• Covalent compounds: Use prefixes to indicate number of atoms, e.g., CO2 (carbon dioxide)

Polyatomic Ions and Nomenclature

Polyatomic ions are important in many common compounds. Their names and formulas must be memorized.

• Common polyatomic ions: Nitrate (NO3-), sulfate (SO42-), phosphate (PO43-), ammonium (NH4+)

• Nomenclature: Name cation first, then anion. For covalent compounds, use prefixes (mono-, di-, tri-, etc.).

Lewis Structures and Molecular Geometry

Lewis structures represent the arrangement of electrons in molecules, which determines molecular shape and polarity.

• Drawing Lewis structures: Show all valence electrons, bonds, and lone pairs.

• Electron group geometry: Determined by the number of electron groups around the central atom.

• VSEPR theory: Predicts molecular shape based on electron repulsion.

• Common shapes: Linear, bent, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

• Example: H2O is bent due to two lone pairs on oxygen.

Polarity and Intermolecular Forces

Molecular polarity affects physical properties and intermolecular interactions.

• Polarity: Determined by difference in electronegativity and molecular geometry.

• Intermolecular forces: Include dispersion (London) forces, dipole-dipole interactions, hydrogen bonding, and ion-dipole forces.

• Example: Water exhibits hydrogen bonding, leading to high boiling point.

Melting and Boiling Points

The melting point (MP) and boiling point (BP) of a substance are influenced by the type and strength of intermolecular forces present.

• Ionic compounds: Generally have high MP and BP due to strong electrostatic forces.

• Molecular compounds: MP and BP vary depending on polarity and intermolecular forces.

• Comparison: NaCl (ionic) has a much higher MP than H2O (molecular).

Table: Common Polyatomic Ions

This table summarizes several important polyatomic ions and their formulas.

Name	Formula	Charge
Nitrate	NO3	-1
Sulfate	SO4	-2
Phosphate	PO4	-3
Ammonium	NH4	+1
Hydroxide	OH	-1

Table: Molecular Shapes (VSEPR Theory)

This table lists common molecular shapes based on the number of electron groups and lone pairs.

Electron Groups	Lone Pairs	Shape	Example
2	0	Linear	CO2
3	0	Trigonal planar	BF3
4	0	Tetrahedral	CH4
4	2	Bent	H2O