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FULL SUMMARY - IB SL Chemistry (Imaginfinity)

❏ Convergence limit​ when lines of hydrogen spectrum converge at
higher freq/energy (n=7 to n=∞) = enough E to remove e- from atom
❏ Continuous spectrum​ contains all wavelengths of light, whereas ​line
spectrum​ contains discrete wavelengths of light
❏ Lyman​ series (UV, n=1), ​Balmer​ series (visible, n=2), ​Paschen
series (IR, n=3)
❏ Energy of photon = change in energy of transitioning
electron, evidence that energy levels get closer as farther
from nucleus (7 to 6 lower energy than 2 to 1, thus closer)
❏ Shorter wavelength = higher frequency = higher energy
❏ Relative atomic mass​ is the weighted mass average of all naturally
occurring isotopes of an atom relative to 1/12 the mass of a C-12
❏ Percent abundanc​e is % of each isotope in composition of an element’s relative atomic mass
❏ Isoelectronic​ means two elements have the same electron configuration, represented by square brackets to
condense electron configuration such as [Ar]4s1 is sodium
❏ Electron configuration shows electrons in each orbital (1s²2s²), electron arrangement shows number of
electrons in each energy level (2, 8, 8), box orbital diagrams show electrons in each orbital and amount of
energy of orbitals
❏ Aufbau’s Principle ​= electrons will fill sublevels in order of
increasing energy since lower levels are more stable
❏ When 3d occupied, 4s electrons removed first in ionization
❏ Pauli’s exclusion principle​ = no electrons can have same 4 quantum
numbers + electrons in same orbital have opposite spin
❏ Hund’s rul​e = electrons of equal energy fill orbitals of sublevel one
at a time until all orbitals have one electron before an orbital will
accept a second electron (minimizes repulsion)
❏ Principle energy levels​ (n=1, n=2) split into ​sublevels​ (1s, 2s, 2p) which
contain ​orbitals
❏ Max electrons in an E level is 2n², Max orbitals in an E level is n²
❏ Electron energy = quantized b/c require a quantum (discrete quantity
of energy) of energy to move to a different sublevel
❏ Promotion​ occurs when an electron is promoted to an orbital with more energy, such as Cu
❏ Radioisotopes = C-14 is used for radiocarbon dating
❏ Atomic # (z) = number of p+ or e- , atomic mass (A) = number of p+ and nº together
❏ Mass spectrometer ​shows the % abundance and mass/charge ratio of an element = calculations
❏ Calculate relative atomic mass from % abundance of each isotope
❏ Calculate % of an isotope (represented by X) from relative atomic mass and % of other isotope(s)
-- could be 1-X if there are >1 unknown % abundances
❏ Calculate % abundance of a molecule from an element (such as % of Cl, then % of Cl2)
❏ Intermolecular forces​ = b/n molecules, ​intramolecular forces​ = b/n atoms in compound
❏ Ionic bonding​ is the electrostatic attraction between oppositely charged ions
❏ Ionic compounds are ​brittle​ (like ions repulse), ​low volatility​ (lots of E to disrupt ionic bonds),
good electrical conductors when molten/aqueou​s (ions), ​soluble in water​, not hexane (partial
charges in polar substance like water attract to ions + pull from lattice)
❏ Covalent bonding​ is the electrostatic attraction b/n a pair of electrons and the positively charged nuclei
❏ Triple bonds shorter and stronger than double bonds, which are shorter and stronger than single
❏ Increasing bond length = decreasing bond strength
❏ Metallic bonding​ is the electrostatic attraction between a lattice of positive ions and delocalized electrons
❏ Magnitude based on​ # of delocalized electrons,​ ​size of cation​ (radius), and ​charge of cation
❏ Metals = ​good conductors​ of heat and electricity, ​malleable​, ​ductile​, ​lustrous​, ​high MP​ (​ mobile
delocalized electrons, random electron movement allows ions to slide/rearrange w/o breaking
❏ Coordinate covalent bond​ is a covalent bond where both electrons originate from one atom (the donor
atom must have a lone pair)
❏ Van der waals ​= london dispersion forces, dipole-dipole interactions, hydrogen bonding
❏ LDFs​ are temporary attraction between adjacent molecules due to instantaneous dipoles
❏ Magnitude based on ​size​ (# moving electrons) and ​shape​ (how tightly packed)
❏ Dip-dips​ is attraction between permanent dipoles (polar molecules)
❏ H-bonding​ is attraction between hydrogen atom and highly electronegative NOF atoms (with LP)
❏ Lewis structures​ = valence electrons (use lewis calculations: Needs, Haves, Electrons, Bonds)
❏ Electronegativity​ is relative attraction of nuclei for shared pair of electrons in covalent bond
❏ Polarity​ 0-0.4 Non-Polar covalent, 0.4-1.8 Polar covalent, >1,8 ionic
❏ Can have polar bond and no net dipole (NP molecule) if symmetrical structure
