Unit 1: Highschool Review

Difference between valence and core electrons:

• Valence: Electrons in the outermost shell.

• Core electrons: Do not participate in chemical bonding. Not in the outer shell.

Determining the number of valence electrons and core electrons based on electron configuration for atoms and ions:

• Electrons in the outer most sub-shell are valence electrons.

What is Zeff? (Effective Nuclear Charge) Pattern?

• Average nuclear charge felt by an individual electron in an atom, taking into consideration shielding

• Z_eff = Z - S (Z = # protons in nucleus) (S = Inner shell electrons)

What is atomic and ionic radius? Pattern?

• Atomic radius: Size increases going left and down the table.

  • This is because as you go down a group, a new shell is added, and at the number of electrons in the valence shell decreases.

• Ionic Radius:

  • Cations have a smaller radius. Larger positive charge, the smaller. Due to electron attraction.

  • Anions have a larger radius. Larger negative charge, the bigger. Due to electron repulsion.

What is ionization energy? Pattern?

• Ionization energy is the amount of energy required to pull an electron from the atom/ion.

• Fuller outer sub-shell, greater ionization energy

• Increases going down shells

• E_ion = (kQ1*Q2)/r

What is electron affinity? Pattern?

• Adding an electron to a gaseous atom releases a lot of energy, very exothermic.

• Energy change that occurs when electron is absorbed by a gaseous atom.

• Higher electro negativity = higher electron affinity

• Generally becomes more exothermic going right

What is electronegativity? Pattern?

• Ability of an atom to attract an electron to itself.

• Increases going up and right the PT.

• M like to give away electrons, NM like to attract electrons.

Ionic vs Covalent bonds

• Ionic

  • Donate electrons

  • Stronger that covalent bonds

• Covalent

  • Share electrons

  • Weaker than ionic bonds

How can you predict the nature of a chemical bond? (ionic/covalent, polar/non polar)

• Compare electro negativities

• Pure covalent (Non-Polar): ΔE_N < 0.4

• Polar covalent 0.4 < ΔE_N < 1.8

• Ionic ΔE_N > 1.8

What is lattice energy? What is the trend in solids?

• Amount of energy needed to separate a mole of solid ionic compound into its gaseous ions

  • Ex: NaCl(s) → Na⁺(g) + Cl^-(g)

• ↑ Charge ↓ Ion Size = ↑ Lattice Energy (More attraction)

• ↓ Charge ↑ Ion Size = ↓ Lattice Energy (Less attraction)

• How to find which compound has the higher Lattice Energy?

  • Multiply the charges of all the ions in the compound and take the absolute value.

Unit 2: Lewis Structures

How to draw Lewis Structures?

1. Count number of Valence E

2. Least electro negative is generally in the centre

3. Use all valence electrons

4. Label Formal Charges

Most stable Lewis Structures:

• Least non-zero FC’s

• Negative charge on the most electronegative atom and positive on the least electronegative atom.

Lewis Structures: Octet Rule

• Elements must have 8 surrounding valence electrons

Lewis Structures: Duet Rule

• Hydrogen must be surrounded by 2 electrons.

Radical Species:

• Molecules with an odd number of electrons (unpaired electron)

Hyper Valence (Expanded Octets):

• Elements in the third row can expand their octet

• (Breaks octet rule)

What is formal charge?

• Difference between number of valence electrons and number of electrons surrounding an at0m in a particular Lewis Structure

• Formal Charge = Valence E - Lone Pair E - Bonding

What is the overall molecular charge?

• Sum of FC = Overall Molecular Charge

What are Resonance Structures?

• Same arrangement of atoms, different arrangement of electrons.

How to make/draw resonance structures?

• Lone pairs and double-bond electrons move around

Unit 3: VSEPR Theory

VSEPR Theory

• Predicts molecular shape as point-charges want to be as fr from each other as possible.

Determining Molecular Polarity:

• Are there any lone pairs on the central atom?

  • Yes: Polar

    • No: Are the lone pairs on the surrounding atoms equal?

