Key Concepts in Chemistry

Chapter 1: Essential Ideas

Section 1.1 Chemistry in Context

• Strive to recognize chemistry in everyday life.

• Understand and apply the scientific method.

• Describe personal experiences using the scientific method.

• Difference between Theory and Law:

  • Theory: A well-substantiated explanation of an aspect of the natural world based on a body of evidence.

  • Law: A statement based on repeated experimental observations that describe some aspects of the universe.

• Domains of Chemistry:

  • Macroscopic domain: The realm of ordinary things that are large enough to see with the naked eye.

  • Microscopic domain: The realm of atoms and molecules that can only be seen through special tools.

  • Symbolic domain: The use of symbols (chemical formulas, equations) to represent elements and compounds.

Section 1.2 Phases and Classification of Matter

• Distinguish between:

  • Solid: Definite shape and volume.

  • Liquid: Definite volume but takes the shape of its container.

  • Gas: No definite volume or shape.

  • Plasma: Ionized gas with free charged particles.

• Mass vs Weight:

  • Mass: Amount of matter in an object (measured in kilograms).

  • Weight: Gravitational force acting on the mass (measured in newtons).

• Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction.

  • Calculations: Can be used to predict the amounts of reactants and products in reactions.

• Distinguish between:

  • Elements: Pure substances that consist of one type of atom.

  • Compounds: Substances formed when two or more elements are chemically bonded.

• Pure Substance vs Mixture:

  • Pure Substance: Material with a uniform composition (elements and compounds).

  • Mixture: A combination of two or more substances that retain their individual properties.

• Physical Change vs Chemical Change:

  • Physical Change: Change that does not affect the chemical composition (e.g., melting, freezing).

  • Chemical Change: Change that results in the formation of new chemical substances (e.g., combustion).

• Physical Properties: Characteristics that can be observed without changing the identity of a substance (e.g., color, boiling point).

• Homogeneous Mixture vs Heterogeneous Mixture:

  • Homogeneous Mixture: Composition is uniform throughout (e.g., saltwater).

  • Heterogeneous Mixture: Composition is not uniform (e.g., salad).

• Understanding phase in context of mixtures and states of matter.

Section 1.3 Physical and Chemical Properties

• Physical Properties: Properties that can be observed or measured without changing the substance's chemical identity (e.g., melting point, boiling point).

• Chemical Properties: Properties that become evident during a chemical reaction, indicating how a substance interacts with others.

• Distinguish between:

  • Extensive Physical Properties: Depend on the amount of substance present (e.g., mass, volume).

  • Intensive Physical Properties: Independent of the amount of substance (e.g., density, temperature).

• Recognize:

  • Physical Changes: Changes that affect one or more physical properties of a substance.

  • Chemical Changes: Changes that result in the formation of new chemical substances.

Section 1.4 Measurements

• Measurements include:

  • Magnitude: The size or amount of a quantity.

  • Unit: Standardized quantity used to specify measurements (e.g., meter, liter).

  • Uncertainty: Means precision or reliability of a measured value.

• Significant Digits: The number of meaningful digits in a measurement that contributes to its accuracy.

• Use of:

  • Significant Figures: Reflects the precision of a measured value.

  • Scientific Notation: Represents very large or very small numbers efficiently.

  • SI Base Units: Seven fundamental units for measurement (e.g., length, mass, time) as listed in Table 1.2 of OpenStax text.

  • SI Prefixes: Represent powers of ten as shown in Table 1.3 of OpenStax text.

  • Derived Units: Units obtained from combining base units (e.g., volume in cubic meters, density in kg/m³).

• Density Calculations: Relation of mass to volume, expressed as:

  • ext{Density} = rac{ ext{Mass}}{ ext{Volume}}.

Section 1.5 Measurement Uncertainty, Accuracy, and Precision

• Scientific notation in calculations helps manage the scale of numbers.

• Report the correct number of significant digits in the results of calculations.