❏ VSEPR theor​y determines ​molecular geometry​: lone pairs take up more space than bonded and electrons
positioned as far apart as
possible to minimize
❏ 2 electron domains - linear
❏ 3 electron domains - trigonal
planar 120º, angular planar
❏ 4 electron domains tetrahedral 109.5º, trigonal
pyramidal 107.5º, angular
planar 104.5º
❏ 5 electron domains - trigonal
bipyramidal 90º + 120º,
see-saw <90º + <120º,
t-shaped 90º, linear 180º
❏ 6 electron domains octahedral 90º, square
pyramidal 90º, square planar
90º, t-shaped, linear 180º
❏ Solubility​ - whether can interact (similar polarity), max solute can dissolve in solvent at certain
❏ Electrical conductivity​ - mobile charges (ions or delocalized electrons), thus molton/aqueous (not solid)
ionic compounds, metals, and acids/bases are good electrical conductors
❏ Thermal Conductivity​ ❏ Melting and boiling points​ - strength of intermolecular forces, whether will disrupt
❏ Strength of IMFs: Van der waals (LDFs < dip-dips < h-bonds) < ionic bonding < metallic bonding
< covalent bonding
❏ Vapour pressure/volatility​ - strength of intermolecular forces
❏ Incomplete octets​ for beryllium (quartet) and boron (sextet) since small/low ShE, stable with <8 electrons
❏ Expanded octet​s when more than 8 electrons around one atom, don’t use lewis calculations (draw out
atom with valence electrons, put other atoms around + expand pairs, determine LPs to find molecular
❏ Resonance​ when >1 option for placement of double/triple bond = delocalized electrons (resonance
❏ Network covalent solids​ formed by carbon and silicon (carbon forms diamond, graphite, graphene,
❏ Diamond = good thermal conductor, hard, high MP; Graphite = good electrical conductor,
soft/slippery due to layers, weak LDFs b/n layers; Graphene = good electrical conductor, best
thermal conductor known (single layer/atom), thin, 100x stronger than steel, high MP; C60 =
semiconductor (some e- mobility), low MP
❏ Alloys​ are compound made up of >1 metal, have enhanced properties = more chemically stable, stronger,
resistant to corrosion (ex. steel, stainless steel, brass, bronze, sterling silver)
❏ Stainless steel = iron + Ni or Cr, high tensile strength and corrosion resistant
❏ Groups​ are vertical columns (# of electrons), ​periods​ are horizontal rows (# occupied principal energy
❏ Modern periodic law​ states that when elements are arranged in increasing atomic number, there is
periodic recurrence​ and ​gradual change​ in both physical and chemical properties
❏ Effective nuclear charge​ (Zeff) is the attraction a nucleus exerts on its valence electrons
❏ Increases across a period, constant down a group
❏ Shielding Effect​ (ShE) is the interference b/n nucleus and valence electrons provided by core electrons and
distance b/n nucleus and valence electrons (# E levels)
❏ Increases down a group, ​relatively​ constant across a period
❏ Zeff/ShE ratio​ reflects overall attraction an atom has for its valence electrons (increase across period,
decrease down group) -- structure dictates function
❏ Atomic radius​ is half the distance
b/n neighbouring nuclei
❏ Increases down group,
decreases across period
❏ Ionic radius ​is half the distance
b/n neighbouring nuclei of ions
(treated as if in lattice)
❏ Increases down group, decreases across period
❏ Anion larger than cation with same parent atom or with same electron configuration
Ionization energy​ is the minimum amount of energy required to remove one mole of electrons from one
mole of gaseous atoms in their ground state ​X(g) + E(kJ/mol) → X+(g) + e❏ Increases across period, decreases down group
❏ Exceptions: group 13 lower than 2 (p subshell farther from nucleus), group 16 lower than 15 (15
afforded greater stability ½ complete p subshell, more e- repulsion in 16)
Electronegativity​ is relative attraction of nucleus for shared pair of electrons in covalent bond
❏ Increases across period, decreases down group
Electron affinity​ is the energy change when one mole of electrons added to one mole of gaseous atoms to
form one mole of gaseous ions ​X(g) + e- → X-(g) + E(+or-)
❏ Increases across period, decreases down group
❏ Exceptions: group 15 and group 2 lower than expected (15 has more e- repulsion/takes more
energy to add since more stable with 3 e- in p subshell, 2 has complete s subshell)
Melting point​ increases across period to max at group 14, then falls to min at group 18
❏ Covalent IMFs for Si and C, increases with metallic bond strength for metals (charge), decreases
for nonmetals b/c of weak van der waals
❏ Decreases down group of metals (metallic bonds weaker since increased ShE)
❏ Increases down group of nonmetals (LDFs stronger down group with more electrons)
Period 3 oxides​ illustrate the transition from metallic to non-metallic character
