IB Chemistry HL memorization material list

• IB Chemistry HL Memorization Material List
• Ideally, I would like to memorize beyond the required level, but due to time constraints, here is a list of things to memorize in the short term.
• Note that this is just a list of things that I personally feel I need to memorize, so there may be other things that require memorization that are not listed here.
  • I have excluded things that have become self-evident to me even without memorization.
• As of June 2021, I am starting to forget things that I used to consider self-evident.
  • For example, the reasons behind the trend of electronegativity.
  • I have always memorized the surface level, but I have neglected the underlying parts (assuming I would understand them) and I am starting to forget them.
  • So, I need to review those things.

Chapter 1

• Relative Atomic Mass
  • Definition: The average mass of an atom compared to 1/12th the mass of a carbon-12 atom.
  • Unit: g mol-1
• 1 atm = 1.01x10^5 Pa

Chapter 2

• Visible Spectrum: Transition between * and n=2, n=3
• Terms
  • “Shell”: 1, 2, 3, etc.
  • “Subshell”: 1s, 2p, 3d, etc.
  • “Orbital”: 2px, 2py, 2pz, etc.
• Irregular (anti-Aufbau)
  • Chromium configuration: [Ar] 3d5 4s1 (not 3d4 4s2)
  • Copper configuration: [Ar] 3d10 4s1 (not 3d9 4s2)
• 1st ionization energy
  • Definition: The energy required to remove 1 mole of electrons from 1 mole of gaseous atoms under standard thermodynamic conditions.
  • It corresponds to the convergence limit of the spectrum from n=1 (use E=hf).

Chapter 3

• Trends
  • Electron shielding = repulsion between energy levels
  • “Nuclear charge” vs repulsion between electrons in the same shell
  • Increase from bottom left: atomic radius, metallic character
  • Increase from top right: electronegativity, magnitude of electron affinity, ionization energy
    • When ionization energy decreases due to a change in subshell or when pairing of electrons starts in an orbital, repulsion increases.
  • Complex: melting point (greatly affected by bonding)
    • Metals on the left have relatively high melting points due to metallic bonds, those in the middle have the highest melting points due to giant covalent bonds, and those on the right have low melting points due to covalent/noble gas bonds.
• Group 1
  • “Alkali metals”
  • Reaction with Oxygen/Water
  • More reactive down the group
• Group 17
  • “Halogens”
  • Reaction with Alkali Metal 2A(s) + H2 -> 2AH(s) (colorless, neutral)
  • Displacement Reaction
    • Lesser period Halogen with Larger period Halide
    • Example: Cl2 + 2KBr -> 2KCl + Br2
  • Less reactive down the group
• Period 3 Oxide
  • 
  • Structure of oxide
    • Na, Mg, Al: Giant Ionic (strong = solid)
    • Si: Giant covalent (solid)
    • P, S, Cl: Molecular compound (weak intermolecular force = gas/liquid)
  • Reaction of oxide with water
    • Basic: XO + H2O -> 2X(OH)n (acidic)
    • Acidic: XO + H2O -> HnXOm (alkaline)
    • (logic: solid dissolved in neutral water = basic product, so the solid is basic)
• Transition Elements
  • Zn is not a transition element
  • Multiple oxidation states
    • All have +2
    • 
  • Magnetism
    • Paramagnetic: Unpaired electrons -> attracted by a magnetic field
    • Diamagnetic: Paired electrons -> repelled by a magnetic field
    • Ferromagnetic: Many unpaired electrons aligned -> Strong magnetism (Ni, Co, Fe)
    • Unpaired electrons cause magnetism
  • Ligands: coordinate covalent bonds
    • Neutral: H2O, NH3, CO
    • 1- Ligands: Cl-, CN-, Br-
  • Catalyst
    • N2 + 3H2 ⇆ 2NH3 with Iron
  • Colored complex
    • Repulsion with electrons of ligands
    • Splitting of 3d orbital, 2 upper and 3 lower: when ligands are attached
    • Absorption of light (low to high energy level)
    • Spectrochemical series of ligands
      • Stronger field -> more splitting -> absorb shorter wavelengths
    • 
    • Remember ligands, coordination number, etc.
      • 

Chapter 4: Chemical bonding and structure

• 

  • Carbonate, chromate, phosphate, and sulfate have a charge other than -1.
  • Chromate, phosphate, and sulfate are exceptions; they have 4 oxygen atoms instead of 3.

