Class 1 - 06/06/24:

• Atoms - smallest unit of any element - has protons, neutrons, and electrons

  • p = +1, mass = 1amu

  • e = -1, mass = 0amu

  • n = 0, mass = 1amu

• Atomic number = Z, the number of p*

• Mass number = p*+n

• The charge = p-e

  • Cations and anions are ions

    • C>0 = cation → +

    • C<0 = anion → -

    • C=0 = atom level

• Isotopes: two atoms of the same element that differ in their number of neutrons - determined by mass number

• Bohr model of the atom:

  • Electrons orbit at a fixed distance from the nucleus - the orbit decreases with distance from the nucleus - as we move away, the ends come closer and energy increases with distance from the nucleus

• Electrons absorb only specific allowed E(due to fixed quantities of E)

  • Current orbit = ground state

  • Higher E orbit = excited state

  • Ephoton = Ef-Ei

  • e- in an excited state can come to a lower level to emit a photon - when dropped, the e- becomes relaxed

• Hydrogen Absorption spectrum = dark bands on a bright background(absorb, so black line)

• Hydrogen Emission spectrum = bright bands on a dark background(emit, so bright lines)

• The Energy of a photon is related to its wavelength(λ lambda) and frequency(f)

  • E = hf = hc/λ → Wavelength and frequency are inversely related - when f is high, λ is low, E is high; when f is low, λ is high, E is low

    

• E- exists in 3D orbitals and 4 quantum number describe their structures(orbitals are the areas around the nucleus where an e- is most likely to be found)

  • s, p, d, f → s being lowest in E and F in highest E

  • ex. boron= 1s2 2s2 2p1 → goes by block and row of the periodic table

• 3 basic rules for Electron filling:

  • Pauli principle: there can be no more than 2 e- in any given orbital(spin up and down)

    • e- cannot be the same, hence, different spins

  • Aufbau principle: E- occupies the lowest orbitals first and is filled in increasing E

    • exception is 3d and 4s → 4s are removed before 3d

  • Hund’s Rule: E- first occupy an orbital singly then pair up(no orbital left empty)

• Anomalous electron configurations: when some elements prefer to be half-filled or filled by taking an e from 4s and putting it into 3d ex. Cr and Cu

• Paramagnetic = at least one unpaired e-

• Diamagnetic = all e- are paired ex. noble gases

• Ground state e-configuration: the ground state is the lowest e configuration → correct amount of e as the element has

• Excited state e-configuration: the element has jumped an orbital but it has not changed the amount of electrons as it originally had

• A half-filled shell is more stable than one that isn’t; filled one is the most stable ex. noble gases

• Valence shell configurations of elements determine the chemical reactivity of the elements → Elements in the same group show similar characteristics ex. noble gases as calm due to their octet shells, while halogens are reactive gases with 1 e-missing

• Valence shell electrons experience electrostatic attraction due to the nucleus → shielding effect

  • force of electrostatic attraction is proportional to Zeff + C/r²

• Atomic radius increases going down right to left

• Ionization E increase going up from left to right

• Electron affinity(negativity) increases going up from left to right

• Electronegativity increases going up from left to right

  • FON=ClBrISC=H

• Acidity increases going down left to right down

Class 2: 11/06/24:

• Molecular Structure

  • Drawing Lewis structures:

    • count the number of valence e-

    • Arrange atoms with the least e-neg in the center(C is always in the middle and H is never in the middle) - Use FON=ClBrISC=H

    • Connect each outer atom to the center(each line is two electrons)

    • add dots of pairs to each electron till none are remaining

    • Complete missing octets

    • Assign formal charges

      • FC = valence e -(1/2 bonding e - sticks)-(lone pair e - dots)

    • Always check the number of e, octet rule obeyed, formal charges add up to total charge, the smallest set of FC, + charges on fewer e-neg atoms and - charges on more e-neg atoms

  • Sets of electrons also determine hybridization → count each bond and lone pair as one group

    • 2 groups = sp linear ex. CO2

    • 3 groups = sp² trigonal planar

    • 4 groups = sp³ tetrahedral

  • The lone pairs on the molecule determine the molecular shape

    • AX4 = tetrahedral ex. CH4

    • AX3E = trigonal pyramidal ex. NH3

    • AX2E2 = bent ex. H2O

      • All have the same bond angles but different shapes

• Chemical bonding: these bonds form when electrons are shared between two atoms as their orbitals overlap

  • More electrons shares make a stronger bond

  • A shorter distance between atoms makes a stronger bond

  • Stronger bonds have higher dissociation energies

  • Breaking bonds is always endothermic

  • Ionic or covalent bond types

  • Electronegativity determines the bond → if higher EN, then the bond is polar and e are not shared equally; if small EN, then the bond is shared equally and non-polar

    • Covalent bonds: formed between non-metal-non-metal → high EN - these compounds are insulators

    • Metallic bonds are formed between atoms with low EN metal-metal - these compounds are conductors and malleable