• Exact Numbers: Counted quantities or defined quantities that have no uncertainty.

• Rounding: Adjusting numbers to reflect the appropriate amount of significant figures.

• Precision vs Accuracy:

  • Precision: How consistent repeated measurements are to one another.

  • Accuracy: How close a measurement is to the actual or true value.

Section 1.6 Mathematical Treatment of Results

• Convert between temperature scales:

  • Celsius (°C), Fahrenheit (°F), and Kelvin (K).

• Perform calculations using the factor-label method or dimensional analysis:

  • Factor-label approach leverages conversion factors to change units in calculations.

• Equality: Provides two conversion factors useful for unit conversions.

• Molar Mass: Acts as a conversion factor linking grams to moles and vice versa.

Chapter 2: Atoms, Molecules, and Ions

Section 2.1 Early Ideas in Atomic Theory

• Understand the postulates of Dalton’s Atomic Theory:

  • Each element is composed of atoms.

  • Atoms of a given element are identical, while atoms of different elements are different.

  • Compounds are formed when atoms of different elements combine in fixed ratios.

  • A chemical reaction is a rearrangement of atoms.

• Atomic Theory Components:

  • Atomic symbols: Recognize and understand their meaning.

  • Become familiar with symbols listed in Table 2.1 of OpenStax text.

• Law of Multiple Proportions: When two elements form more than one compound, the ratios of the masses can be expressed as small whole numbers.

Section 2.2 Evolution of Atomic Theory

• Important experiments and findings by:

  • J. J. Thompson: Demonstrated that atoms are divisible and discovered the electron's mass to charge ratio (~1897).

  • Robert Milliken: Conducted oil drop experiments to determine the charge of the electron (~1909).

  • Ernest Rutherford: Known for the gold foil experiments leading to the nuclear model of the atom (~1911).

  • Frederick Soddy: Introduced the concept of isotopes (early 1900s).

  • James Chadwick: Provided evidence for the existence of neutrons (~1932).

Section 2.3 Atomic Structure and Symbolism

• Structure of the atom:

  • Nucleus: Comprised of protons (positive charge) and neutrons (no charge).

  • Electrons: Negatively charged particles surrounding the nucleus.

• Relative Masses and Charges of atomic particles (see Table 2.2, page 80, OpenStax text).

• Understanding Atomic Number (Z) and Mass Number (A):

  • Atomic Number (Z): The number of protons in the nucleus.

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

• Distinction between:

  • Atoms: Generally neutral entities having balanced charges.

  • Ions: Charged particles formed when atoms gain or lose electrons.

  • Cations: Positively charged ions (loss of electrons).

  • Anions: Negatively charged ions (gain of electrons).

• Recognition of element symbols in the periodic table and their corresponding elements.

• Isotopes: Variants of elements that have the same number of protons but different numbers of neutrons.

• Be proficient in calculating:

  • Average mass of atoms of elements.

  • Percent composition of samples of compounds.

Section 2.4 Chemical Formulas

• Recognition and understanding of various chemical formulas.

• Distinguish between formulas for:

  • Ionic Compounds: Compounds formed from ionic bonds (e.g., NaCl).

  • Molecular Compounds: Compounds formed from covalent bonds (e.g., H₂O).

• Isomers:

  • Structural Isomers: Different connectivity among the same atoms.

  • Geometric Isomers: Different spatial arrangements of atoms in a molecule.

Section 2.5 The Periodic Table

• Historical figures:

  • Dmitri Mendeleev: Developed an early version of the periodic table (1869).

  • Julius Lothar Meyer: Independently developed a periodic table (1870).

• Periodic Law: The properties of elements are periodic functions of their atomic numbers.

• Organization of the Periodic Table:

  • Elements arranged according to increasing atomic number.

  • Elements in the same period (row) share similar properties and energy levels.

  • Elements in the same group (column) have similar chemical behaviors.

• Metals vs Nonmetals:

  • Metals: Typically found on the left side, possess good conductivity and malleability.

