The Student Room Group

Chemistry OCR A

Module 2
Proton (p+) has relative charge 1+ and relative mass of 1. Electron (e-) has relative charge 1- and relative mass of 1/1836. Neutron (n) has relative mass and relative charge of 1.

Atomic number (Z): p+ number. Mass number (A): p+ + n. Isotope: atom of an element, different number of neutrons, different physical property, same chemical property (due to same e- number). Cation: atom lose e-, p+ > e-. Anion: atom gain e-, p+ < e-.

Standard isotope: C12. 12 atomic mass unit. Standard mass = 1/12th mass of atom of C12. Relative isotopic mass: mass of isotope relative to 1/12th mass of atom of C12. Relative atomic mass: weighted mean mass of atom of an element relative to 1/12th mass of atom of C12. Weighted mean mass: relative isotopic mass * % abundance.

Binary compound: contain 2 elements only. Common ions: NH4+, OH-, NO3-, NO2-, HCO3-, MnO4-, CO3^2-, SO4^2-, SO3^2-, Cr2O7^2-, PO4^3-. Diatomic molecule: 2 atoms covalently bonded together. States: (s), (l), (g), (aq).

Amount of substance (n): number of particles of __, measured in mole. Mole is 6.02 * 10^23 particles. Avogadro constant: 6.02 * 10^23 mol^-1, number of particles per mole of C12.

Example: A question may say how many molecules of H2 does 2.5 moles of C6H12O6 contain? C6H12O6 contains 6 moles of H2 molecules (notice H2, not H atoms, which would be 12 moles. So you would do 6 * 2.5 moles * 6.02*10^23.

Molar mass: mass per mole of substance (gmol^-1). m (mass) = mole (n) * Molar mass (M).

Molecular formula: number of atoms of each element in a compound. Empirical formula: simplest whole number ratio of atoms of each element in a compound.

For empirical formula: find moles of each element using mass/Mr. Then divide all the moles by the smallest mole number. If elements are given as percentages, e.g 13% H, treat it as 13g.

Relative molecular mass: weighted mean mass of a molecule of a compound relative to 1/12th mass of atom of C12.
Relative formula mass: weighted mean mass of formula unit of a compound relative to 1/12th mass of atom of C12.

1cm^3 = 1ml. 1dm^3 = 1000 cm^3 = 1L.
Moles (n) = concentration * volume. If concentration is mol/dm^3, and question wants answer in g/dm^3, you can use mass = moles * Mr: mol/dm^3 * Mr -> g/dm^3.

Standard solution: solution of known concentration.
Molar gas volume (Vm): at stated temperature + Pressure, volume per mole of gas molecules. At RTP (298K, 101KPa), 1 mole of gas = 24dm^3 = 24000 cm^3.
Volume = mole (n) * Vm.

For gases not at RTP, use ideal gas equation. Assume random motion of particles, elastic collision, negligible size, no intermolecular forces.
P (Pa) x V (m^3) = n * R * T(K)
Unit conversions: cm^3 -> m^3 (x10^-6). dm^3 -> m^3 (x10^-3). Degrees -> Kelvin (+273). Kpa -> Pa (x10^3).

Stoichiometry: ratio of moles. E.g 4Al + 3O2 -> 2Al2O3. If I have moles of Al, and want to find moles of Al2O3, use (x want/divide have). Moles Al *4/2.
% yield: actual yield/theoretical yield x 100.
Theoretical yield: complete conversion reactant to product. Reasons for not getting theoretical yield: incomplete reaction, side reaction, reactants stuck on apparatus. Actual yield normally is lower, but may be higher due to water.

Limiting reagent: reactant not in excess, used up first, reaction stops. Whenever you have a reaction, and want to find moles of product, calculate moles of all the reactants, and use the stoichiometric ratio between reactants to find the limiting reagent.

Atom economy: Sum of Mr desired product (include the stoichiometric number)/sum of Mr of all product *100. High atom economy for more desired product, reduce waste product, more sustainable, greater use of natural resource.

