OCR (A) Chemistry A-level Module 2 Notes (PDF)

OCR (A) Chemistry A-level Module 2: Foundations in Chemistry Notes by Amie Campbell www.pmt.education 2.1 atomic structure and isotopes Isotopes: atoms of same element w/ different no of neutrons + different masses Mass number=p + n, atomi...

OCR (A) Chemistry A-level Module 2: Foundations in Chemistry Notes by Amie Campbell www.pmt.education 2.1 atomic structure and isotopes Isotopes: atoms of same element w/ different no of neutrons + different masses Mass number=p + n, atomic number=p Different isotopes react in similar ways: have same number of electrons, chemical reactions involve electrons & number neutrons has no effect on reactions Heavy water: used to control nuclear processes, for heavy water 2​​ 1​H isotopes (deuterium /D) formula can be written D​2​O. slightly higher mpt, bpt + density. If all water were heavy water, would see ice more often since higher mpt Cations: +ve, anions: -ve 2.2 relative masses Mass defect: small amount of mass lost due to strong nuclear force holding together protons/neutrons Relative isotopic mass: mass of an isotope relative to 1/12 mass of an atom carbon-12 Relative atomic mass: weighted mean mass of an atom of an element relative to 1/12 mass of an atom of carbon-12 % abundances isotopes can be found w/ mass spectrometer. Work by: ○ Sample placed in mass spec, vapourised then ionised to form +ve ions ○ Ions accelerated (heavier move more slowly + difficult to deflect so ions of different isotopes separated) ○ Ions detected as a mass to charge ratio (m/z). Greater the abundance, greater the signal ○ For ion w/ +1 charge, ratio equivalent to relative isotopic mass 2.3 formulae and equations Binary compound: contains 2 elements only. To name: change ending 2nd element to -ide. Metals come first in ionic compound names Polyatomic ions: an ion w/ more than one element bonded together: ○ Ammonium NH​4​+ ○ Hydroxide OH​-​, nitrate NO​3​-​, nitrite NO​2​-​, hydrogencarbonate HCO​3​-​, manganate (VII) (permanganate) MnO​4​- ○ Carbonate CO​3​2-​, sulfate SO​4​2-​, sulfite SO​3​2-​, dichromate (VI) Cr​2​O​7​2- ○ Phosphate PO​4​3- Diatomic molecules: H​2​, N​2​, O​2​, F​2​, Cl​2​, Br​2​, I​2 Other small molecules: P​4​ or S​8​ (normal practise to write as S) 3.1 amount of substance and the mole Avogadro constant: number of particles in each mol of carbon-12 One mole: amount of substance that contains 6.02x10​23​ particles ​ Molar mass: mass in g in each mole of a substance (units= g mol​-1​) n=m÷Mr 3.2 determination of formulae Molecular formula: no of atoms of each element in a molecule Empirical formula: simplest whole number ratio of atoms of each element in a compound (important for giant crystalline structures where it would be impossible to use actual numbers- used for metals, some non metals (e.g. C) and ionic compounds) www.pmt.education  Relative molecular mass: compares mass molecule w/ mass carbon-12 Relative formula mass: compares mass formula unit w/ mass carbon-12 Analysis: investigating chemical composition of substance Water of crystallisation: water molecules which are part of crystalline structure When hydrated crystals are heated, bonds holding water w/in crystals are broken + water driven off E.g. CuSO​4​。5H​2​O (s) (blue) -> CuSO​4​ (s) (white) + 5H​2​O (l) Experimentally finding formula hydrated salt: ○ Weigh empty crucible, then add hydrated salt & reweigh ○ Use pipe clay triangle to support crucible on a tripod. Heat for 1 min strong, 3 mins gentle ○ Leave to cool then weigh Accuracy of experimental formula: assumes all water has been lost (solution= heat to constant mass) & assumes no further decomposition (difficult if no colour change) 3.3 moles and volume 1 mol dm​-3​ sol contains 1 mol of solute dissolved in each 1 dm​3​ solution mol=conc x vol Standard solution= solution of a known conc. Prepared by: dissolving exact mass solute in solvent + making it up to an exact volume Can have mass conc (g dm​-3​), mol dm​-3​ → mass dm​-3​: x Mr Molar gas volume V​m​: the volume per mole of gas molecules a stated temperature and pressure (24 dm​3​ mol​-1​ at RTP) RTP: about 20℃ and 101 kPa Assumptions for molecules making up ideal gas: random