❏ Acid/base character​ - metal oxides (Na2O and MgO) dissolve in water to form basic solutions,
amphoteric oxides (Al2O3) react as acid or base (when with acid or base), nonmetal oxides (SiO2,
P4O10, P4O6, SO3, SO2, Cl2O7, Cl2O) dissolve in water to form acidic solutions
❏ Electrical conductivity​ high for metal oxides, very low for SiO2, none for nonmetal oxides
❏ Structure​ is giant ionic for metal oxides (ionic bonds), giant covalent for SiO2 (network covalent
bonds), molecular covalent (covalent bonds) for nonmetal oxides
Metallic character​ (luster, high density, conductive, malleable, ductile) increases down group when more
available delocalized electrons, decreases across period when become nonmetals
❏ Halide displacement ​X¹2(g, l, or s) + 2MX²(aq) → 2MX¹(aq) + X²2(g, l, or s)​, only more reactive
halogen displaces (F>Cl>Br>I), colour change, darkens/orange when Cl displaces I, orange when
Cl displaces Br, darkens when Br displaces I
❏ Alkali metals with water ​2M(s) + H2O(l) → 2MOH(aq) + H2
❏ Halogens with silver ​2Ag(s) + X2(g, l, or s) → 2AgX(s) ​= silver halide precipitates
❏ Alkali metals with halogens​ 2M(s) + X2(g, l, or s) → 2MX(s) ​= ionic halides
Reactivity​ of alkali metals increases down group (more ShE) but decreases down group for halogens (more
ShE, but need to gain e- to form anions)
❏ Empirical formula​ is the smallest whole number ratio of the elements in a compound: ​percent to mass,
mass to moles, divide by small, multiply to whole
❏ Molecular formula​ is the actual ratio of elements in a compound (may be same or different than
empirical), use ​Ma​ ctual​/M​empirical​ to determine
❏ Molecular structure, naming, ​structural​ = draw structure, ​condensed ​= CH3COOH
❏ The ​mole​ allows for observable measurements of chemical entities (atoms, molecules, FUs)
❏ Molar mass​ is the mass (g) of one mole of a substance ​M=m/n​, ​from mass to mole and moles to
mass, all you need is molar mass
❏ Avogadro’s constant ​is 6.02 × 1023 entities, ​NA=entities/n
❏ Mole ratio​ is used in stoich grid (coefficients) to convert b/n substances
❏ Percentage composition by mass​ is % of each element in a compound ​Me​ lement​/M​total​ x 100%
❏ Limiting reagent​ is reactant that is used up completely in reactions (reaction stops when used)
❏ Use stoich grid to calc amt (n or m) of product for all reactants, LF yields smallest amt
❏ Use LF as initial value to calculate amt product formed
❏ Excess reagent ​is the reactant that is in excess, therefore there is some left over after reaction
❏ Calculate mass consumed by reaction with LF and subtract from total mass to know how much XS
remaining (add 10% to know how much sufficient for LF consumption)
❏ Stoich grid used to perform unit analysis, fill in info you know
❏ Avogadro’s​ ​law​ states that equal volumes of gases, measured at the same temperature and pressure, have
the same number of particles V ∝ n
❏ Molar volume​ is the volume (dm³) occupied by one mole of any ​gas​ Vm=V(dm³)/n(mol)
❏ STP​ = 100kPa ( 105 Pa), 273K (0ºC)
❏ SATP​ = 100kPa ( 105 Pa), 298K (25ºC)
❏ Molar volume of an ideal gas at STP​ is 22.7dm³/mol or 2.27 × 10−2 m³/mol
❏ n (mol) = volume (V) x molar volume (dm³/mol)
❏ Ideal gases​ - gases deviate from ideal gas behaviour most at low temps b/c particles slower, more
opportunity for potential IMFs to take effect
❏ Kinetic theory of matter​: gases as largely empty space containing free-moving particles of
negligible volume having no inter-particle forces
❏ Ideal gas equation: ​PV=nRT
·kP a
m ·P a
❏ R​ is the ​universal gas constant​: ​8.31 dm
mol·K or 8.3​1 mol·K
P 1V 1
P 2V 2
❏ Boyle’s law P ∝ 1/V , the pressure of a gas is inversely proportional to its volume (increasing
volume decreases the frequency of collisions/pressure)
❏ Charles’ law V ∝ T the volume of a gas is directly proportional to its absolute temperature
(increasing temperature increases kinetic energy, volume increases to keep pressure constant)
❏ Gay-lussac’s law P ∝ T the pressure of a gas is directly proportional to its absolute temperature
(increasing temperature increases kinetic energy, more frequent collisions)
❏ Solutions
❏ Electrolytes​ are solutes that conduct electricity - ionic compounds (dissociate), acids/bases
(ionize) since have free-flowing charges (not molecules since neutral)
❏ Aqueous solutions have water as solvent, always transparent, sometimes colourless
❏ Dilute​ solution has small amt solute relative to solvent, ​concentrated​ solution has large amt
solute relative to solvent
❏ Dilution equation​ C 1 V 1 = C 2 V 2 ​‘says R2D2’
❏ When know initial C and V and add two together, have to find the total moles, then the
total volume, then divide n/V for the overall concentration!