Bronsted-Lowry acid: H+ donor. Bronsted Lowry base: H+ acceptor. Strong acid: completely dissociate in (aq) to release H+. Weak acid: partially dissociate in (aq). Base: neutralize acid to form a salt. Metal oxides, metal hydroxides, metal carbonates.Salt: H+ in acid replaced by metal ion or ammonium ion. Alkali: base dissolve in water, release OH-.

Titration: measure volume of 1 solution that react with another. To find concentration, find purity, identify unknown solution. Volumetric flask: measure standard solution accurately. Pipette/burette + drop (bottom of meniscus at eye level). Repeat titration for concordant results to find mean (within 0.1 dm^3).

Oxidation number: set of rules, e- used in bonding of elements. Pure element: 0. Based on electronegativity: more electronegative atom gain e-, less electronegative atom lose electron.
Systemic name. Example iron(II) = Fe^2+.
Oxidation = gain O2 = lose e- = increase oxidation number.
Reduction = lose O2 = gain e- = decrease oxidation number.

Redox reaction: reduction with oxidization. A species is oxidized, with that electron being transferred to be species that is reduced. Oxidation is always present with reduction (as that electron has to be accepted by something).
Reducing agent helps another species be reduced (gain e-), so the reducing agent must be oxidized. For the similar reasoning: Oxidising agent is reduced.

Shell: energy level. 2n^2 (number of electrons in each shell). Made of atomic orbitals. Atomic orbital: region of space around nucleus, hold 2- with opposite spins. S orbital: spherical shape. P orbital: dumbbell shape.

Subshell: group of orbital of same type. Shell 1: 1s subshell. Shell 2: 2s and 2p subshell. Shell 3: 3s, 3p, 3d subshell. Shell 4: 4s,4p,4d,4f subshell.
S subshell has one S orbital. P subshell has three p orbitals. d subshell have five d orbitals. f subshell has seven f orbitals.

Rules: fill lowest energy level orbital first, all orbitals in a subshell has 1 e- before pairing, an orbital can only hold 2e- with opposite spins.

General: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p. Note this is only a general trend in orbital energy levels. Different elements may have different orbital energy levels, or this may change if the orbital has an electron. As you will need to write the electron configuration for elements up to 36 electrons, this is useful for most elements.

Exceptions: Cr is [Ar]4s1 3d5. Cu is [Ar] 4s1 3d10. As different elements have different number of protons (different electrostatic attraction to e-), different number of electrons (different electrostatic repulsion between e-), it is not surprising that different element orbitals will have different energy levels and order of filling.

Shorthand: use previous noble gas. [He], [Ar]. Periodic table has s, d, p block. For d block: 4s fill first, 4s empty first (for when forming ions).

Ionic bonding: electrostatic attraction between oppositely charged ions (cations and anions). Electrostatic attraction in all direction. Strength depend on ionic charge and ionic radius. Smaller ionic radius, greater relative charge. Form giant ionic lattice. Repeating pattern of oppositely charged ions.

Has high melting/boiling point. To dissolve: break ionic lattice, water is polar, attracted to ions. Not conduct electricity in solid, ions fixed position. Conduct in liquid, ions are mobile.

Covalent bond. Electrostatic attraction between shared pair of e- and nuclei of both atom. Overlap of atomic orbitals. Localized attraction. Lone pairs, pair of electron not shared.

3S shell has expansion of octet, by promotion of e- to orbitals. Dative covalent (coordinate bond): covalent bond, shared pair of e- from 1 bonding atom. E.g N in NH4+. Shown in diagram as -> bond.

Average bond enthalpy: measure covalent bond strength.
E- pair repulsion theory: e- around central atom determine shape, e- pair repel furthest apart, greatest stability, minimize repulsion.
Bond in plane (-), bond out of plane (solid wedge), bond into plane (dotted line).

Tetrahedral. 4 bond pair, 0 lone. 109.5 degree bond angle. CH4.
Pyramidal. 3 bond, 1 lone. 107 degree. NH3.
Non-linear. 2 bond, 2 lone. 104.5. H2O.
Multiple bond = bonding region.
Linear. 2 bonding region. 180 degree. CO2.
Trigonal planar. 3 bonding region. 120 degree. BF3.
Octahedral. 6 bonding region. 90 degree. SF6.