  • Nonmetals: Found on the right, generally poor conductors and brittle.

• Identification of:

  • Alkali Metals: Group 1, highly reactive.

  • Alkaline Earth Metals: Group 2, reactive metals.

  • Chalcogens: Group 16, diverse group including oxygen and sulfur.

  • Halogens: Group 17, highly reactive nonmetals.

  • Noble Gases: Group 18, inert gases with low reactivity.

  • Transition Metals: Groups 3-12, metals with variable oxidation states.

  • Inner Transition Metals: Lanthanides and Actinides, known for radioactive properties.

• Recognize trends in Effective Nuclear Charge and Electronegativity across the table.

Section 2.6 Molecular and Ionic Compounds

• Understanding chemical reactions as a rearrangement of atoms and electrons.

• General rule: Atoms move towards achieving an octet configuration:

  • Some atoms become ions to form ionic bonds (transfer of electrons).

  • Some share electrons to establish covalent bonds.

• Ionic Compounds: Form from ionic bonds involving the transfer of electrons.

• Molecular Compounds: Form from covalent bonds involving shared electrons.

Section 2.7 Chemical Nomenclature (Naming)

• Distinction between:

  • Organic Compounds: Typically contain carbon and must follow specific naming conventions.

  • Inorganic Compounds: Do not predominantly consist of carbon.

• Naming rules include:

  • Binary compounds: Use prefixes to denote the number of atoms present.

  • Naming ionic compounds, including those with metals that have variable charges.

  • Naming ionic hydrates, molecular inorganic compounds, binary acids, and oxoacids.

• Strong Acids and Bases: Familiarity with the six common strong acids (e.g., HCl, H₂SO₄) and strong bases (e.g., NaOH, KOH)

  • Remember that stronger acids correspond to weaker conjugate bases.

Chapter 3: Composition of Substances and Solutions

Section 3.1 Formula Mass and the Mole Concept

• Determine the formula mass (molar mass) of:

  • Molecules, ionic compounds, polyatomic ions.

• Mole: The SI unit for chemical quantity.

  • 1 ext{ mol} = 6.02 imes 10^{23} ext{ atoms/molecules/etc.}.

  • Avogadro’s Number: 6.02214179 imes 10^{23}.

• Understand:

  • The correlation between the number of atoms in moles of different substances.

  • The mass of one mole of various substances differs.

• Recognize that Molar Mass is the mass (in grams) of one mole, expressed in g/mol, and is numerically equivalent to the atomic mass in atomic mass units (amu).

• Proficiency in conversions between mass, moles, number of molecules, and number of atoms.

Section 3.2 Determining Empirical and Molecular Formulas

• Ability to determine:

  • Percent composition from atomic masses.

  • Empirical formulas from grams of elements in samples of compounds (Example 3.9, 3.10, 3.11).

  • Convert empirical formulas to molecular formulas using molecular mass (Example 3.12, 3.13).

Section 3.3 Molarity

• Concentration: The amount of solute present in a given volume of solution.

  • Solute: The substance that is dissolved.

  • Solvent: The liquid that dissolves the solute.

  • Solution: The homogenous mixture of solute and solvent.

• Distinction between:

  • Concentrated and Dilute Solutions.

• Recognize Molarity (M) as:

  • M = rac{ ext{moles of solute}}{ ext{volume of solution in liters}}.

• Proficiency in calculating:

  • Molar Concentration (Example 3.14).

  • Moles and volumes based on molar concentrations (Example 3.15).

  • Molar concentration from solute mass (Example 3.16).

  • Mass of solute in given volumes (Example 3.17).

  • Volume of solution from given mass of solute (Example 3.18).

Section 3.4 Other Units for Solution Concentrations

• Become proficient at calculations involving:

  • Mass Percentage: (mass of solute / total mass of solution) × 100% (Examples 3.22, 3.23).

  • Volume Percentage: (volume of solute / total volume of solution) × 100% (Example 3.24).