❏ PPM = m​solute​/m​solution​ x 1,000,000​ (need to be same unit)
❏ Molar concentration​ is moles of solute per volume (dm³) of a solution ​C = n/V(dm³)
❏ Standard solution​ is a solution with a known concentration
❏ Represent concentration with square brackets around substance
❏ Dissociation equations ex.​ Al₂O₃(s) → 2Al³+(aq) + 3O²-(aq)
❏ Titrations
❏ Acid-base titration
❏ React solution with unknown concentration (​analyte​) with a
strong acid/base (​titrant​, standard solution) in neutralization
reaction, determine volume titrant needed to neutralize known
volume of analyte (of unknown conc.), use mole ratio to calculate
❏ Neutralization reaction​ is a reaction b/n an acid + a base to
produce salt + water
❏ Total ionic equations​: acids/bases/ionic cmpds ionize/dissociate
into ions, g/l/s phases don’t form ions (solid NaCl has no free
ions, H2O molecule has no ions) ​H+(aq) + Cl-(aq) + Na+(aq) +
OH-(aq) → H2O(l) + Na+(aq) + Cl-(aq)
❏ Net ionic equations​ remove spectator ions (identical on both sides
of equation), same in all neutralization reactions: ​H+(aq) +
OH-(aq) → H2O(l)
❏ Back titrations
❏ React analyte with known amount XS reagent, remaining XS titrated with another
reagent, calculate how much XS reagent reacted in first titration by using second titration
results, then calculate original sample’s concentration (go past endpoint on purpose in
1st, then return back to endpoint in 2nd)
❏ Behaviour of water: OH- electronic environment is different, H+ leaves behind e-, H in COOH is
acidic hydrogen (since other, ex in CH3 are not in EN environment)
❏ Random errors​ occur when there is equal probability of data being too high or too low; caused by
readability of instruments, insufficient data, misinterpretations; reduced by repeated measures
❏ Systematic errors ​occur as a result of poor experimental procedure, reduced by careful design
❏ Measured values ​have limited certainty, ​counted/exact values​ have infinite certainty
❏ Variables and graphs
❏ Quantitative data​ is numerical/from measurements, ​qualitative data ​is non-numerical/from observations
❏ Accuracy​ is how close to true value, ​precision​ is how close values are together
❏ Rounding​: greater than 5 round up, less than 5 unchanged, # is 5 digits after round up, 5 with zeros after
round up if odd and unchanged if even
❏ Significant Digits:​ non-zero are significant, zeros b/n sig digs and final zeros in decimal portion are
significant, all digits in scientific notation are significant (before exponent)
❏ Adding/subtracting = round to original with fewest ​decimal places
❏ Multiplying/dividing = round to original with fewest ​sig digs
❏ Uncertainty​: ​fraction​ (absolute/measured value), ​percent​ (fraction uncertainty x 100%), ​absolute​ (from
scale, ex. +/- 0.1)
❏ Analogue instruments​ = half the smallest division
❏ Digital instruments​ = smallest scale division
❏ Adding/subtracting = sum of absolute uncertainties
❏ Multiplying/dividing = sum of individual % uncertainties, then calculate absolute
❏ Conversion factors​ are an equality b/n two units (some are counted, some are measured)
❏ Index of hydrogen deficiency​ used to measure how many molecules of H2 are needed to convert a
molecule to its corresponding saturated, non-cyclic molecule
❏ Count number of rings and multiple bonds (each adds 1, triple adds 2)
❏ Halogens treated like H atoms, sulfur + oxygen don’t affect, add nitrogens to C and H
(2C + 2 − H + N )
❏ IHD =
❏ Mass spectroscopy​ - determine empirical formula from compositional analysis, molecular formula from
molecular ion​ mass (greatest), possible structures, analyse fragmentation pattern (mass difference), identify
fragments (table 28) *fragments present have positive charge*
❏ Infrared spectroscopy​ - what you expect to see (​polar bonds​ in molecule, only polar excited by E in IR
region), where peaks should be (table 26), identify peaks, fingerprint region 500-1500 cm-1 (wavenumber)
❏ H NMR spectroscopy​ - different ​chemical shifts​ in different chemical environments (H positions),
measured against standard signal of 12 H nuclei in tetramethylsilane (TMS), integrated trace is relative # of
H atoms in each environment (CH3, CH2, CH, OH, if same, look at atoms bonded too)
❏ Chemical System ​is the chemicals of interest undergoing a change and the ​surroundings ​are the rest of the
universe (including boundary)
❏ Thermal energy ​(Q) is energy available from substance as result of motion of particles
❏ Heat ​(q) is the transfer of thermal energy b/n substances with diff. temperatures until thermal
equilibrium achieved
❏ Temperature​ is a measure of the average kinetic energy of particles
❏ Heat change formula​ q = mcΔT
❏ Heat capacity​ is quantity of heat to raise temperature of a substance by 1K or 1ºC
❏ Specific heat capacity is of 1g ( J · g −1 · ºC )
❏ Molar heat capacity is of 1mol ( J · mol−1 · ºC )
❏ First Law of Thermodynamics​: energy conserved (not created/destroyed)
❏ Euniverse = Esystem + Esurroundings therefore ΔHsystem =− q surroundings