motion, elastic collisions, negligible size, no intermolecular forces pV=nRT: ○ p= pressure Pa (kPa to Pa x 10​3​) ○ V=vol m​3​ (cm​3​ to m​3​ x 10​-6​ // dm​3​ to m​3​ x 10​-3​) ○ n=mol of gas ○ R=ideal gas constant 8.314 J mol​-1​ K​-1 ○ T= temp in K Experimentally finding a relative molecular mass of volatile liquid (liquid room temp, boils below 100℃): ○ Add sample liquid to small syringe via needle + weigh ○ Inject into gas syringe through self-sealing rubber cap + reweigh small syringe to find mass added ○ Place in boiling water bath (100℃). Liquid vapourises + record pressure ○ Use ideal gas eq to find moles, then use moles & mass to find Mr 3.4 reacting quantities Stoichiometry: ratio of amount/moles each substance in balanced eq Experimentally identifying unknown metal: ○ Weigh sample metal + add to conical flask ○ Add known conc + vol HCl to flask + quickly replace bung www.pmt.education  ○ Measure max vol gas in syringe , work out moles gas then use eq to find moles unknown, use moles + mass to find Mr Theoretical yield: maximum possible product if all reactants converted to product May not achieve theoretical yield because: ○ Reaction may not go to completion ○ Side reactions may take place ○ Purification may lead to loss of product Percentage yield=actual yield/theoretical yield x100 Limiting reagent: reactant not in excess which will be used up first + stop reaction. Calculations must use limiting reagent Atom economy: measure of how well atoms have been utilised Atom economy=(sum molar masses desired product)/(sum molar masses all products) x100 High atom economy: large proportion desired products, few waste products, important for sustainability-make most of resources. Makes industrial processes more efficient, preserves raw material + reduces waste Atom economy only part of sustainability: should use readily available reactants w/ low obtaining costs & depends on percentage yield 4.1 acids, bases and neutralisation Strong acid fully dissociates in aq solution, weak acid partially dissociates in aq solution Alkali: base that dissolves in water releasing OH​-​ ions into solution Neutralisation of an acid: the H​+​ ions react w/ a base to form a salt + water. H​+​ replaced by metal or ammonium ions from base with alkalis, all reactants for neutralisation are aq. Ionic eq: H​+​ + OH​-​ → H​2​O 4.2 acid-base titrations Titration: technique to accurately measure vol one solution reacting exactly with another Can be used to check purity- important for pharmaceuticals etc Volumetric flask typical tolerances: 100 cm​3​ +/- 0.2 cm​3​, 250 cm​3​ +/- 0.3 cm​3 Experimentally preparing standard solutions: ○ Solid weighed + dissolved in beaker using less distilled water than needed to fill volumetric flask ○ Transfer to volumetric flask + last traces rinsed into flask with distilled water ○ Add distilled water dropwise until bottom of meniscus matches up with mark ○ Flask inverted slowly several times to mix, if not titration results will be inconsistent Typical tolerances pipette: 10 cm​3​ +/- 0.04 cm​3​, 25 cm​3​ +/- 0.06. Burette: 50 cm​3​ +/- 0.1 Burette recorded to nearest half division, to 2dp (last 5 or 0) Acid-base titration procedure: ○ Add measured volume one solution to conical flask w/ pipette ○ Add other solution to burette, record initial reading ○ Add few drops indicator to conical flask ○ Run solution from burette into conical flask, swirling it, until it reaches the end point www.pmt.education  ○ Record final reading. ○ 1st titre carried out quickly to get approx, then repeat accurately adding solution dropwise as end point approached. Carry out until two accurate titres are concordant (within 0.1 cm​3​) 4.3 redox Sign for oxidation number placed before the number Oxidation numbers: ○ O -2 ○ H +1 ○ F -1 ○ Ions: charge on ions ○ Special cases: H in metal hydrides: -1, O in peroxides -1, O bonded to F +2 Roman numerals used for compounds w/ elements that form different ions w/ different charges Reduction:gain of electrons (decrease oxd no), oxidation: loss of electrons (increase oxd no) 5.1 electron structure shells=energy levels, energy increases as shell number increases Principal quantum number: the shell number or