Electronegativity - affinity of atom to attract bonded pair of e- in covalent bond. Greater nuclear charge, smaller atomic radius, greater electronegativity. Increase up and across to F (most electronegative). Does not include noble gas, does not normally form bonds.

Non-polar bond. Bonded e- shared equally. Similar electronegativity. Polar bond - different electronegativity. Has permanent dipole dipole interaction. Dipole - separation of charge. Whether something is polar, depends on shape of molecule: dipoles may cancel out or reinforce each other.

Large difference in electronegativity, ionic bond forms. Small difference in electronegativity, covalent bond forms. Not a clear divide between ionic and covalent. E.g HCl has covalent bond, but has some ionic properties (dissocistes in aq soluton).

Internolecular forces: induced dipole dipole (all molecule have), permament dipole dipole (polar), hydrogen bonding.

Induced dipole dipole: movement e- produce dipole charge, induce adjacent molecule. Greater e-, greater induced. Greater surface area of contact, greater induced.

[link to organic. For structural isomers of a hydrocarbon, the less branched structural isomer will have greater SA contact than the more branched one, has higher boiling point as more energy required to break].

Simple molecular substance: form simple molecular lattice. Held by weak intermolecular forces. Low boiling point. Only contain localised e-, no conduct electricity.

Like dissolves like. Non-polar + non-polar can interact, as induced dipole interactions form between them. Polar + polar can interact due to dipole charges. [polar + non-polar. Polar too strong for non-polar to interact with]. Useful solvents contain both polar + non-polar, e.g ethanol has non-polar carbon chain and polar OH, to interact with different solutes.

Hydrogen bond: type of permanent dipole interaction. Has H bonded to an electronegative atom with a lone pair of electrons, such as O, F, N. Causes ice to be less dense than water, form tetrahedral lattice with air gap. Responsible for molecules having high boiling points relative for simple molecular substances.

Practicals list (mostly)
Determine relative atomic mass. Sample in mass spectrometer -> vaporize + ionize to cation -> accelerate. Heavier ion more difficult to deflect than lighter ion -> mass/charge ratio.

Analysis - empirical formula. n = m/Mr of element. Divide by smallest. If %, convert to mass. [if].

Hydrated salts. Water of crystallization - water part or crystalline structure. CuSO4.5H2O (dot to show water of crystallisation, unable to find symbol for it) -> CuSO4 + 5H2O.
Weigh empty crucible -> add hydrated salt and reweigh -> heat proof mat, bunsen burner, tripod, pipe-clay triangle. Heat strongly. Re-weigh crucible (now be anhydrous salt).
Calculate n of anhydrous salt. Calculate n of water. Simplest whole number ratio.
Assumptions:
All water is evaporated. Confirm by heating, and reweigh, mass no change.
No further decomposition of salt (you know the formula of the salt).

Determine relative molecular mass for volatile liquid. Assume liquid is pure, with boiling point < 373K. Add liquid to syringe using needle -> weigh syrine -> transfer to gas syringe through self-sealing rubber cap -> reweigh syringe -> water bath at 373K. Record volume. Assume 101000 Pa.
Find n using pV/RT. Find Mr using mass/moles.

Identify group 2 metal. Mass balance, weigh metal. To measuring cylinder add excess HCl. Transfer to flask connected with gas syringe. Add metal. Record volume of H2.
X + 2HCl -> XCl2 + H2. Volume = n x Vm. Mr = mass/moles.

Prepare standard solution. Weigh solid on mass balance using container -> transfer to beaker (rinse container into beaker) -> add distilled water, stir with glass rod until dissolve -> transfer to volumetric flask (rinse beaker with distilled water into volumetric flask) -> use pipette to fill until graduation line, with bottom of meniscus at eye level. -> add stopper and invert volumetric flask for constant concentration.