  • Mass-Volume Percentage: (mass of solute/volume of solution) × 100% (Example 3.25).

  • Parts per Million (ppm) and Parts per Billion (ppb) as measures of concentration (Example 3.25).

Chapter 4: Stoichiometry of Chemical Reactions

Section 4.1 Writing and Balancing Chemical Equations

• Ability to write balanced chemical equations from reactants.

• Recognize:

  • Reactants: Substances that undergo change.

  • Products: Substances formed from the reaction.

  • Phase Labels: Indicate the physical state of the reactants and products.

  • Stoichiometric Coefficients: Numbers indicating the ratio of reactants to products.

• Balancing Equations: Ensure mass and charge are conserved in reactions.

• Write balanced:

  • Molecular Equations: Full formulas for all reactants and products.

  • Complete Ionic Equations: Show all ions present in solution.

  • Net Ionic Equations: Identify only the species that undergo change.

• Importance of net ionic equations in predicting outcomes of reactions.

Section 4.2 Classifying Chemical Reactions

• Distinctions between different reaction types:

  • Redox Reactions: Involves changes in oxidation states.

  • Precipitation Reactions: Formation of solid from a solution.

  • Acid-Base Reactions: Transfer of protons (H⁺ ions).

• Recognize whether a precipitate is formed from balanced equations.

• Differentiate between strong and weak acids/bases:

  • Neutralization reactions involve strong acids and bases reacting.

  • Use oxidation state assignments to determine redox occurrence.

  • Identify species oxidized and reduced in redox reactions.

• Recognize that single displacement and combustion reactions are types of redox reactions.

  • Distinction between decomposition and combination reactions based on redox characteristics.

Section 4.3 Reaction Stoichiometry

• Utilize stoichiometric coefficients to:

  • Determine moles of reactants needed (Example 4.8).

  • Relate masses of reactants to products (Example 4.10).

  • Relate masses between reactants themselves (Example 4.11).

Section 4.4 Reaction Yields

• Identify the Limiting Reactant:

  • Determines the maximum amount of product formed in a reaction (Example 4.12).

• Calculate Percent Yield:

  • ext{Percent Yield} = rac{ ext{Actual Yield}}{ ext{Theoretical Yield}} imes 100 (Example 4.13).

Section 4.5 Quantitative Chemical Analysis

• Techniques for analytical chemistry:

  • Titration: Method to determine concentration based on neutralization reactions.

  • Gravimetric Analysis: Measure of mass to determine concentration.

  • Combustion Analysis: Used for determining empirical formulas through combustion.

Chapter 5: Thermochemistry

Section 5.1 Energy Basics

• Define and identify examples of:

  • Kinetic Energy: Energy due to motion.

  • Potential Energy: Stored energy based on position.

  • Thermal Energy: Energy within a system due to temperature.

• Conversion between Temperature Scales:

  • Celsius (°C), Fahrenheit (°F), Kelvin (K).

• Units for Energy:

  • Calorie (cal): Amount of energy required to raise the temperature of 1 gram of water 1°C.

  • Joule (J): SI unit of energy (1 cal = 4.184 J).

• Distinction between:

  • Heat Capacity: Total heat required to change the temperature by a degree.

  • Specific Heat: Heat needed to raise the temperature of 1 gram by 1°C.

  • Molar Heat Capacity: Amount of heat needed for one mole.

• Calculate:

  • C = rac{q}{ riangle T} and q = c imes m imes riangle T: where q = heat energy, c = specific heat, m = mass, ΔT = change in temperature.

Section 5.2 Calorimetry

• Describe how Calorimetry measures heat transfer.

• Distinctions between:

  • Endothermic Processes: Absorb heat.

  • Exothermic Processes: Release heat.

  • System vs Surroundings: The system is the part of the universe we focus on.

• Apply the relationship:

  • q{ ext{reaction}} = -q{ ext{solution}}

• Understand the use of bomb calorimeter for constant volume calorimetry.