    • Coordinate covalent bonds are formed between atoms with lone pairs and e-deficient - ligands ex. hemoglobin

    • Ionic bonds are formed with particles of opposite charges anions-cations - these are insulators and brittle

• Intermolecular Forces: when opposite charges attract each other - larger charges, stronger IMF and more tightly held together

  • Ion-dipole forces - between ions and polar molecules

  • Dipole-Dipole forces - between polar molecules and easier to cleave

  • Dipole-induced-dipole - between polar and non-polar - bigger e cloud and polarity from polar molecule but very easily cleaved

  • London Dispersion - temporary small dipoles formed by collisions - very weak

  • Hydrogen bonding - between very polar molecules with donors being (N-H, O-H, F-H) and acceptors being (O:, N:, F:)

    • Strongest bonds highest to lowest: Ionic>Covalent>Coordinate Covalent>Metallic

    • Strongest IMFs: Ion-dipoles>H-bond>Di-di>di-induced-di>LDF

• Chemical Thermodynamics

  • Enthalpy(H) → The energy stored within chemical bonds or any attractive forces

  • the change in enthalpy of a reaction is the difference between energy stored in reactant vs products - delta H reaction can be high or low

    • Low H is exothermic(-) Forming bonds

    • High H is endothermic(+)Breaking bonds

  • Enthalpies of formation is the amount of energy associated with forming One mole of compound = sum of products-sum of reactants

    • Enthalpy is independent of its pathways - Hess’s law requires a combination of two or more reactions to find the enthalpy for the overall reaction

  • Entropy(S) = potential randomness - more particles, changing phases, increase temperature → all increases entropy

  • Gibbs free energy: energy available to do work

    • G = H - TS

      • Spontaneous process is exergonic → G is negative

      • non-spontaneous is endergonic → G is positive

Class 3 - 27/06/24:

• Phases transitions → solid, liquid, gas, ideal gas

  • The IMFs decrease as the substance becomes more ‘liquidy’

  • Heat is absorbed going to gas, and released going to solid

• Triple Point: the point where all 3 phases coexist

• critical point: where the difference between liquid and gas is no longer distinct

• Calorimetry: the science of measuring the changes to determine heat transfer - using a calorimeter

• Heat of transition: the amount of E to complete a transition

• heat of fusion: the amount of heat to be absorbed to change from solid to liquid

• Heat of vapourization: teh amount of heat to be absorbed to change from liquid to gas

• ΔH = molar enthalpy of change(kJ/mol)

• n = number of moles of substance

• q = mcΔT → adding heat can also raise the temp of a substance by increasing the kinetic E

• Density: a measure of how condensed a substance is

  • p = m/v

  • IMF is inversely proportional to p

• Boiling point: the temperature at which the condensation/vapourization phase transition happens

• Melting and freezing point: the temperature at which fusion and crystallization occurs

• Solution: a homogenous mixture of two or more substances → solvent is usually water

• Electrolyte: a solute that is dissolved

• van’T Hoff Factor: the number of articles produced in a solution per mole of a substance

• electrolytes dissolve in water

• polar non-electrolytes dissolve in water

• non-polar non-electrolytes do not dissolve in water

  • Like-dissolves-like

• Solubility: the amount of a substance that can dissolve in a specific solvent at a specific temperature

  • Unsaturated - too cold → still can add more solute

  • Saturated - just right → solvent=solute

  • Supersaturated → too hot → too much solute, produces a precipitate

• Electrolytes in WATER:

• Ideal gas: has ZERO IMFs - particle size is negligible volume compared to container size

  • Has Ek proportional to temp

  • ideality favoured with high temp and low pressure

• Avogadro’s law: the volume of an ideal gas is proportional to the number of particles in the container at the given time, regardless of the identity of the gas

• STP = 0°C and 1 atm

• Two variable Gas Law’s

  • Boyle’s Law: P is inversely related to V → p1v1 = p2v2

  • Charles Law: t is directly proportional to V → v1/t1 = v2/t2

  • Gay-Lussacs Law: P is directly proportional to T → p1/t1 = p2/t2

• Pressure: 1atm = 1000mL = 1000 cm³ = 0.001 m³

• Temperature must be in the absolute temperature scale = Kelvin → °C + 273 = K

• Combined Gas Law → Boyle’s and Charle’s

  • p1v1/t = p2v2/t

• Ideal Gas Law: PV = nRT

  • R = gas constant(0.08 Lxatm/Kxmol)

• Ideal pressure will always be greater than real pressure(since real gas do experience IMFs)

• Ideal volume will have more free space in a container than a real gas

• Dalton’s Law: states that the total pressure of a mixture of gases is equal to the sum of their partial pressures

• Graham’s Law of Diffusion/Effusion: the rate of diffuse/effusion of a gas is inversely related to the square root of its molar mass

  • Rate gas 1/rate gas 2 = √(molar mass gas2/molar mass gas1)