• Metallic Bonding Factors

  • Number of delocalized electrons
  • Charge of the cation
  • Ionic radius of the cation

• Covalent Bonding

  • VSEPR
  • Lone pair-lone pair > lone pair-bonded pair > bonded pair-bonded pair
  • Expanded octet is possible for period 3 and beyond.
  • 
  • Hybridization
    • In the end, it is about adding +1 for each lone pair and +1 for each σ bond.
    • In other words, when drawing the structure, count how many sides the lines representing electron pairs are extending from.
  • Sigma bond: round + round
    • Single bond
  • Pi bond: dumbbell + dumbbell
    • Double or triple bond between 2 p-orbitals
    • Cloud above and below the σ bond

• Formal Charge- (Valence electron in uncombined form) - (number of non-bonding electrons) - ½(number of bonding electrons)

• The closer the result is to 0, the more preferable it is.

• Intermolecular forces (only in simple covalent compounds) (Lower = Stronger)

  • London Force: Instantaneous dipole-induced dipole
  • Dipole-Induced-dipole
  • Dipole-dipole
  • Hydrogen Bond: strong Dipole-dipole (between H and N, O, F)
  • More electrons = stronger intermolecular forces
  • Similar intermolecular forces = solubility

Chapter 5: Energetics

• Open system: allows for exchange of all substances
• Closed system: allows for exchange of energy only
• Isolated system: does not allow for any exchange

Chapter 6: Kinetics

• Conditions for a reaction to occur:
  • Correct orientation of reactant molecules
  • Kinetic energy (KE) greater than activation energy (AE)
• Definitions of specific enthalpy
  • Default:
    • Direction: as the name suggests
      • For example:
      • atomize = convert to a single gas atom
      • ionization = convert to a positive ion
    • Amount: 1 mole of a measurable substance (sometimes reactant, sometimes product)
  • Lattice enthalpy (Irregular)
    • Direction: from broken ions to lattice
    • Amount: 1 mole of lattice
  • It is important to understand the trend of enthalpy (can usually be understood intuitively)

Chapter 7: Equilibrium

• Conditions for equilibrium:
  • Closed system
  • Rate of forward and backward reactions are equal
  • Concentration of reactants/products remains constant
• Haber Process (Ammonia)
  • N2 + 3H2 ⇄ 2NH3
    • Exothermic, decrease in the number of moles
  • Conditions for the forward reaction:
    • Low temperature
      • However, some temperature is needed for a high reaction speed
      • Optimum temperature: 450°C (not K)
      • (A high reaction speed means faster reaching equilibrium)
    • High pressure
      • Optimum pressure: 200-250 atm
      • Too high pressure can cause the plant to break
    • Catalyst
      • Catalyst is needed, otherwise lower temperature or higher pressure must be used
    • “Fresh” iron should be used
  • Usage: Fertilizer, explosives, etc.
• Contact Process (Sulfuric Acid)
  • S + O2 ⇄ SO2
  • SO2 + O2 ⇄ 2SO3 (Important step)
    • Exothermic, decrease in the number of moles
  • SO3 + H2O ⇄ H2SO4 / SO3 + H2SO4 ⇄ H2S2O7
  • Conditions for the forward reaction:
    • Low temperature
      • 400°C or higher
      • Catalyst only works at 400°C or higher
    • Catalyst
      • Vanadium (V) Oxide
    • No high pressure
      • The rate is already sufficient
      • High pressure is not cost-effective and dangerous
  • Usage: Fertilizer, Paints, etc.
• 

• At equilibrium, the Gibbs free energy is at a minimum and entropy is at a maximum.