• Nutritional calorie (Calorie) is defined as:

  • 1 ext{kcal} = 1,000 ext{cal}.

Section 5.3 Enthalpy

• State First Law of Thermodynamics:

  • Energy cannot be created or destroyed, only transformed.

• Concepts and terminology:

  • Internal Energy: Total energy contained within a system.

  • Expansion Work: Work done by a system when it expands against external pressure.

  • State Function: Property determined by state variables (e.g., internal energy, enthalpy).

  • Enthalpy (H) and Change in Enthalpy (ΔH): The heat content of a system at constant pressure.

  • Standard State Conditions: Standard for measuring enthalpy changes at 1 atm and specified temperatures (usually 25°C).

  • Standard Enthalpy of Combustion (ΔH°ₕ): Heat released during combustion of 1 mole of substance.

  • Standard Enthalpy of Formation (ΔH°ₓ): Heat change when one mole of a compound is formed from its elements.

  • Hess’s Law: Total enthalpy change for a reaction is the sum of the enthalpy changes for individual steps.

• Write thermochemical equations while manipulating them for calculations.

• Determine heat of reactions from standard enthalpies of formation using Hess's law.

Chapter 6: Electronic Structure and Periodic Trends

Section 6.1 Electromagnetic Energy

• Convert between wavelength and frequency:

  • Use the equation c =
    u imes ext{λ} where c = speed of light, ν = frequency, and λ = wavelength.

• Understand relationships between:

  • Energy: Directly proportional to frequency and inversely proportional to wavelength.

• Familiarize yourself with the electromagnetic spectrum: Range of all wavelengths of electromagnetic radiation.

• Recognize and analyze Constructive and Destructive Interference: Wave behaviors affecting intensity and amplitude.

• Identify nodes in standing waves indicating points of no displacement.

• Distinguish between Continuous Spectrum: All wavelengths vs Line Spectrum: Specific wavelengths emitted by substances.

• Discuss phenomena such as:

  • Ultraviolet Crisis: Inadequacy of classical physics for electron behavior.

  • Photoelectric Effect: Emission of electrons when light hits a material.

  • Wave-Particle Duality of Light: Light behavior exhibiting both wave-like and particle-like properties.

• Contributions of early scientists:

  • Balmer: Developed formula for hydrogen spectral lines.

  • Rydberg: Expanded understanding of spectral series.

  • Bohr: Introduced quantized energy levels in star model of the hydrogen atom.

Section 6.2 The Bohr Model

• Describe Bohr Model of the atom, focusing on discrete energy levels of electrons.

• Distinction between Ground State (lowest energy configuration) and Excited State (higher energy configurations).

• Utilize the Rydberg Equation for calculating transition energies of electrons between levels:

  • rac{1}{ ext{λ}} = RH imes igg( rac{1}{n1^2}- rac{1}{n_2^2}igg).

Section 6.3 Development of Quantum Theory

• Relate wave-particle duality for both light and electrons.

• Understand and apply Quantum Numbers:

  • Define each of the four quantum numbers:

  • Principal Quantum Number (n): Indicates energy level.

  • Angular Momentum Quantum Number (l): Specifies subshell shape.

  • Magnetic Quantum Number (m): Orientation of orbital.

  • Spin Quantum Number (ms): Direction of electron spin.

• Define deBroglie Wavelength: Indicates wave nature of particles:

  • ext{λ} = rac{h}{mv}, where h = Planck's constant, m = mass, v = velocity.

• Understand the Heisenberg Uncertainty Principle: Limits on simultaneously knowing position and momentum of particles.

• Comprehend Pauli Exclusion Principle: No two electrons in an atom can have the same four quantum numbers.

Section 6.4 Electronic Structure of Atoms

• Write electron configurations for both atoms and ions based on their energy levels.

• Draw Orbital Diagrams illustrating electron arrangement in different subshells.

• Apply principles of:

  • Aufbau Principle: Electrons fill lower-energy orbitals first.