Acid-base Titration. One solution to conical flask on white tile with indicator (phenolphatelin. Pink alkali, colourless acid). Fill burette - run solution to remove air bubble. Record initial reading. Add solution + swirl to endpoint (colour change). Record final reading. Repeat for concordant result to calculate mean.
The moles of the solution in the conical flask remains the same, so you can add more distilled water without changing the titre volume.

For calculation: use n = cv, using mean titre volume. Use stoichiometric ratio to find mole of other solution. If question said a standard solution of 250cm3 was prepared, and a sample of 30cm^3 was used in titration, and you want to find mass of original solution, you will need to multiply moles by 250/30.
Module 3
Periodicity - repeating trend of properties across periods. e- configuration, ionization energy, melting point, structure.
S and p subshell fill same way. Ionisation energy - energy required to remove 1 mole of electrons from 1 mole of gaseous atom/cation to form 1 mole of gaseous _ cation. E.g Ag(g) -> Ag+(g) + e-.
Depend on atomic radius, nuclear charge, e- shielding. Successive - endothermic. Big jump in energy, to next shell (you can deduce number of outer e-).
Down group: increase atomic radius, increase e- shielding, reduce electrostatic attraction to outer electron, decrease ionization energy.
Across period: same electron shielding, increase nuclear charge, smaller atomic radius, greater electrostatic attraction, increase ionisation energy.
Exceptions across period increasing (drops in ionisation energy). 2s2 -> 2p3 or 3s2 -> 3p3: p orbital higher energy, less energy required to remove electron. 2p3 -> 2p4 or 3p3 -> 3p4: paired electron repel, less energy required to remove.
Giant metallic lattice. Metallic bonding - electrostatic attraction between cation (fixed) and delocalised electron. Conduct electricity.
Giant covalent lattice: high melting point. Break covalent bond.
Group 1 to 4 are giant structures . Groups 5 to 8 are simple structures.
Group 2 metals are reducing agent. Ca + 2H2O -> Ca(OH2) + H2.
CaO + H2O -> Ca(OH)2. [notice the atom economy]. For hydroxides: down group, more soluble, more dissociate OH-, higher pH. Used in agriculture (neutralise acidic soil, Ca(OH)2), used as antacid (CaCO3).
Group 7. Boiling point increase down group, more e-, more induced dipole interaction. An oxidising agent.Reactivity decrease down group (increase atomic radius, increase e- shielding, reduce electrostatic attraction to e- on another atom). Displacement reactions:
Cl2 + 2Br- -> 2Cl- + Br2 (green to orange).
Cl2 + 2I- -> 2Cl- + I2 (green to brown in water, or violet in cyclohexane).
Br2 + 2I- -> 2Br- + I2 (orange to brown in water, or violet in cyclohexane).
Cl2 + H20 -> HClO + HCl. HClO <-> H^+ + ClO^-. ClO- kills bacteria. Greater solubility in cold, dilute NaOH.
Cl2 + 2NaOH -> NaCl + NaClO + H2O. Negatives: toxic, respiratory irritant, chlorinated hydrocarbon cause cancer.
Tests. Carbonate -> sulfate -> halide.
CO3^2- + 2H+ -> CO2 + H2O. CO2 + Ca(OH)2 -> CaCO3(s) + H2O. White precipitate.
Ba^2+ + SO4^2- -> BaSO4 (s). White precipitate.
Ag+ + X- -> AgX (S). AgCl (white, soluble dilute NH3). AgBr (cream, soluble concentrated NH3). AgI (yellow, insoluble concentrated NH3).
Use HCl (carbonate) -> Ba(NO3)2 for sulfate -> AgNO3 (halide).
BaCO3, Ag2SO4, Ag2CO3 also precipitate, hence that order.
NH4+ test. Add warm NaOH. NH4+ +OH- -> NH3 + H2O. Heat as NH3 form hydrogen bond with water. Turn damp litmus paper blue, from red.
Enthalpy. ΔH = H(product) - H(reactant).Change in energy between system and surroundings. Exothermic: negative ΔH. H(reactant) > H(product). Endothermic: positive ΔH. H(product) > H(reactant).
Standard condition: 100KPa, 298K, 1 mol dm^-3, standard state.
ΔHreaction = enthalpy change reaction in molar quantity in stated chemical equation under standard condition, reactant/product standard state.
ΔHformation = 1 mole of compound formed from its elements under standard condition, reactant/product standard state.
ΔHcombustion = 1 mole of substance react completely with oxygen under standard condition, reactant/product standard state.
ΔHneutralisation = acid + base react to form 1 mole of water under standard condition, reactant/product standard state.
Q = m( mass) * c (specific heat capacity) *ΔT( final - initial temperature). From experiment, ΔHcombustion less exothermic then expected due to heat loss to surroundings, incomplete combustion, non-standard conditions, evaporation from wick. Use draught screen, O2 cylinder.
On cooling curve, extrapolate to when reactants were added together.
Calculate Q in J, divide by 1000 into KJ, divide by moles to find enthalpy, use whether temperature increased to add negative sign to show exothermic, or + sign if temperature decreased (endothermic). [remember to use moles of limiting reagent. remember the enthalpy definitions. if a question had you calculate the enthalpy change of a reaction that produced 3 moles of water with an acid and base, then had you calculate the enthalpy change of neutralisation, you would had to divide your previous answer by 3].
Average bond enthalpy: break 1 mole of covalent bond in gaseous molecule. Average - depend on chemical environment. Endothermic - so has + values.
Sum of ΔH reactant - sum of ΔH product. Must be (g).
Hess Law. It is like displacement. Formation: - reactants + product. Combustion: reactant - product. [remember the enthalpy definitions to help you. for combustion, ensure 1 mole of _ reacts with oxygen. for formation, ensure 1 mole of _ is formed.]
Unknown Hess cycle: add together reactant + product. [from overall reaction identify a reactant or product that is only in one of the reactions you have been given. see the stoichiometric number, and see if it has changed from a reactant <-> product to help start you off].
Rate = change in concentration of reactant or product/time. Increase by catalyst, temperature, surface area, concentration, pressure. Collision - depend on orientation and activation energy. Measure rate by volume gas produced, or mass loss.
Catalyst - alternate reaction pathway with lower activation energy. Homogenous - same physical state. Esterify COOH + alcohol with H2SO4. Photodissociation.
Heterogenous - different physical state. Haber (Fe, for NH3), reforming (alkene into alkane), hydrogenate alkene (423K use Ni), V2O5 for contact process. Allows lower temperature, less energy, less fossil fuel.
Boltzmann distribution. Energy on x axis, molecules on Y axis. Shows particles have different energy levels. Line marks Ea indicates activation energy. Increase temperature, peak shift right and lower. Increase frequency of collision + more particle with energy > activation. Use catalyst: greater proportion of molecules energy > activation.