Chapter 8 Acid and bases

• Definitions

  • Bronsted-Lowry

    • Acid: donates H+ ions
    • Base: accepts H+ ions
      • has lone pair of electrons or forms a coordinate bond with H+
    • Acid + Base → Conjugate base and acid (only for Bronsted-Lowry)

  • “amphiprotic” vs “amphoteric”

    • “amphiprotic” refers to substances that can act as both acids and bases (narrow condition)
    • “amphoteric” refers to substances that can act as either acids or bases (broad condition)
      • Both are possible (not none, but both)

  • Lewis Acid

  • Acids to remember

    • 

  • Bases to remember

    • 

  • Classification

    • mono/di/triprotic acid: number of donatable protons
      • Note: CH3COOH is monoprotic
    • strong/weak
      • strong: completely dissociates / weak: does not completely dissociate
      • organic acids are weak
      • how to distinguish
        • pH
        • conductivity (strong acids have higher conductivity)
        • relative rate with metal (only for acids)
      • strong acid → weak conjugate base, weak acid → strong conjugate base

  • Reactions

    • Salt
      • Definition: an ionic compound formed between a metal and non-metal
    • Acid + Metal → Salt + Water
    • Acid + Base → Salt + “messy” product
      • Acid + Metal Oxide (O) → Salt + Water
      • Acid + Metal Hydroxide (OH) → Salt + Water
      • Acid + Metal carbonate (CO) → Salt + Water + CO2
      • Acid + Metal Hydrogencarbonate (HCO) → Salt + Water + CO2
    • Neutralization (acid + base → salt + water)
      • Exothermic
        • always around -57 kJ/mol
      • Application: antacids
      • Titration
        • “titrant” = alkali in burette (known concentration)
        • “analyte” = acid in beaker (unknown concentration)
        • indicator: phenolphthalein, etc.

  • Ionic product of water Kw = 1*10^-14

    • Kw increases (water dissociates more) when temperature increases

  • larger Ka/Kb = smaller pKa/pKb = stronger acid/base

    • Ka + Kb = Kw = 14 at 297K

  • Buffer Solution

    • Definition: When (1) a SMALL amount of acid/base is added to
    • weak acid : salt
    • 1 weak acid : 0.5 strong base
• Causes of Acid Deposition (i.e. acid rain)
  • nitric/ous acid: 2NO• from lightning and engines + O2 → NO2• / NO2• + water → nitric/ous acid
  • sulfuric/ous acid: S from power stations / volcanoes + O2 → SO2 / SO2 + H2O → sulfuric/ous acid
  • (There are other details such as environmental impacts and reaction explanations, but let’s leave them for now)
  • Memorization of many facts is required
• Equivalence Point
  • Strong acid, weak base: <7
  • Strong base, weak acid: >7
• Other
  • If H2CO3 appears, it is probably H2O + CO2

Chapter 10 Organic Chemistry

• Homologous series: have the same difference in CH2, same functional group, same general formula

  • Physical properties gradually change, chemical properties are similar

• Isomers: branch, functional, positional

  • Functional isomers: aldehyde/ketone, alcohol/ether
  • Trend: More branching = harder to approach = weaker intermolecular forces = lower boiling point

• ![image](https://gyazo.com/b521f597d0b06ce- Benzene

  • Only substitution reactions, no addition reactions
  • Benzene is classified as “aromatic” (not aliphatic)

• Alkane

  • Non-reactive except with functional groups | Name | Combustion | Addition | Substitution | Polymerization | Oxidation | Reduction | | — | — | — | — | — | — | — | | Benzene | | | ✔︎ Electrophilic | | | | | Alkane | ✔︎ | | ✔︎ Free radical | | | | | Alkene | | ✔︎ Electrophilic | | ✔︎ | | | | Ester | | | ✔︎ Nucleophilic | | | | | Alcohol | ✔︎ | | ✔︎ Nucleophilic | | ✔︎ | | | Aldehyde | | ✔︎ Nucleophilic | | | ✔︎ | ✔︎ | | Ketone | | ✔︎ Nucleophilic | | | ✔︎ | ✔︎ | | Carboxylic acid | | | ✔︎ Nucleophilic | | | ✔︎ | | Halogenoalkanes | | | ✔︎ Nucleophilic | | | |

• I made a table for better understanding

  • Checked with the teacher and excluded topics not in the syllabus
  • Now I have a general understanding, I want to understand the details
    • As a result, I want to be able to derive this table using logic without memorizing it

• Reactions

  • Combustion

    • Complete combustion produces CO2 and water
    • Incomplete combustion produces CO/C and water (lacks oxygen)
    • Water is always produced!

