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112 Cards in this Set
- Front
- Back
- 3rd side (hint)
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List the types of electromagnetic radiation in order of decreasing frequency |
Gamma rays, X-rays, UV, visible light IR, microwaves, radiowaves |
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State the differences between wavelength and wave number |
Wave number is the reciprocal of wavelength (distance between 2 wave crests). i.e. wave number = 1/wavelength (waves per metre - units: m^-1) |
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Describe atomic emission spectrum |
☆sufficient energy supplied to atom ☆electron promoted to higher level ☆electron falls back down ☆photon emitted (energy = difference in energy of levels) ●position of lines = what is present; intensity = how much is present ●no colours seen: insufficient energy supplied or light emitted not visible ●energy levels further from nucleus = closer together to form continuum |
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State 3 transition series |
• Lyman: transitions fall from higher energy levels to n = 1 (UV emitted) • Balmer: n = 2 (visible light) • Paschen: n = 3 (IR) |
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What are emission & absorption spectra used for? |
• to quantify/identify elements ▪ emission: measures radiation emitted when electrons fall ▪ absorption: measures radiation required to promote electrons |
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In absorption spec, how does ion concentration effect radiation absorbed? |
As concentration of ion in solution increases, there is an increase in intensity of radiation absorbed. |
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Define principal quantum number |
What shell is the electron in? • higher value of "n" - further electrons are from nucleus • increased "n": atomic orbitals increase in size & energy |
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Discuss angular momentum quantum number (l) |
What subshell is electron in? (And therefore what shape is orbital?) •4 subshells: s (spherical), p (dumbell), d (dumbell) & f • s = 2e; 3 × p = 6e; 5 × d = 10e •1st shell: s, 2nd: s + p, 3rd: s, p + d •quantified as values of "l" - s: l = 0, p: l = 1, d: l = 2 (l = n - 1) |
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Define magnetic quantum number (m) |
What orientation in space are they? •spherical are non-directional •values of "m" range from -l ➡ +l ○e.g. d subshell: l = 2, so m = [-2, -1, 0, 1, 2] (5 orientations) |
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Define spin quantum number (s) |
Is the electron spinning clockwise or anti-clockwise? •spins are equal and opposite, so that they can co-exist in 1 orbital • s = +1/2 (clockwise) or -1/2 (anti) |
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Define degenerate orbitals & state how many electrons an orbital holds |
•degenerate orbitals = same energy •orbital holds a max. of 2 electrons |
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Define atomic orbital |
Region in space where probability of finding an electron is high (over 90%) |
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State how to calculate number of orbitals & electrons in energy level |
How many orbitals? = n2 • e.g. n = 3 ➡ 9 orbitals (1 x s, 3 x p, 5 x d) How many electrons? = 2n2 • e.g. n = 3 ➡ 18 electrons |
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State the Pauli Exclusion Principle |
an orbital can't have more than 2 electrons AND they must have opposite spins |
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State Hund's Rule |
In degenerate orbitals, electrons will fill each orbital singly with parallel spins, before pairing (electrons won't favour a particular orbital) |
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State the Aufbau Principle |
When electrons are placed into orbitals, electrons are filled in order of increasing energy |
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State Heinsberg's Uncertainty Principle |
The more precisely you know one quantity, the less precisely you can know another associated quantity |
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Define ionisation energy |
The energy required to removed 1 mole of electrons from 1 mole of gaseous atoms/ions |
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State and explain trends in ionisation energy across the periodic table |
• increases across period ▪ more protons = increased nuclear charge ▪ exceptions due to stability associated with full/half-full or empty subshells • decreases down group ▪ electrons in larger atoms are further from nucleus ▪ inner electrons "shield" nucleus |
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What is the oxidation number of an atom? |