energy levels number Atomic orbital: region around the nucleus that can hold up to 2 electrons w/ opposite spins (make up shells) S- sphere, P-dumbbell, can have 3 separate orbitals P​x​, P​y​, P​z Orbitals of same type are grouped as sub-shells Two electrons in an orbital must have opposite spins (up/down) Periodic table can be divided into blocks corresponding to highest energy subshell 4s sub shell at lower energy than 3s, so is filled first. Once filled, energy 3d falls below 4s, so it empties before 5.2 ionic bonding and structure Ionic bonding: the electrostatic attraction between positive and negative ions High melting point and boiling point: high temp needed to provide large quantity energy needed to overcome strong electrostatic attraction between ions Melting points higher for lattices with greater charges on ions, because of stronger attraction between ions Many dissolve in polar solvents (e.g. water), they break down the lattice + surround each ion in solution. If made of ions with strong charges, attraction too strong to be broken down www.pmt.education  Solubility depends on relative strengths attractions w/in lattice & attraction between ions + water molecules so predictions of solubility should be treated with caution Solid state: ions in fixed position, no mobile charge carriers so doesn’t conduct electricity liquid/molten: solid lattice breaks down so ions free to move as mobile charge carriers so can conduct electricity Ions in tooth enamel removed in acid conditions, gaps can allow toothy decay to develop. Saliva helps neutralise acidic food + replace ions but not always enough so toothpaste contains fluoride ions 5.3 covalent bonding Covalent bonding: strong electrostatic attraction between a shared pair of electrons and the nuclei of bonded atoms Attraction is localised between electrons In BF​3​, there are 6 electrons in boron’s outer shell→ bonding predictions can’t be made solely on noble gas structure For phosphorus, sulfur and fluorine their outer electrons are in n=3 outer shell, which can hold up to 18 electrons, so more electrons are available for bonding. This means you can get: PF​3​ PF​5​ SF​2​ SF​4​ SF​6​ ClF ClF​3​ ClF​5​ ClF​7​. Different numbers unpaired electrons→ different possibilities of compounds (e.g. sulfur) Double covalent bond: electrostatic attraction is between 2 shared pairs of electrons and the nuclei of bonding atoms e.g. O​2​ or CO​2 Triple covalent bond: electrostatic attraction is between 3 shared pairs of electrons and the nuclei of bonded atoms e.g. N​2​ or HCN Dative covalent/coordinate bond: covalent bond in which shared pair of electrons has been supplied by one of the bonding atoms only. Shown by →arrow, e.g. NH​4 6.1 shapes of molecules and ions Electron pair repulsion theory: ○ Electron pairs surrounding central atom determine molecule/ion shape ○ Pairs repel each other so they are arranged as far apart as possible ○ Arrangement minimises repulsion, so holds atoms in definite shape ○ Different number of electron pairs→ different shape www.pmt.education  Solid line=in plane of paper, wedge=comes out of plane of paper, dotted=goes into plane Lone pair slightly closer to central atom + occupies more space than bonded atom→ repels more strongly than bonding pair. Bond angle is reduced by about 2.5° per lone pair CH​4​= tetrahedral 109.5°, NH​3​ pyramidal 107°, H​2​O nonlinear 104.5° Multiple bonds treated as a bonding region Different shapes/bond angles:* ○ 2: 180°, linear e.g. CO​2 ○ 3: 120°, trigonal planar e.g. BF​3 ○ 4: 109.5°, tetrahedral e.g. CH​4 ○ 6: 90°, octahedral e.g. SF​6 Ammonium ion has 4 bonding pairs so tetrahedral CO​3​2-​ and NO​3​-​ ions are trigonal planar* SO​4​2-​ ions are tetrahedral 6.2 electronegativity and polarity Electronegativity: attraction of a bonded atom for the pair of electrons in a covalent bond Pauling scale (& pauling electronegativity values) used to compare electronegativity Increases upwards + across towards fluorine (F, O, N, Cl most, group 1 least) If electronegativity difference is large, bond becomes ionic rather than covalent Non-polar bond: bonded electron pair shared equally between bonded atoms (when atoms same element/ similar