❏ Change in enthalpy​ is the heat transferred b/n system and surroundings, carried out under
constant pressure = ​state function​ (phase matters) and ​pathway independent
❏ Second Law of Thermodynamics​ when 2 objects in thermal contact, heat always transferred from object
at higher temperature to object at lower temperature until same (thermal eqm)
❏ Endothermic: + ΔH , feels cold, absorbs heat E (melting, vaporizing)
❏ Exothermic: − ΔH , feels hot, releases heat E (condensing, freezing)
❏ Use potential energy diagrams for endo/exothermic reactions
❏ Heating curve represents state changes
❏ Enthalpy Change depends on ​change in energy to break/form bonds a​ nd the
strength of IMFs​ (temp/pressure dependant)
❏ Standard conditions: pressure ​100kPa​, concentration ​1mol/dm³ ​and
all substances in standard states (temp usually 298K but not part)
❏ Molar Enthalpy Change​ in K J · mol−1 (c for combustion, f for
formation, fus for melting, vap for vapourization, general= ΔHrxn )
❏ Put ΔH after​ or ​in ​balanced equation (on right for exothermic, on left for endo)
❏ Thermochemical eqn​: species of interest as 1 mol, fractional coefficients
❏ Calorimetry​: measure heat flow with a calorimeter
❏ Bomb or high school calorimeter​ ΔHrxn = − (mcΔT )/nsys
❏ Metal or flame calorimeter​ ΔHrxn = − [(mcΔT )H2O + (mcΔT )cal]/nsys
❏ Metal/flame calorimeter: surroundings are water and metal calorimeter so add
❏ Dissolving: water is surroundings, solute is system
❏ Neutralization: surroundings are dilute acid/base solutions (​assumed​ to be pure water so
1.00g/cm³ and 4.18J/g/ºC)
❏ Compensating for heat loss (to surr/calorimeter) ​Tsurr = T3 - T1
❏ Hess’s Law​: the ΔH for a rxn carried out in a series of steps is
equal to the sum of the ΔH ’s for the individual steps
❏ Use Hess’s Law to determine ΔHrxn (cross out finished
rxns + write new ones as go):
1. Are reactants on the left? (if not, multiply by
2. Are products on right? (if not, multiply by -1)
3. Manipulate so intermediaries are on opposite sides of the reaction
4. Multiply intermediaries to eliminate
5. Write final rxn and add ΔH ’s
❏ Gasification of Coal​ = coal burning produces CO2, steam passed over bed of coal at 600ºC in
anaerobic conditions (doesn’t change ΔH , infeasible in developing nations)
❏ Molar Enthalpy of Formation ​( ΔHºf ) is the enthalpy change associated with the formation of mole of a
compound in its standard state from its constituent elements in their standard states
∑(ΔHºf ·nproducts)−∑(ΔHºf ·nreactants)
ΔHrxn =
n species of interest
❏ Formation equations​ have only 1 mole of product and only elements in standard states are
reactants (all are synthesis reactions)
1. 1 mole of product (write → write next to it)
2. ΔH right beside it
3. Reactants on left (use H2, O2, not H2O) and BALANCE
❏ Average Bond Enthalpy​: the energy needed to break 1 mole of a bond in a gaseous molecule averaged
over similar compounds under standard conditions
∑ E (bonds broken)−∑ E (bonds f ormed)
ΔHrxn =
n species of interest
❏ Breaking bonds: in ​reactants​, E input, endothermic (separating attracted atoms ex. Cl2)
❏ Forming bonds: in ​products​, E output, exothermic (bringing together attracted atoms, thus lower
energy state, ex. H+ ions together to form H2)
❏ Average bond enthalpies depend on ​number of bonds in molecule ​(e- milieu) and the ​molecule
the bond is located in ​(ex. minor vs large hydrocarbon)
❏ Ozone depletion: O2 double bond is stronger than the 1.5 bond in O3 (ozone), more E to break, keeps
stratosphere warm (absorbs UV light) O2→2O・; O・+O2→O3 ; O3+O・→ 2O2
❏ Rate of Reaction​: change in [c] of reactant or product per unit time ( mol · dm−3 · s−1 )
❏ Rate of disappearance of reactants r(average) = − Δ[c]/Δt (since rate is +ve)
❏ Rate of appearance of products r(average) = Δ[c]/Δt
y −y
❏ Slope of tangent = x2 −x1 (draw line though pt where finding rate)
❏ Determined by ​frequency of collisions ​and​ fraction of effective collisions
❏ Collision Theory​: successful collisions require ​sufficient energy +
​ ​correct
❏ Kinetic Energy Theory​ states that all particles have kinetic energy, increased temp = increased kinetic
energy = moves faster (temp = average b/c particles in substance have diff kinetic Es)
❏ Maxwell-Boltzmann Distribution of Particle Energy
❏ Activation energy is min E in initiate a rxn, lower = faster rate
❏ [Reactants]:​ increased ​freq of collisions​ (increased with increased
[c] in soln, increased pressure in gas, increased SA of solid)
❏ Temperature:​ increased ​freq of collisions ​+ ​# effective collisions
❏ Catalyst:​ increased reaction rate by lowering Ea (alternate
❏ Equilibrium state​ when forward reaction and reverse reaction at same rate
❏ Dynamic​ b/c fwd/rvs reactions still occuring
❏ In closed system​ b/c no matter exchange with surroundings
❏ [c] reactants/products constant​ but not necessarily equal
❏ No change in macroscopic properties​ b/c same observable
❏ Either direction​ (starts with reactants or products)
❏ Equilibrium position​ is proportion of reactions/products