      • No: Non-polar

      • Yes: Polar

Unit 4: Intermolecular Interactions & Phases of Matter

Intermolecular Forces:

• Intermolecular forces are the attractions between molecules.

Types of Intermolecular forces

• London Dispersion

  • Exist in all atoms

  • Atoms temporarily inducing dipoles

  • Weakest force

• Dipole-Dipole

  • Molecules with a permanent dipole

  • Strong force

• Hydrogen Bonding

  • Molecules with H bonded to N, O, F.

  • Strong Force

• Charge-Charge (ion-ion)

  • Very strong force

  • Ionic solids or ionic liquids

• Charge-Dipole (ion-dipole)

What is Polarizability?

• How easily an electron cloud can be disturbed by an electron field

• Larger atoms/molecules are more Polarizable

Relationships between intermolecular forces and:

• Melting point

  • Stronger intermolecular forces, more energy, higher melting point

• Boiling point

  • More branching of central atom, lower boiling point

• Vapor Pressure

  • Stronger intermolecular forces, lower rate of evaporation, lower vapour pressure

How to predict the types of intermolecular forces in a particular substance?

• Lewis Diagrams

• H - FON

  • Hydrogen Bonding

  • Dipole Dipole

  • London Dispersion

• Polar Molecules

  • Dipole Dipole

  • London Dispersion

• Non-Polar

  • London Dispersion

Interpreting Phase diagrams:

• Critical point

  • Liquid and gas phase are indistinguishable

Phase Changes:

• Trends from going Gas → Liquid → Solid

  • Average IMF increases

  • Molecular Spacing decreases

  • Entropy decreases

• Sublimation

  • Solid to vapour

• Deposition

  • Vapour to gas

• Melting

  • Solid to liquid

• Freezing

  • Liquid to solid

• Vaporization

  • Liquid to vapour

• Condensation

  • Vapour to liquid

Thermodynamic Equilibrium

• No net macroscopic flows of matter of energy within a system or between system and surroundings

• Macroscopic properties of system remain constant

Unit 5: Polymers

What is a monomer?

• Small molecules

What is a polymer?

• Molecule built up from monomers

  • Short-Hand notation

    • Find where polymer pattern is repeating

    • Take that part and put brackets around it

    • Bottom right, outside the bracket write the letter ‘n’.

    • n = number of monomers

What is a oligomer?

• An oligomer is a molecule that consists of a few monomer units

What is a degree of polymerization?

• Amount of repeating units in a polymer chain

What is crosslinking?

• Formation of covalent bonds that hold together polymers

What is elastomers?

• Can be stretched with little or without permanent deformation

How to draw and interpret Line Bond Structures?

• Covalent bonds are the lines

• Unless specified, end of the lines represents a carbon atom

• Each carbon is bonded to H atoms, unless otherwise stated

• Lone pairs can be omitted

How to draw and interpret Condensed Lewis Structures?

• Condensed lewis structures leave out lone pairs and outer bonds.

Functional Groups

• Group of atoms with distinct properties

Types of Functional Groups:

Types of Polymer Linkage

• Addition

  • Two or more molecules join and create a larger molecule without the loss of atoms/molecules

  • 

  • Polymerization:

    1. Initiation

       • Number of radicals increases

       • Starts with the formation of a radical (Refer to unit 1) that is very reactive and an odd number of electrons

       • Abbreviated as R^.

    2. Propagation

       • Number of radicals remains constant

       • Growing polymer reacts with monomeric unit, growing the polymer.

    3. Termination

       • Number of radicals decreases

       • There is a reaction between a growing chain and another radical species

• Condensation:

  • Amide linkage

    • Carboxylic acids or an acid chloride react with AMINES

    • Two monomers join and form a molecule byproduct (water or HCL)

      • Ex/ OH +H₂N forms a byproduct of water

      • Ex/ Cl + H₂N forms byproduct of HCL

  • Ester linkage

    • Carboxylic acid or an acid chloride reacts with ALCOHOLS

    • Forms byproduct of water or HCL

• Reactant 1	Reactant 2	Linkage	Byproduct
  Carboxylic Acid	Amine	Amide	Water
  Carboxylic Acid	Amine	Amide	HCL
  Acid Chloride	Alcohol	Ester	Water
  Acid Chloride	Alcohol	Ester	HCL

What affects polymer properties?