  • Substitution

    • Free radical

      • Initiation (increase in free radicals)
        • Homolytic fission
      • Propagation (no change in the number of free radicals)
      • Termination (decrease in free radicals)
      • Condition: UV light

    • Nucleophilic

      • Attacked by a nucleophile (prefers positive charge)

      • Nucleophile attacks the polar C-X bond (X: δ-)

      • Halogenoalkanes

        • SN2 (bimolecular)
          • rate = [halogenoalkane][nucleophile]
          • Lower order halogenoalkanes prefer SN2
            • Hard to attack due to bulkiness of alkyl groups in primary halogenoalkanes
        • Secondary: Depends
          • Neither too bulky nor too open
        • Tertiary: SN1 (unimolecular)
          • rate = [halogenoalkane]
          • Higher order halogenoalkanes prefer SN1
            • Carbocation is more stable in tertiary halogenoalkanes
        • Trend of rate
          • C-X bond of Halogenoalkane
            • X with a larger radius has a lower bond enthalpy, resulting in a higher rate
            • The trend of electronegativity difference is negligible
          • Nucleophile (only for SN2 reaction, because SN1 is independent)
            • Anion is better than a polar molecule
          • Solvent
            • Polar protic solvents prefer SN1
              • Protic solvents help break the H-X bond by pulling both sides
              • Protic solvents solvate intermediate carbocations and stabilize them
            • Polar aprotic solvents prefer SN2
              • Aprotic solvents do not dissociate nucleophiles, so nucleophiles are not surrounded by atoms and are more reactive
                • Nucleophile’s reactability is important in SN2 but not in SN1

      • Special: Esterification

        • Nucleophile: Alcohol
        • Condition: Heat, concentrated H2SO4
        • Written as a reversible reaction
        • Also a condensation reaction

      • Special: Hydrolysis

        • SN1/SN2

    • Electrophilic

      • Nitration of benzene
        • Condition: Heat under reflux, 60℃
          • Reflux prevents OH and other substances from escaping

  • Addition

    • Types of possible reactions and their conditions are skipped for now

      • Memorize if necessary

    • Electrophilic

      • Alkenes (C=C bond) undergo this reaction
      • Electrophile: Cation or the positive side of a dipole
      • Even if two atoms come together, if the initial attack is by an electrophile, it is called electrophilic addition
      • Major product: The carbon with fewer hydrogens becomes a carbocation
        • Less hydrogen = more atoms with positive inductive effect = more stable carbocation
      • Examples
        • H-Br, Br-Cl (dipole)
        • Br-Br (induced dipole by high electron density)
        • H-H (hydrogenation)

    • Nucleophilic

    • Application: Test with bromine water

      • Alkanes: React with Br2 + H2O, decolorizes
      • Alkane/Benzene: No reaction, no decolorization

  • Polymerization

  • Oxidation

    • Agent: Acidified K2Cr2O7 or acidified KMnO4
      • Acidified means H+ is present (and used for oxidation)
    • Primary alcohol -> Aldehyde -> Carboxylic acid
      • Complete oxidation under reflux
      • Partial oxidation under distillation
        • Aldehyde has a lower boiling point
    • Secondary alcohol -> Ketone

  • Reduction

    • Aldehyde / Ketone -> Alcohol
      • Reducing agent: LiAlH4 (stronger), NaBH4
    • Carboxylic acid -> Alcohol
      • Reducing agent: LiAlH4 ONLY
      • NaBH4 is weaker

  • Esterification

    • Catalyst is required

• General Laws

  • Polar bonds are weaker
  • Things that cause charges
    • Electron density
      • C=C bond has high electron density
    • Dipole
    • Ions

Chapter 11

• NMR
  • Instead of 1,2,3, we use terms like singlet, doublet, triplet
  • Why TMS (tetramethylsilane):
    • High boiling point (easy to remove)
    • Gives a single strong signal
    • Non-reactive