  • Hund’s Rule: Electrons occupy degenerate orbitals singly before pairing.

• Identify Valence Electrons (electrons in the outermost shell) versus Core Electrons (inner shell electrons).

• Classify:

  • Main Group Elements: Elements in Groups 1, 2 and 13-18.

  • Transition Elements: Elements in Groups 3-12.

  • Inner Transition Elements: Lanthanides and Actinides.

Section 6.5 Periodic Variations in Element Properties

• Describe and explain trends within the periodic table:

  • Covalent Radius: Decreases across a period, increases down a group.

  • Effective Nuclear Charge (Z_eff): Increases across a period due to greater nuclear attraction.

  • Ionic Radii: Varies based on charge and environment.

  • Ionization Energy: Energy required to remove an electron.

  • Electron Affinity: Energy change when an electron is added.

  • Metallic Character: Increases down a group, decreases across a period.

• Recognize Isoelectronic Species: Different elements or ions having the same electron configuration.

Chapter 7: Chemical Bonding and Molecular Geometry

Section 7.1 Ionic Bonding

• Distinguish between Cations (positively charged ions) and Anions (negatively charged ions), atoms, and molecules.

• Explain the formation processes of cations and anions:

  • Cations form through loss of electrons (often metals).

  • Anions form through gain of electrons (often nonmetals).

• Identify properties of ionic compounds:

  • High melting and boiling points, conduct electricity in solution.

• Compare ions and atoms based on properties such as size and charge.

• Identify and describe Binary Ionic Compounds: Combustions of two elements.

• Recognize that attractive forces between ions are isotropic: Equal in all directions.

• Explain the Inert Pair Effect: Tendency of the two electrons of the outermost s subshell to remain unbonded.

• Write electron configurations for:

  • Atoms.

  • Anions (gaining electrons).

  • Cations (losing electrons).

Section 7.2 Covalent Bonding

• Explain the distinctions between covalent and ionic bonds:

  • Covalent bonds form when atoms share electrons (typically nonmetals).

  • Ionic bonds result from the transfer of electrons (typically between metals and nonmetals).

• Compare properties of molecular compounds with ionic compounds:

  • Molecular compounds typically have lower melting/boiling points.

• Discuss how covalent bonds form based on electron sharing.

• Recognize factors affecting bond lengths:

  • Electronegativity, number of shared electron pairs.

• Understand that:

  • Bond Breaking: Endothermic process (energy absorbed).

  • Bond Making: Exothermic process (energy released).

• Distinguish between:

  • Polar Covalent Bond: Unequal sharing of electrons.

  • Pure Covalent Bond: Equal sharing of electrons (e.g., in diatomic molecules).

• Differentiate properties of Effective Nuclear Charge, Electronegativity, and Electron Affinity.

• Recognize trends in Effective Nuclear Charge, Electronegativity, and Electron Affinity across the periodic table.

• Identify Polyatomic Ions: Groups of covalently bonded atoms carrying a charge.

Section 7.3 Lewis Symbols and Structures

• Write Lewis Symbols based on valence shell electron configurations for atoms.

• Show electron arrangements in ions using Lewis symbols, distinguishing between lone pair and bonding electrons.

• Explain the Octet Rule: Atoms tend to bond until they have eight electrons in their outermost shell.

• Distinction between:

  • Single Bonds, Double Bonds, Triple Bonds: Represent different degrees of electron sharing.

• Draw Lewis Structures by following valence shell arrangements.

• Recognize exceptions to the octet rule:

  • Electron Deficient Species: Molecules with fewer than eight electrons around an atom.

  • Hypervalent Species: Molecules where central atoms hold more than eight electrons.

Section 7.4 Formal Charge and Resonance

• Assign Formal Charge to atoms in molecules to determine stability.

• Use formal charge to evaluate which Lewis structures are more favorable.

• Explain Resonance: Configuration where two or more valid structures exist for a molecule, not representing real distribution but stability.