• Polymer architecture

  • Linear

    • Stronger

    • Can pack together easily

  • Branched

    • Weaker

    • Cannot pack together easily

• Molecular weight

  • Greater molecular weight, mechanical strength increases

  • The lower the molecular weight, lower the transition temperature, viscosity, and the mechanical properties

• Crosslinking

  • High-stiffness

  • Stronger

**What is weight distribution?**

Unit 6: Gases

Ideal Gas Law:

→ PV = nRT

• Units

  • Pa, m³, 8.31 j/mol*K

  • kPa, L, 8.31 j/mol*K

  • arm, L, 0.08206 L*atm/K*mol

Ideal vs Real gases

• Ideal:

  • High temperatures and Low Pressures

  • Molecules do not interact

  • Accurate at low pressure because of low interactions

  • Not accurate at high pressures, more interactions

• Real:

  • Molecules do have interactions

  • Accurate at high pressures.

How to use the Van der Waals equation

• (P + an²/V²)(V - nb) = nRT

• a = constant for strength of attractions

• b = constant for size of gas particles

Dalton’s Law of partial pressures

• P_total = P₁ + P₂ + P₃ + P₄…

Mole Fraction

• Ratio of number of moles in a mixture to total number of moles.

Kinetic Molecular Theory:

• Gas is made up many particles in constant random motion

• Gas particles occupy no volume

• Collisions are elastic

• Only interact during collisions

• Average KE is proportional to temperature

When temperature increases

• Average KE increases

Unit 7: Energy and Chemistry

First law of thermodynamics

• Explanation

  • Because heat and work account for all energy exchange between a system and surroundings, energy change must equal the change in the total energy of a system

• Equation

  • ΔU = Q + W

Hess’s Law

• Sum of ΔH for all steps gives overall reaction enthalpy

• ΔH = State Function

What is ΔH

• Enthalpy

• Equal to ΔU

Work

• Object does work on a system = NEGATIVE EXOTHERMIC

• System does work on the object = POSITIVE ENDOTHERMIC

• What is it?

  • Energy transferred between system and surroundings

• PV work

  • W = -PextΔV

    • If ΔV is positive

      • Work is done by system → Negative

    • if ΔV is negative

      • Work is done on system → Positive

Extensive vs Intensive Properties

• Extensive

  • Scaled with the size (QUANTITY) of the system

• Intensive

  • Does not scale with the size (QUANTITY) of the system

Heat capacity

• Quantity of heat required to change the temperature of a substance

• Molar heat capacity → C_pm (mols)

• Specific heat capacity → C_p (grams)

ΔU: Internal Energy

• Total energy of particles in a system

Unit 8: Entropy and the Second and Third Laws of Thermodynamics

Second Law of Thermodynamics

• Reactions increase the entropy in the universe

Third Law of Thermodynamics

• The entropy of a system approaches a constant value as the temperature approaches absolute zero

What is Entropy?

• The amount of disorder in a system

• If increasing, reaction is irreversible

What is Enthalpy

• Amount of internal energy in a compound

Deduce the sign of ΔS for many chemical reactions by examining the physical states of the reactants and products.

How to determine whether Entropy is increasing or decreasing?

• Increasing

  • Number of gas moles increasing

  • Temperature increases

  • \Increasing volume

  • Going from a more ordered state to disordered (solid to liquid to gas)

• Decreasing

  • Number of gas moles decreases

  • Going from a disordered state to more ordered (gas to liquid to solid)

S = klnW

• What is this formula?

  • k is the Boltzmann constant

  • W = is the number of arrangements that are possible in the system

ΔS_universe = ΔS_surr + ΔS >= 0

• ΔS can be positive or negative, but ΔSuniverse must NEVER be negative.

• ΔSuniverse = 0 if all processes are reversible

Standard Molar Entropy

• State function.