Dynamic equilibrium. Rate forward = rate backward. In a closed system.
Le Chatelier principle. System at equilibrium, counter act change.
Increase reactant, shift product. Increase product, shift reactant.
Increase temperature, shift endothermic. Decrease temperature, shift exothermic.
Increase Pa, shift to side with lower stoichiometric number of gases. Decrease Pa, shift to side with higher stoichiometric number of gases.
Catalyst - increase rate forward and backward equally.
Equilibrium constant - value for position at equilibrium. Constant at a constant temperature. Kc > 1, shift to product. Kc < 1, shift to reactant. Kc use [], using [products]/[reactants].
Kc does not include (S) as they have a constant concentration, or liquids due to their high concentration there is a negligible change. Only include (aq) and (g).
(edited 1 month ago)
Module 5
Rate = change in concentration/time. Rate proportional to [a]^m^n. Overall order = m+ n. k = rate constant. Determined experimentally. Initial rate when t =0. Use continuous monitoring, gas collection, mass loss, colorimeter (plot calibration curve first).
[a]^2 means if concentration of A tripled, rate would increase 3^2 times.
Concentration time graph. Zero order reaction: Linear, negative gradient. First order reaction: has constant half life. Then k = ln(2)/half life.
Rate concentration graph. 0 order: straight line where y = number. 1st order: linear line through origin, positive gradient. 2nd order: positive curve. Initial rate when t =0 on [] time graph or use clock reaction. Rate = 1/t, where t is time taken for visual change. Assume average rate = initial rate.
Rate determining step is slow step. Involve species stated in rate equation.
k = Ae^(-Ea/RT) links temperature and rate constant. For plotting, use ln(k) = -Ea/R *1/T + ln(A).
Equilibrium constant. Homogenous equilibrium include all. Heterogenous equilibrium include only (aq) and (g). Kp - involve only (g). Mole fraction = moles of X/total number of moles. Partial pressure = total pressure * mole fraction. [remember only temperature changes k value].
Conjugate acid-base pair. Acid in forward direction, has conjugate base in backward direction (linked by H+). Acid can be mono/di/triprotic.
pH = -log[h+]. 10^-pH = [h+]. Ka = 10^-pKa. pKa = -log(Ka). Greater Ka, great shift equilibrium right, greater H+ released.
Strong acid - high Ka, low pKa.
pH weak acid. Ka = [h+][a-]/[ha]. Assume negligible change in [ha], so [ha] equilibrium = {HA] at start. Ignore dissociation of water. As [h+] = [a-], can be written as [h+]^2.
Kw = [h+][oh-]. At RTP, 10^-14. Neutral solution when [h+] = [oh-]. At RTP, neutral is pH 7.
Buffer solution. Minimize pH change when acid/alkali added. Weak acid + salt. [can be made by excess weak acid + base]. Add acid, increase H+, shift to HA. Add base, decrease H+, shift to A-. Buffer effective when [ha] = [a-]. 50% is associated and 50% is dissociated.
Ka = [h+][a-][ha] -> [h+] = Ka * [ha]/[a-] -> [h+] = Ka -> pH = pka.
Blood pH. H2CO3 <-> HCO3- + H+.
pH titration curve. Vertical section = pH change when small acid/base added. Equivalence point: center vertical section, volume acid react completely with base. Acid-base indicator. [ha] colour1 <-> [a-] different colour + H+.