  • Heavy particles move slowly and light particles move quickly

• At 1 mole He gas, takes up 22.4L of volume at ideal temperatures

• P vs. T = direct relationship → higher pressure, higher temperature

• P vs. V = indirect relation → higher pressure, lower volume

• V vs. T = direct relation → higher volume, higher temperature

• P vs. n = direct relation → higher pressure, moles increase

• V vs. n = direct relation → as volume increases, moles increase

Class 4 - Kinetics and Equilibrium

• Reaction Coordinate graph

• Ea is the difference between the reactant E and the highest E transition

• Highest Peak = slowest step(rate-limiting step)

• Reaction rate = change in concentration of a reactant or product over a change in time

  • rate = -1/r x [R]/t = + 1/p x [P]/t → reactants will be (-) and products positive since we gain them

  • The rate of the reaction will always have a positive rate in the forward direction

  • rR → pP

• Anything with the same coefficients will have the same rate of change

• Rate of reaction = to the compound in the stoichiometric equation with a 1 in front of it

• Collision Theory: reactants must collide at a faster rate

  • Increase temp

  • increase pressure, lower volume

  • increase concentration of reactants

  • adjusting molecules to a proper orientation by the addition of a catalyst → quicker reaction

  • cause activation E to change → lower it by adding a catalyst or increasing temperature

• Rate constant = the more successful the reaction, the higher the K value

  • K is inversely proportional to the Ea → Ea higher, lower K

  • K is directly proportional to the temperature

• Reaction Mechanism:

  • Catalysts

  • Intermediate

  • Transition states

• Rate Laws: give rate in terms of initial [C] and rate constant for the process

  • Rate = k[R]^r rR → pP

  • Only for elementary steps and not the overall reaction and solids and solvents not included

• Write out rate law, make sure powers of 10 are the same, and make sure the order of two different trials is the same

  • 0th - no collision - different [C] but same rate overall

  • 1st - one collision - double [C], doubles rate

  • 2nd - anything else that doesn’t match

  • rate constant can be found from a single trial

• Dynamic Equilibrium = forward and reverse rates are equal - [C] will be consistent at equilibrium

  • K = [P]^p/[R]^r → K can only be calculated at equilibrium

• Any K value will always be at equilibrium → K is unitless

  • K >1 → More products

  • K < 1 → More reactant

  • K = 1 → Both P and R are the same

• Kc = [C]

• Kp = pressure

• Kb = dissociations

• Ka = acids

• Kb = base

• Kf = formation

• Ksp = solubility

• Reaction Quotients → described the distance from Equilibrium → ratio of initial product/reactant [C]

  • Q > K → more products, the reaction goes in the reverse direction

  • Q = K → at equilibrium

  • Q < K → more reactants, the reaction goes in the forward direction

• Delta G = G° + RT lnQ

  • Q>K → reverse → delta G is + → non-spontaneous in forward, but spontaneous in reverse

  • Q<K → forward → delta G is - → spontaneous in forward direction

  • Delta G = 0 → Q=K at equilibrium

• delta G° = -RT lnK

  • K»1 → more products, delta G° becomes negative

  • K«1 → more reactants, delta G° becomes positive

  • K = 1 → delta G° is 0

• Le Châtelier’s Principle

  • a system in equilibrium, will shift to decrease stress when stressed

  • Solids and liquids don’t affect equilibrium

• Increase V, less pressure → shift to more moles of gas side

• Decrease V, increase pressure → shift to less gas side

• Treat Temperature like a product/reactant

  • If Endothermic, heat is → used up → reactant

    • + delta H

  • If Exothermic, heat is → produced → product

    • - delta H

  • Temp change Q and K → only way to change K value(applies to all K values)

• Multiple Equilibria → Some things can change other reactions(chain reaction)

• At equilibrium, forward = reverse

• Formation Equilibrium → equilibrium of a coordination complex

  • Reversing equilibrium

  • Combining equilibrium means multiplying the K

• Solubility Equilibrium: dissolving salts into water → cationic and anionic parts

  • Ksp(solubility) → use of ions only that are aqueous

  • Molar solubility = mol solute/L solution

    • More molar solubility, more solubility of the salt

• ICE Tables

• Qsp = [Ion A][Ion B]…

  • Qsp > Ksp → supersaturated → can form precipitate

  • Qsp = Ksp → saturated = equilibrium

  • Qsp < Ksp → unsaturated

• Common Ion effect: salt’s solubility decreases if a common ion is added to the solution

  • If add an acid, acid/base reaction could happen → solubility will increase if common ion removed

• Kd = dissociation

• Kaff → affinity

  • putting it together, reaction reversed → inverse of the dissociation constant

• Kaff = 1/Kd

• Kinetics = rate, mechanism, catalyst, Ea

• Thermodynamics = stability, equilibrium, spontaneity, enthalpy/entropy, free E

Class 5 - Acids and Bases