• ΔS(knot) = vSproducts - vSreactants (v = coefficient)

What is ΔG?

• Predicts spontaneity of reaction

Ecell positive = spontaneous

• Negative = non-spontaneous

ΔG⁰ = ΔH⁰ - TΔS⁰

• Used to calculate Gibbs

Unit 9: Chemical Equilibrium

Describe the relationship between ΔG°, ΔGrxn, Q, and K,

and and apply this relationship to gain information about

chemical reactions.

What is K?

• K is the way to determine which way the reaction shifts

• Large K → products are favoured, goes to completion

• Small K → reactants are favoured, reactants are favored

• K close to 1 → around equilibrium

ΔG_rxn = ΔG⁰ - RTlnQ

• Difference between this and ΔG⁰ = ΔH⁰ - TΔS⁰

  • ΔG⁰ = ΔH⁰ - TΔS⁰ is for reactions under non-standard conditions (any conditions)

ΔGrxn vs ΔG⁰

• ΔG⁰ tells you about equilibrium position

• ΔGrxn tells you which way it’ll proceed to reach equilibrium

  • At equilibrium equals 0

    • Therefore:

    • ΔG⁰ = - RTlnK

    • K = Q at equiibrium

ln(K2/K1) = - ΔH⁰/R (1/T2 - 1/T1)

• Assuming ΔH knot doesn’t change with temperature we can use this formula.

Solubility Ice Tables and Keq

• Keq = [products]/[reactants] to the power of their coefficients

• Solids and liquids are not in the eq expression

– Differentiate between a reaction quotient and equilibrium

constant and describe the meaning of both.

Reaction Quotient Q vs Keq

• The reaction quotient (Q) is a way to measure the relative amounts of products and reactants present during a reaction at any given moment. It's like a snapshot of the reaction's progress.

• Q compares the current situation to the perfect recipe. If Q doesn't match K, the reaction isn't at equilibrium yet, and something will change to try to get there.

Unit 10: Chemical Kinetics

Define the rate of a chemical reaction and express the rate in terms of the concentrations of individual reactants or products.

Use the method of initial rates to determine rate laws from experimental data.

Use graphical methods to determine rate laws from experimental data.

Use the rate-determining-step to calculate reaction rates in multi-step reactions.

Relate the qualitative principles of collision theory to the quantitative treatment of rates by the Arrhenius equation.

Calculate the activation energy for a reaction from experimental data.

Explain the role of a catalyst in the design of practical chemical reactions.

Explain the importance of both kinetic and equilibrium considerations in the design of industrial processes.

Factors that affect Reaction Rate

• Heterogenous reactants

  • Reactants that have different states (one gas one solid)

• Homogenous reactants

  • Reactants that have the same state (both gases)

• Surface Area

• Concentration

• Pressure (gasses)

• Nature of Elements

• Catalysts

• Increase in Temp

Role of Catalysts:

• Speed up the reaction

• Decrease the activation energy required

Reaction Rate: k[A]^x[B]^y

• A and B are molarity of reactants

• x and y are their order of reaction

Elementary and Intermediates

• Intermediate

  • Is formed and consumed in the reaction

Unit 11: Electrochemistry

Oxidation Number Rules

1. Neutral compound

   • 0

2. Ion

   • Oxidation number adds up to charge on the ion

3. Free Elements

   • 0

4. Fluorine

   • -1

5. Group 1 = +1 Group 2 = +2

6. Hydrogen with Non-Metals

   • +1

7. Hydrogen with Metals and Boron

   • -1

8. Oxygen except with Fluorine or Peroxides

   • -2

9. Group 17 = -1

10. Group 16 = -2

11. Group 15 = -3

LEO GER

• Loss E Oxid

• Gain E Reduc

RED CAT POS

• Reduction is at Cathode which is positive

Writing ElectroChemical Equations

Ecell = E cathode - E anode

• Ecell​=Ecell∘​−nFRT​lnQ

ΔG⁰ = -nFE^0_cell

• F = Faradays constant

  • 96485 C mol^-1