❏ Mostly products = lies to right; mostly reactants = lies to left
❏ Calculating [c] at Equilibrium
1. Balanced chemical equation + state conditions (ex. at SATP)
2. RICE table
3. LF Calculation
4. Solve for ‘x’ (write --- if initial is zero, choose x for change with stoich coefficient of 1)
5. Find [product]eq
6. % Reaction = [completion] × 100%
❏ % rxn > 50% = products favoured
❏ % rxn < 50% = reactants favoured
Equilibrium Constant ​is ratio of reactants to products
(constant) at particular temperature at eqm
(UNITLESS), in terms of [c] in mol · dm3
K c = [C]
[A] [B]b
❏ Larger K: more products at eqm (lies to the right)
❏ Smaller K: more reactants at eqm (lies to the left)
❏ K > 1: favours products (virtually to completion when K >>> 1)
❏ K < 1: favours reactants (basically doesn’t go forward when K <<< 1)
❏ Solids and liquids​ have constant [c] b/c density + molar mass not variables so ignore for Kc
Changing Kc → for ​reverse rxn​ is reciprocal of Kc (with same conditions), when ​doubling rxn
coefficients​ square Kc (tripling cubes, etc), for ​halving ​square root Kc, ​adding 2 rxns​ multiply Kc of both
The ​Reaction Quotient​ is at any time in rxn (when Kc=Q, rxn at eqm), used to predict direction rxn needs
to shift to achieve eqm (larger=greater [products])
Q = [C]
[A] [B]b
❏ Kc then Q (alphabetical order), arrow points to shift direction
❏ Kc <-- Q ​= Q has greater [products] than needed so rvs reaction favoured (left shift)
❏ Kc --> Q ​= Q has lower [products] than needed to fwd reaction favoured (right shift)
Le Chatelier’s Principle​: if system at eqm disturbed, system responds to eliminate the stress
❏ Draw arrow on equation and follow it (→ or ←)
❏ Ask: ​shift?​ (Q to Kc) ​is Kc affected?​ (only temperature)
❏ Concentration​: increase = shift to consume added rxt/product; decrease = shift to replace removed
❏ Temperature​: increase = shift to consume added thermal E; decrease = shift to replace removed
thermal E *​affects Kc (right = greater, left = smaller​)
❏ Volume:​ increase (less pressure) = shift to side with greater total entities; decrease (more pressure)
= shift to side with fewer total entities
❏ Catalyst: increases rate of rxn but no effect on Kc or eqm position
❏ The Haber Process​: production of NH3, Fritz Haber, Germany before WWI, exothermic
N2(g)+3H2(g) ⇌ 2NH3(g) ΔH =-93KJ/mol
❏ [c]:​ in molar ratio 1N2:3H2 (NH3 removed as formed)
❏ Pressure​: high ( ≈ 2 × 107 P a)
❏ Temperature​: moderate (~450ºC, too low is too slow economically)
❏ Catalyst​: finely divided iron with Ag/Mg oxides (to increase rxn rate)
❏ Brønsted-Lowry Theory​: acid is a proton donor, base is a proton acceptor
❏ Conjugate acid-base pair​ differ by 1 proton (c. base remove H+, c. acid add H+)
❏ Amphiprotic​: substance that can be p+ donor or a p+ accepter (not strong acid/base)
❏ Ionic compound​: first dissociate the salt, then show acid/base formation
❏ Autoionization of water: H2O is amphiprotic, contains small but equal amounts H+ and OH- (1 in
10 million each), some of each needed in acidic/basic soln so Kw =/ 0
❏ Kw​ is the equilibrium constant for water, Kw = [H30+][OH-] = ( 1 × 10−7 )( 1 × 10−7 )
❏ Since Kw = ( 1 × 10−14 ), pKw = -log(Kw) = -log( 1 × 10−14 ) = 14 (pOH+pH)
❏ pH Scale​ between 0-14 (acidic to basic); ​pH = -log[H3O+] ​and [H 3 O +] = 10−pH ​(same but with [OH-]
for pOH), sig digs for pH is number of decimal places
❏ Inverse relationship: increase pH by 1 = decrease [H3O+] a factor of 10
❏ Determined by ​strength ​(% ionization) and ​concentration ​( mol · dm−3 )
❏ 1.00 mol · dm−3 HCl=concentrated strong acid; 0.10 mol · dm−3 HCl=dilute strong acid
❏ Find [H3O+] from [Acid] ex. 0.10 mol · dm−3 HCl × 1.00 = 0.10 mol · dm−3
❏ Weak Acid/Base​ ionize < 50% (K <<< 1, in equilibrium)
❏ Strong Acid/Base ​ionize 100% (K >>> 1, not in equilibrium)
❏ Acids 7 taps: nitric, sulfuric, hydrochloric, perchloric, chloric, hydrobromic, hydroiodic
❏ Bases: all soluble hydroxides (alkali metal hydroxides, heavier alkaline earth metal hydroxides
Ca(OH)2, Sr(OH)2, Ba(OH)2
❏ Strong acid​ has weak attraction for p+ (donor), ​stronger acid​ = weaker conjugate base
❏ Strong base​ has strong attraction for p+ (acceptor), ​stronger base​ = weaker c. acid
❏ Equilibrium favours the reaction that moves the proton to the stronger base​ (thus rxn in
opposite direction since an anion)
❏ Features of Strong and Weak Acids/Bases
❏ Electrical Conductivity ​- strongs have more mobile ions since ionize ~100%
❏ Rate of Reaction​ - strongs higher, nature of p+ transfer (more attraction b/n base/p+)
❏ pH​ - stongs more extreme pH, weaks have moderate pH around 7 (a bit above/below)
❏ Physical Properties: acids ​sour, sticky, sharp​; bases b
​ itter, slippery, odourless ​(except NH3)
❏ Chemical Properties
❏ Bases:​ metal oxides, metal hydroxides, metal carbonates, metal bicarbonates, ammonia; ​alkalis
dissolve in water to produce OH- (soluble → many other bases are not)
❏ Acids​: react with metals, bases, and carbonates to form salts
