Types of Chemical Reactions

• Chemical reactions undergo several classifications

Combustion Reactions

• A reactant combines with oxygen, releasing energy in the form of light and heat
• An example is 4Fe(s) + 3O2(g) → 2Fe2O3(s) + heat

Precipitation Reactions

• Cations and anions in an aqueous solution combine, creating an insoluble ionic solid (precipitate)
• An example is 2KI(aq) + Pb(NO3)2(aq) → 2KNO3(aq) + PbI2(s)
• Precipitate reactions hold importance in analytical chemistry

Acid-Base Reactions

• Acids and bases mix
• Reactions form water and a salt (an ionic compound consisting of a base's cation and an acid's anion)
• Example: HCI(aq) + NaOH(aq) → H2O(l) + NaCl(aq)

Oxidation-Reduction (Redox) Reactions

• Reactions entail the transfer of electrons between two species
• Oxidation is Losing electrons (OIL)
• Reduction is Gaining electrons (RIG)
• An example is Zn + Cu2+ → Zn2+ + Cu

Solubility and Precipitation

• Solubility refers to the maximum solute concentration dissolving in a solvent at a given temperature
• Soluble substances exhibit relatively high solubility and do not form a precipitate unless supersaturated
• Insoluble substances possess low solubility and readily precipitate from solution
• Solubility tables aid in predicting if ionic compounds dissolve in water or form precipitates

Generally Soluble Compounds

• Compounds containing alkali metals (Li+, Na+, K+, etc.) and NH4+ are generally soluble
• Nitrates (NO3-), acetates (C2H3O2-), etc. are also generally soluble
• Most chlorides, bromides, and iodides are generally soluble, except when combined with Ag+, Hg22+, Pb2+
• Most sulfates are generally soluble, except when combined with Sr2+, Ba2+, Pb2+, Ag+, Ca2+

Generally Insoluble Compounds

• Hydroxides (OH-) are generally insoluble, except when combined with alkali metals and NH4+
• Sulfides (S2-) are generally insoluble, except when combined with alkali metals, NH4+, and alkaline earth metals
• Carbonates (CO32-), phosphates (PO43-), etc. are generally insoluble, except when combined with alkali metals and NH4+

Ionic Equations

• Ionic reactions utilize three types of equations

Molecular Equation

• Displays all reactants and products as neutral substances
• Example: CaCl2(aq) + 2AgNO3(aq) → Ca(NO3)2(aq) + 2AgCl(s)

Complete Ionic Equation

• Displays all dissolved substances as their component ions
• Example: Ca2+(aq) + 2Cl-(aq) + 2Ag+(aq) + 2NO3-(aq) → Ca2+(aq) + 2NO3-(aq) + 2AgCl(s)

Net Ionic Equation

• Shows only the species directly participating in the reaction, excluding spectator ions
• Example: Cl(aq) + Ag+(aq) → AgCl(s)
• Spectator ions are those that remain unchanged on both sides of the reaction

Acids and Bases: Arrhenius Definitions

• Acid: A substance that produces H+ ions (protons) in an aqueous solution
• Base: A substance that produces OH- ions in an aqueous solution

Strong vs. Weak Acids

• Strong acids completely ionize in solution (HCl, HBr, HI, HNO3, H2SO4, HClO4)
• Weak acids only partially ionize in solution (HCOOH, CH3COOH, HF)

Strong vs. Weak Bases

• Strong bases ionize completely in aqueous solution (NaOH, KOH, LiOH, Ca(OH)2, etc.)
• Weak bases partially ionize in aqueous solution (NH3, CH3NH2, etc.)