• Distinguish resonance structures from isomers (different connectivity of atoms).

• Define Resonance Hybrid: The actual molecule represented as a blend of all possible resonance structures.

Chapter 8: Advanced Theories of Covalent Bonding

Section 8.1 Valence Bond Theory

• Describe the concept of Covalent Bond Formation:

  • Formation through overlap of atomic orbitals, with electrons pairing.

• Conditions for bond formation include:

  • Proper orientation and sufficient energy during overlap.

• Discuss the energy changes occurring when forming bonds

  • Closer to stability with bond formation.

• Distinguish between types of bonds:

  • Sigma Bonds: Formed by head-on overlap of orbitals.

  • Pi Bonds: Formed by side-on overlap, occurring in double/triple bonds.

• Determine the number of sigma and pi bonds in molecular structures.

Section 8.2 Hybrid Atomic Orbitals

• Explain limitations of valence bond theory in describes carbon bonding.

• Define Hybridization: The process in which atomic orbitals mix to form new hybrid orbitals suitable for bonding.

• Describe hybrid orbital formations for:

  • sp³ Hybridization: One s and three p orbitals (e.g., in methane).

  • sp² Hybridization: One s and two p orbitals.

  • sp Hybridization: One s and one p orbital.

• Determine the hybridization of a carbon atom based on its bonding structure.

Section 8.3 Multiple Bonds

• Describe covalent bonds in terms of Orbital Overlap where multiple bonds involve both sigma and pi bonds.

• Relate Resonance to electron delocalization through pi bonds in structures.

• Draw atomic orbital energy diagrams displaying formation of sp³, sp², and sp hybrid orbitals.

• Draw and label Orbital Overlap in molecules such as Ethene, Benzene, and Ethyne.

Section 8.4 Molecular Orbital Theory

• Characteristics explained by Molecular Orbital Theory include:

  • This theory elucidates the nature of bonding and accounts for paramagnetic and diamagnetic properties of substances.

• Compare and contrast Valence Bond Theory with Molecular Orbital Theory outlined in Table 8.2 (Page 433).

• Distinguish between Bonding and Antibonding Orbitals, and sigma and pi interactions.

• Recognize Molecular Orbital Energy Diagrams for diatomic molecules to determine stability and bond order:

  • Bond Order calculated by:

  • ext{Bond Order} = rac{ ext{(number of bonding electrons - number of antibonding electrons)}}{2}.

Chapter 9: Gases

Section 9.1 Gas Pressure

• Define Gas Pressure: Force per unit area exerted by gas on walls of its container.

  • Pressure usually measured in pascals (Pa).

• Introduce the Ideal Gas Law:

  • PV = nRT,

  • where P is pressure, V is volume, n is number of moles, R is the ideal gas constant, and T is temperature in Kelvin.

Section 9.3 Stoichiometry of Gaseous Substances, Mixtures, and Reactions

• Use the Ideal Gas Law to derive:

  • Gas Density: density can be expressed based on molar mass and conditions of gas.

  • Molar Mass: Can be calculated from gas density and the ideal gas law.

  • Molecular Formula: Determined from experimental mass data using molar relationships.

Section 9.5 Interpret Molecular Speed Diagrams

• Discuss interpretation of molecular speed diagrams, indicating the distribution of molecular speeds in gas samples (see Figure 9.33 in OpenStax text).

Chapter 10: Liquids and Solids

Intermolecular Forces

• Describe various types of interactions present in liquids and solids:

  • Dispersion Forces: Weak forces arising from temporary shifts in electron density.

  • Dipole-Dipole Interactions: Occur between polar molecules due to their permanent dipoles.

  • Hydrogen Bonding Interactions: Specifically strong dipole-dipole interactions that occur when hydrogen is bonded to electronegative atoms (e.g., N, O, F).

• Relate intermolecular forces to physical properties including:

  • Melting Point: The temperature at which a solid becomes a liquid.

  • Boiling Point: The temperature at which a liquid becomes a gas.