acid + metal → salt + hydrogen ​(single displacement)
acid + base → salt + water ​(neutralization)
acid + carbonate → salt + water + carbon dioxide ​(type of acid-base rxn)
❏ Metal Corrosion​: acids cause metal to corrode b/c ​anion​ of acid displaced, forming salt → metals higher
on activity series more susceptible to corrosion, metals below H not corroded by acids (not reactive enough
to displace acidic H+)
❏ Spontaneously Decompose: ​carbonic acid​ (H2CO3) to CO2 and H2O upon formation, sulfurous acid
(H2SO3) to SO2 and H2O, ammonium hydroxide (NH4OH) to NH4 and H2O
❏ Indicators​ change colour reversibly according to
[H3O+], weak acids/bases whose conjugate has
different colour (litmus, methyl orange,
❏ Universal indicator​ is a mix of many
indicators, many colours so can ~pH
❏ Acid Decomposition​ is all processes by which acidic
components leave atmosphere (gas or precipitate)
❏ Rain is naturally acidic (pH 5.6) b/c
contains dissolved CO2 which reacts with water to form carbonic acid ​CO2 + H2O ⇌ H2CO3
then ​H2CO3 ⇌ HCO3- + H+
❏ Effects of acid rain: ​eutrophication​ (nitrates, algal bloom, Al(OH)3 in rocks reacts with H+ to
form toxic Al3+; ​leaching​ causes minerals in soil to become soluble so washed away (less Mg2+
= less chlorophyll); ​erosion of buildings​ (CaCO3 rxts with acids to form soluble salts, cracks stone
CaCO3 + H2SO4 → CaSO4 + CO2 + H2O​); ​respiratory problems​ due to sulfate/nitrate
precipitates, fish die when pH lake < 5
❏ Production of sulfur and nitrogen oxides
❏ Reduction of sulfur and nitrogen oxides:​ catalytic converters, exhaust gas recirculation,
switch to renewable E, liming lakes to neutralize acidity
❏ Redox Reactions​ is the complete or partial transfer of electrons, coupled reaction
❏ Oxidation​ is loss of electrons (more positive because breathing is good)
❏ Reduction​ ​is gain of electrons (more negative because reducing)
❏ Agents: ​oxidizing agent​ is reduced substance (gains electrons, oxidizes other substances), ​reducing agent
is oxidized substance (loses electrons, reduces other substances)
❏ Oxidation State​ is assigned value that measures electron control/possession of atom in cmpd
1. Zero for​ pure elements
2. Monoatomic ions​ have oxidation state of their charge
3. Sum of all in ​neutral compound​ = zero
4. Sum of all in a​ polyatomic ion​ = charge
5. In compounds, ​group 1​ is +1, ​group 2​ is +2, ​aluminium​ is +3
6. Hydrogen​ is +1 (unless with a metal)
7. Oxygen​ is -2 (unless in peroxides where it is -1 or with fluorine, then +1 or +2)
8. Fluorine​ is -1
9. Chlorine​ is -1 (unless with oxygen or fluorine)
❏ Half equations​: redox reactions broken into oxidation and reduction half-reactions
❏ Electrons on right for oxidation (always write oxidation
first to keep consistent), ​not​ in eqm
❏ Can use table 24 in data booklet (forward rxn is
reduction​, reverse is oxidation)
❏ Balancing in Solution​: only possible in acidic solutions b/c need
H+ and H2O to balance
1. Assign oxidation states to determine what’s oxidized and reduced
2. Write separate half equations
3. Balance all elements except O and H in each half equation
4. Balance O using H2O
5. Balance H using H+
6. Balance charge with electrons and multiply equations to make electrons equal
7. Add equations together and cancel identical species
❏ Activity Series​ is a list of metals (and C and H) in order of relative reducing agent strength
❏ Higher on series = better reducing agent b/c more reactive metals lose e- more easily
❏ Determined experimentally with single displacement reactions
❏ Non-metals more reactive up a group (gain e- more easily so stronger oxidizing agents)
❏ Winkler Method​ used to measure biological oxygen demand (BOD), a measure of dissolved oxygen in
ppm required to decompose organic matter in water biologically over a set time
❏ Water saturated with oxygen so initial [O2] known, measured volume incubates for 5 days while
microorganisms oxidize organic material, manganese(II) salt added in XS, can calc [O2] from amt
1. Under alkaline conditions, Mn2+ oxidized by remaining oxygen
2Mn2+ (aq) + 4OH- (aq) + O2(aq) → 2MnO2(s) + 2H2O(l)
2. Potassium iodide added and oxidized by manganese(IV) oxide in acidic solution
MnO2(s) + 2I-(aq) + 4H+(aq) → Mn2+(aq) + I2(aq) + 2H2O(l)
3. Iodine released then titrated with standard sodium thiosulfate solution
I2(aq)+2S2O32-(aq) → S4O62-(aq)+2I-(aq)
❏ Electrochemical Cells:
❏ Voltaic Cells
❏ Cathode - positive; Anode - negative
❏ Metals higher on activity series
at anode (better reducing agent
is oxidized)
❏ Spontaneous redox reactions
❏ Ion movement through ​salt bridge
neutralizes charge build up and
maintains potential difference
❏ Electrodes are ​not inert
❏ Change potential difference by changing
metals/ions or [c] of metal ions in
❏ Convention to write anode on left, thus electrons flow to right
❏ Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
❏ Aqueous next to salt bridge
❏ Electrolytic Cells
❏ Cathode - negative; Anode - positive
❏ Ionic: nonmetals at anode (oxidized/lose