Lattice enthalpy. Enthalpy change 1 mole of giant ionic compound formed from its gaseous ions under standard condition. Bond formation, always exothermic.
Enthalpy of atomisation: form 1 mole of gaseous atom from its element (in standard state) under standard condition. Always endothermic.
Enthalpy of Ea: 1mole of electron added to 1 mole of gaseous atom to form 1 mole of gaseous anion. 1st is endothermic (bond form), successive is exothermic (e- repulsion).
Born haber cycle. 1 pathway: using formation. Second pathway: use atomisation, then ionisation energy to form cation and Ea for anion, then lattice enthalpy.
Enthalpy of solution. 1 mole of giant ionic lattice dissolved in water to form (aq) ions.
Enthalpy of hydration. 1 mole of (g) ion dissolved in water to form 1 mole (aq) ion.
1 pathway: salt + water -> (aq) ions using enthalpy of solution. Other pathway: -Lattice enthalpy + hydration.
Lattice enthalpy/hydration depend on ionic charge and radius. Increase charge density, increase electrostatic attraction to delocalised e- (lattice) or to water (hydration).
Entropy. Dispersal of energy within chemical system. Entropy (g) > (l) > (s). In JK-1mol-1. Entropy change = sum product - sum reactant.
Feasibility if ΔG < 0 [product has less energy then reactant. less energy is more stable. favorable].
ΔG = ΔH - TΔS. You can plot a graph of Temperature on X axis, and ΔG on Y axis to visually see feasibility.
[this is thermodynamic feasibility. it does not account for kinetics: reaction rate, activation energy. even if a reaction is feasible, it may have a very low rate and wont occur].
Oxidising agent - is reduced. Reducing agent - is oxidized. Redox titration use burette/pipette/ standard solution, conical flask.
MnO4- (oxidizing agent): read from top of meniscus due to purple colour.
I2/S2O3- for oxidizing agent. 2I- + oxidizing agent -> I2 + 2e-. I2 react with S2O3-. I2 - orange brown. Near the end point add starch (turn blue-black).
Electrochemical cell. Half cell. Metal (aq)/metal (s) or metal 1 (aq), metal 2 (aq)/ inert metal (Pt). Standard electrode potential - emf of a half cell connected to a standard hydrogen half cell under standard condition. Salt bridge: concentrated electrolyte not react with electrode. KNO3. Use High resistance voltmeter.
Electrode potential = E (reduction) - E (oxidation). Less positive electrode is oxidised. More positive is reduced. Limitations: reaction rate, concentration not 1, only for aq equilibria.
Primary cell - non-rechargable. In smoke detector. Secondary cell: rechargable. In lead acid battery, Li+ battery.
Fuel cell: energy produced by fuel react with oxygen. Continuous fuel, not rechargable. Electrolyte remain in cell.
H2 fuel cell - 1.23V, only produce water as waste product.
For acid:
Anode: H2 -> 2H+ +2e-
Cathode: 2H+ + 2e- + 1/2O2 -> H2O.
For alkali:
Anode: O2 + H2O + 4e- -> 4OH-
Cathode: 2H2O + 4OH- + H2 -> 1/2O2 + 4e-.
d block element: element highest energy subshell is d subshell. Transition element: d block element which forms at least 1 stable ion with incomplete d subshell. Sc/Zn not transition element.
Has different oxidative state, catalyst, coloured compound, complex ions.
Complex ion. Ligand form dative covalent bond with central metal ion, use lone pair of electron. Coordination number: number of dative covalent bond. Ligand: species donate lone pair electron to central metal ion.
Shape: tetrahedral, octahedral, square planar (90 degree, has 3d8 in subshell). Square planar has cis-trans. Cis: 2 identical ligand adjacent. Trans: 2 identical ligand opposite.
Octahedral: Has cis and trans. 4 of 1 ligand, 2 of another ligand. Or 2 bidentate and 2 monodentate. Has optical isomerism using 3 bidentate of cis isomer.
Cis-platin. Pt with NH3 on 1 side, and Cl on other side. Square planar, 90 degrees. Anti-cancer, disrupt DNA.
Ligand substitution.
Fe2+ in haem group of hemoglobin, can form dative covalent with oxygen (normal), CO2 (normal), and CO (not normal, lead to suffocation). CO forms an irreversible covalent bond, does not dissociate