electronegativity) Pure covalent bond: bonded atoms are same element Polar covalent bond: bonded electron pair shared unequally (different atoms with different electronegativity). More electronegative has greater attraction for bonded pair electrons. Polar bonds are polarised with 𝛅+/- signs (dipoles) Permanent dipole: dipole in a polar covalent bond which doesn’t change www.pmt.education  Polar bonds may reinforce one another to produce larger dipole over whole molecule or cancel out: ○ H​2​O: polar- OH bonds have permanent dipole, act in opposite directions but don’t exactly oppose each other, overall O end is 𝛅- and H end is 𝛅+ ○ CO​2​: nonpolar- C=O bonds have permanent dipole, act in opposite directions + exactly oppose each other, overall dipoles cancel Ionic lattices dissolving in polar solvents: water molecules attract +/- ions, ionic lattice breaks down as dissolves, water molecules surround ions, + ions attracted towards 𝜹- oxygen of water & - ions attracted to 𝜹+ hydrogen 6.3 intermolecular forces Intermolecular forces: weak interactions between dipoles of different molecules Induced dipole-dipole interactions (London forces): ○ Exist between all molecules, only temporary ○ Movement electrons produces changing dipole in any molecule, at any instant an instantaneous dipole will exist but position shifts constantly ○ Instantaneous dipole induces a dipole on neighbouring molecule, which induces dipoles on further molecules, they then attract one another ○ More electrons in each molecule→ larger instantaneous + induced dipoles→ greater induced dipole-dipole interactions→ stronger attractive forces (explains increased bpt noble gases) Permanent dipole-dipole interactions: ○ Act between permanent dipoles in polar molecules ○ Mean boiling point of polar molecules is much greater: have both London and permanent dipole-dipole interactions, extra energy needed to break additional permanent interactions, so boiling point is higher Simple molecular substance: made of of simple molecules (small units w/ definite number of atoms). In solid state form simple molecular lattices, held together by weak intermolecular forces, but atoms within molecules bonded strongly with covalent bonds Simple molecular substances have low mpt/bpt: only weak intermolecular forces break, not strong covalent bonds Nonpolar simple molecules tend to be soluble in nonpolar solvents (hexane): intermolecular forces form between molecules and solvent, weakening intermolecular forces in simple lattice, so they break & compound dissolves Tend to be insoluble in polar solvents (water): little interaction between molecules in lattice & solvent molecules, intermolecular bonding in solvent too strong to be broken Solubility of polar simple molecular substances is hard to predict: depends on strength of dipole. May dissolve because polar solute/solvent molecules can attract each other (similar to ionic dissolving). Some with part polar part nonpolar dissolve Simple molecular structures don’t conduct electricity: no mobile charged particles within structure, so nothing to complete an electrical circuit 6.4 hydrogen bonding Found in molecules containing: electronegative atom w/ lone pair electrons (O,N,F) attached to a hydrogen atom www.pmt.education  Shape around hydrogen atom involved in bond is linear, bond shown by dashed line Solid is less dense than liquid: hydrogen bonds hold water molecules apart in open lattice structure, water molecules in ice further apart than in water, so ice is less dense than liquid water & floats. 2 lone pairs on oxygen and 2 hydrogens, so each molecule can form 4 bonds→ open tetrahedral lattice full of holes. Bond angle H involved 180° Relatively high mpt/bpt: hydrogen bonds & london forces, large quantity of energy is needed to break hydrogen bonds. When ice lattice breaks, arrangement hydrogen bonds is broken & when boils hydrogen bonds break completely Other anomalous properties: relatively high surface tension + viscosity DNA is held together by hydrogen bonds: AT pair form 2 hydrogen bonds, CG pair forms 3. Pairs match up correctly because bases must fit together so hydrogen atom one molecule and O/N from other align correctly www.pmt.education