electrons to go from anion to neutral
atom) and metals at cathode
(reduced/gain electrons to go from
cations to neutral atoms)
❏ Non-spontaneous redox reactions
❏ Conducting wires connect to external power
❏ Anions to anode b/c anode +ve, cations to
❏ Electrodes are​ inert
❏ Electrolysis used to obtain reactive metals (like Na) from their common ores
❏ Colours of ions: Cu2+ is blue, Cr3+ is green, chromate and
dichromate are orange (O in CrO42- and CrO72-), Mn2+ is pale
pink, MnO4- is purple, Fe2+ is green, Fe3+ is brown
❏ Homologous series​ is series of compounds with same functional group + similar chemical properties:
methanol, ethanol, propanol, butanol (alcohols, -OH group), grow by -CH2- group
❏ Similar reactivity since have same functional group
❏ Graduation in physical properties
❏ Can be represented by same general formula ( C n H 2n+1 OH for alcohols)
❏ Primary, secondary, and tertiary compounds
❏ Structural isomers ​are different arrangements of same atoms
❏ Drawing organic compounds
❏ Empirical, molecular, and structural formulas
❏ Full structural formul​a shows every bond and
atom (using 90º and 120ª angles)
❏ Condensed structural formula​ omits bonds where they can be assumed
❏ Stereochemical formula​ shows relative positions of atoms around carbon in 3D
❏ Naming organic compounds
❏ Precedence of classes, functional groups, naming with prefixes/suffixes
❏ Carboxylic acids
❏ Esters
❏ Amides
❏ Aldehydes
❏ Ketones
❏ Alcohols
❏ Amines
❏ Alkenes
❏ Alkynes
❏ Alkanes
❏ Cycloalkanes
❏ Ethers
❏ Benzene/aromatics
Substituents/side chains
❏ Halogenoalkanes
❏ Amine
Functional Group
❏ Carboxyl
❏ Ester
❏ Carboxyamide
❏ Carbonyl
❏ Carbonyl
❏ Hydroxyl
❏ Amine
❏ Alkenyl
❏ Alkynyl
❏ Alkyl
❏ Cycloalkyl
❏ Ether
❏ benzene ring/phenyl
❏ Halogeno
❏ Amine
-oic acid
iodo, chloro, bromo, fluoro
❏ Benzenes​:
❏ Bond lengths​ - all C-C bonds equal/intermediate in length (1.5)
❏ Enthalpy of hydrogenation​ - more stable than predicted, delocalization minimizes repulsion b/n e❏ Type of reactivity​ - reluctant to undergo addition (substitution more likely) b/c addition would
disrupt cloud of delocalized electrons (not energetically favourable)
❏ Isomers​ - only 1 isomer exists of cmpds such as 1,2-dibromobenzene b/c symmetrical
❏ Physical properties of organic compounds​: made of hydrocarbon skeleton and functional group, longer C
chain has higher BP (stronger LDFs), branching = lower BP (less tightly packed, weaker IMFs), polar =
soluble, longer non-polar chain = insoluble
❏ Most to least volatile, increasing BP/IMF strength​:
alkane>halogenoalkane>aldehyde>ketone>alcohol>carboxylic acid
❏ Reactions of Alkanes
❏ Combustion​ with O2 → ​complete​ forms CO2 and H2O, ​incomplete​ also
forms CO and C
❏ Substitution with halogens in presence of UV light​ to form haloalkane
❏ Mechanism of the chlorination of methane (​free radical
substitution)​ , homolytic fission = each atom retains an electron
forming 2 free radicals (reactive, unpaired electrons)
❏ Initiation = produce Cl free radical
❏ Propagation = free radical on each side
❏ Termination = 3 ways to get rid of free radicals
❏ Reactions of Alkenes
❏ Hydrogenation with Ni(s) catalyst and 150ºC​ → add hydrogen to
❏ Halogenation​ to produce dihalogeno cmpds by adding halogen (quickly
at SATP, no UV radiation, loss of colour of reacting halogen)
❏ Addition of hydrogen halides​ to produce halogenoalkanes, HI>HBr>HCl
(quickly in soln at SATP)
❏ Hydration reactions with concentrated H2SO4 catalyst in presence of H2O as steam​ → adding
water to produce ​alcohols​ (ethanol is an important industrial
❏ Polymerization​ → produce addition polymers (addition
reactions with itself), monomer to polymer
❏ Reactions of Alcohols
❏ Combustion​ to produce CO2 and H2O (complete), longer
hydrocarbon chain = more energy released
❏ Oxidation with acidified solution of potassium dichromate(VI) and heat​ with distillation or reflux,
orange to green from Cr(VI) to Cr(III), write +[O] or ​H+/Cr(VIO, heat​ above arrow, oxidizing
agents selectively oxidize C atom attached to the -OH group, distillation/reflux = boiling processes
❏ Primary alcohol → ​(distillation) ​aldehyde → ​(reflux)​ carboxylic acid
❏ Secondary alcohol → ​(reflux)​ ketone
❏ Tertiary alcohol → no reaction
❏ Esterification with concentrated H2SO4​ react carboxylic acid + alcohol to produce ester + water
(type of nucleophilic substitution), ​methanoic acid + ethanol → ethyl methanoate + H2O
❏ Reactions of Halogenoalkanes
❏ Nucleophilic substitution​ → substitute a ​nucleophile​ (loves
+ve), e- rich species that is attracted to e- deficient parts of a
molecule, halogenalkanes react with alkalis to from alcohol
and a salt, ex. ​chloromethane + NaOH → methanol + NaCl
❏ Reactions of Benzenes
❏ Electrophilic substitution​ → substitute an ​electrophile​ (loves
-ve) such as NO2+, H+, to protect stability of arene ring
❏ With concentrated H2SO4, 50ºC for nitric acid HNO3 (subs in NO2, remainder as H2O)
❏ With AlCl3 in dry ether for Cl2 (subs in Cl+ and remainder as HCl)