Precipitation reaction:
[Cu(H2O)6]^2+ blue. Add OH-: [Cu(H2O)4(OH)2] blue. Excess NH3: [Cu(NH3)4(H2O)2]2+ dark blue.
[Fe(H2O)6]^2+ green. Add OH-: [Fe(H2O)4(OH)2] green.
[Fe(H2O)6]^3+ yellow. Add OH-: [Fe(H2O)3(OH)3] yellow.
[Mn(H2O)6]^2+ pink. Add OH-: [Mn(H2O)4(OH)2] brown.
[Cr(H2O)6]^3+ violet. Add OH-: [Fe(H2O)3(OH)3] gray-green. Add excess OH-: [Fe((OH)6]3- green (aq). Add excess NH3: [Fe(NH3)6]3+ purple (aq).
[Cu(H2O)6]^2+ +4Cl- -> [Cu(Cl)4]2- yellow tetrahedral + 6H2O.
Ion changes
Fe2+ -> Fe3+ with MnO4- (purple) -> Mn2+ (colorless).
Fe3+ -> Fe2+ with I- -> I2(aq).
Cr2O7^2- (orange) -> Cr3+ (green) -> Cr2+ (blue) with Zn.
Cr3+ + hot alkaline H2O2 -> CrO4^2- (yellow)
Cu2+ +3I-m-> CuI (white precipitate) + I2.
Disporportation. Same element oxidized + reduced. [seen in module 3. Cl2 with cold, dilute NaOH]
Cu2O + H2SO4 -> CuSO4 (blue) + Cu (yellow) + H2O.
Cu2O + 2H+ -> Cu2+ + Cu(s) + H2O.
Remember that Fe2+ (aq) means [Fe(H2O)6]2+
(edited 1 month ago)

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