Oxidation Numbers

• Rules for assigning oxidation numbers include:
• Elemental substances: 0
• Monoatomic ions: Equal to their charge
• Sum of oxidation states in a neutral molecule: 0
• Sum of oxidation states in an ion: Equal to its charge
• Metals in compounds: Group 1A: +1, Group 2A: +2
• Nonmetals: H: +1, O: -2, F: -1, others vary

Stoichiometry and Reaction Calculations

• Stoichiometry involves the relationships between the amounts (moles) of reactants and products in a chemical reaction
• Limiting Reactant and Theoretical Yield

Limiting Reactant

• The reactant that produces the least amount of product (completely consumed)

Reactant in Excess

• The reactant present in greater quantity than needed (has leftover)

Theoretical Yield

• The maximum amount of product that can be made based on the limiting reactant

Actual Yield

• The amount of product actually produced, typically less than the theoretical yield

Percent Yield

• (Actual yield / Theoretical yield) × 100%

Solution Stoichiometry

• Utilizes molarity (M = moles of solute / liters of solution) to calculate amounts
• Relevant in titrations and other quantitative analyses

Analytical Chemistry Techniques: Titration

• Involves a titrant, which is a solution of known concentration, and an analyte, which is a solution of unknown concentration
• The equivalence point is when chemically equivalent quantities of reactants have been mixed
• The end point is observed using indicators to show when the equivalence point has been reached

Gravimetric Analysis

• Determines the quantity of an analyte based on its mass
• Often involves precipitation, followed by filtration, drying, and weighing of the precipitate

Introduction to Thermochemistry

• Thermochemistry is the study of heat absorbed or released during chemical and physical changes
• Thermochemistry focuses on how chemical reactions exchange energy with their surroundings and how to quantify these energy exchanges
• A practical example is the flameless ration heater (FRH) used in military MREs, where Mg + 2H2O → Mg(OH)2 + H2 + heat
• The flameless ration heater is an exothermic reaction that releases 440.9 kJ/mol of energy as heat

Key Concepts in Thermochemistry

• Energy (E) signifies the capacity to supply heat or do work and is measured in joules (J) or kilojoules (kJ)

Forms of Energy

• Kinetic Energy is energy due to motion -Potential Energy is energy due to position or composition -Thermal Energy is energy associated with temperature -Chemical Energy is energy stored in chemical bonds
• Heat (q) is the transfer of thermal energy between two bodies, flowing from high to low temperature
• Heat is measured in joules (J), kilojoules (kJ), or calories (cal)
• +q indicates heat is absorbed (endothermic process)
• -q indicates heat is released (exothermic process)
• Work (w) represents the result of a force acting over a distance, where work = force × distance = F × d
• Temperature (T) measures the thermal energy within a sample in °C or K
• Higher temperature signifies faster moving molecules

Systems and Surroundings

• System refers to the material or process being investigated
• Surroundings include everything with which the system can exchange energy

Types of Systems

• Open systems can exchange both mass and energy with surroundings -Closed systems can exchange energy but not mass -Isolated systems cannot exchange energy or mass

First Law of Thermodynamics

• The law of conservation of energy states that energy can be transformed but cannot be created or destroyed
• The total energy of an isolated system remains constant

Endothermic and Exothermic Reactions

• Endothermic reactions (+ΔH) absorb energy/heat from the surroundings, resulting in products having higher energy than reactants
• Exothermic reactions (-ΔH) release energy/heat to the surroundings, resulting in reactants having higher energy than products

Specific Heat Capacity

• The specific heat capacity (c1) is the amount of heat required to raise the temperature of 1 gram of a substance by 1°C
• Formula: q = m × c₁ × ΔT, where:
• q = amount of heat (J)
• m = mass (g) -C₁ = specific heat capacity [J/(g·°C)] -ΔT = temperature change (°C)
• Water has a high specific heat capacity (4.18 J/g·°C), which means it requires more energy to heat up and retains heat longer than many other substances

Calorimetry

• Calorimetry is the process of measuring heat released or absorbed during chemical reactions using a calorimeter
• There are two main types:

Coffee Cup Calorimetry

• Used for measuring heat changes in solution reactions
• System: the chemical reaction -Surroundings: the solution (usually water) -q() = m(。) × c() × ΔΤ -q(rx) = -qo□□) (energy is conserved)

Bomb Calorimetry

• Bomb Calorimetry is used for measuring heat changes in combustion reactions
• The reaction occurs in a sealed container

Enthalpy (H)

• Enthalpy is a measure of the total heat content of a system
• Enthalpy change (ΔH) is the heat absorbed or released at constant pressure

Three methods to determine ΔH

• Experimentally through calorimetry
• Hess's law
• Using standard enthalpies of formation

Hess's Law

• Hess's law states that if a chemical reaction can be expressed as the sum of several steps, then ΔH for the overall reaction is the sum of the enthalpy changes for each step

Three key relationships

• If a reaction is multiplied by a factor, ΔH is also multiplied by that same factor
• If a reaction is reversed, the sign of ∆H is changed
• Enthalpy changes are additive for sequential reactions

Standard Enthalpy of Formation (ΔH°f)

• The enthalpy change when 1 mole of a compound forms from its elements in their standard states
• Standard state: Reference conditions (typically 1 atm pressure, 25°C)
• For elements in their standard state: AH°f = 0
• Formula to calculate enthalpy of reaction: ΔΗ°rxn = ∑(n□ΔΗ°f products) - ∑(nr∆Η°f reactants)

Chapter 6: Electronic Structure and Periodic Properties

Wave Nature of Light

Light as a Wave

• Electromagnetic radiation is electromagnetic radiation consisting of oscillating electric and magnetic fields
• All electromagnetic waves travel at the speed of light: 3.00 × 108 m/s

Properties of Waves

-Wavelength (λ) is the distance between consecutive peaks -Frequency (v) is the number of waves passing a point per second (Hz or s-1) -Amplitude is half the height from peak to trough, which determines brightness -Relationship: v = c/λ, where c is the speed of light

• Electromagnetic Spectrum: Range from gamma rays (shortest λ) to radio waves (longest λ)
• Visible light is just a small portion of the spectrum

Energy of Light

• Energy of a photon: E = hv = hc/λ
• h is Planck's constant: 6.626 × 10-34 J·s
• Higher frequency (shorter wavelength) = higher energy

Particle-Wave Duality

• Light and other quantum entities can behave as both particles and waves
• Light behaves as waves (interference, diffraction)
• Light also behaves as particles (photons) in the photoelectric effect

Photoelectric Effect

• Photoelectric Effect ejects, electrons when light shines on certain metals if:
• The frequency is greater than the threshold frequency
• Each photon has enough energy to overcome the binding energy (Φ)
• Key equation: KE = hv - Φ
• If hv < Φ: No electrons are ejected
• If hv = Φ: Electrons are just ejected with zero KE
• If hv > Φ: Electrons are ejected with KE = hv - Φ

Atomic Spectra

• When atoms are excited, they emit light at specific wavelengths, creating an emission spectrum unique to each element
• Bohr's Model of the Atom

Niels Bohr proposed

-Electrons travel in fixed circular orbits around the nucleus -These orbits exist only at specific distances from the nucleus -Each orbit has a fixed, quantized energy -When electrons move between orbits, they absorb or emit photons with specific energies

• For hydrogen, the energy of an electron in the nth orbit is: ΔΕ = -2.18 × 10-18 J × (1/nf² - 1/ni²)

Wave Behavior of Matter

• De Broglie proposed that all matter has wave properties
• Wavelength of a particle: x = h/(mxv)
• This is especially significant for very small particles like electrons
• For large objects, the wavelength is too small to be observed

Quantum Mechanics and Orbitals

• In modern quantum mechanics:
• Electrons are described by wave functions (ψ)
• The probability of finding an electron in a particular location is given by |Ψ|²
• Orbitals are three-dimensional probability distributions where electrons are likely to be found
• Orbitals are defined by quantum numbers