The number of electrons needed to be added/removed from an element for it to combine to form a compound. e.g. ox. number of Cl in HClO2 is +3 - 3 electrons need to be removed from Cl |
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Define oxidation & reduction |
•oxidation is loss of electrons (increase in O:H ratio - more O for every H) •reduction is gain of electrons (decrease in O:H ratio - less O for every H) |
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Define transition metals |
Metallic elements with an incomplete d subshell in at least 1 of their ions |
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State which is filled first in transition metal subshells - 4s or 3d? |
•when filling, 4s is filled before 3d - 4s has lower energy than 3d even though it is further from the nucleus •when emptying, 4s emptied before 3d - repulsion of electrons in 3d push 4s to higher energy level |
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State two transition metals which don't follow the Aufbau Principle |
Chromium - 4s13d5 (not 4s23d4) Copper - 4s13d10 (not 4s23d9) ○special stability of full/half full 3d |
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State 2 transition metal elements that don't comply with the definition |
Scandium - only ion is 3+ (empty d) Zinc - only ion is 2+ (full d) |
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State 4 properties of transition metals |
○ form coloured ions ○ form complexes ○ have variable oxidation states ○ show catalytic activity |
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Define a complex |
Central metal ion surrounded by ligands |
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Define ligand |
Negative ions or neutral molecules with 1 or more non-bonding pairs |
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Define and provide examples of monodentate, bidentate & hexadentate ligands |
○ monodentate: 1 pair of electrons donated to central ion (Cl- & H2O) ○ bidentate: 2 pairs (oxalate ion) - look for 2 functional groups ○ hexadentate: 6 pairs (EDTA) • bind to central ion in 1:1 ratio • chelating agent ("claws" at ions) |
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Define coordination number |
Number of direct links to the central metal ion - specied outside [] aren't directly links, simply maintain charge • e.g. 6:6 coordination - each X ion surrounded by 6 Y (& vice versa) |
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Compare covalent and dative covalent bonds |
Covalent: shared pair of electrons Dative (coordinate bond): both electrons come from the same atom. (e.g. NH4+) *once formed, indistinguishable |
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Describe the process of naming complexes |
○ ligands named first (alphabetically), then metal & its oxidation state ○ ligand preceded by prefix showing number of ligands (di, tri, etc...) ○ metal name followed by oxidation number (I, II, III, etc...) ○ negative ligand ending in -ide ➡ -o (e.g. chloride ➡ chloro) ○ negative complex, transition metal ends in -ate (e.g. copper ➡ cuprate) ○ salt complex, +ve then -ve (Na+Cl-) |
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State 3 neutral ligands & their name in a complex |
○ ammonia (NH3) ➡ ammine ○ carbon monoxide (CO) ➡ carbonyl ○ water (H2O) ➡ aqua |
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Describe how we see colour |
○ white light shines, some light absorbed ○ transmitted light = white - absorbed ○ transmitted is complementary colour to absorbed light. ● e.g. red absorbed ➡ blue + green (cyan) transmitted ● blue absorbed ➡ red + green (yellow) transmitted ● green absorbed ➡ blue + red (magenta) transmitted |
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Describe colour in transition metals |
• 5 × 3d orbitals degenerate • ligand arrives & bonds to metal • orbitals split due to repulsion • 2 × 3d orbitals raised energy levels; 3 × 3d orbitals lowered energy levels • electrons in lower 3d can absorb energy & be promoted to higher 3d • if energy absorbed has wavelength equal to visible light, complementary colour is transmitted |
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Explain the effect of different ligands on d-d transitions |
• stronger ligands cause a bigger split in 3d subshells • bigger split: more energy difference between levels - shorter wavelength absorbed |
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List ligands in order of decreasing strength |
CN- > NH3 > H2O > OH- > F- > Cl- > Br- > I- |
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State the name of the theory of d-d transitions and describe the conditions of the theory |
Crystal Field Theory - only for octahedral complexes with incomplete 3d subshells with at least 1 × 3d electrons |
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Describe why transition metals are used as catalysts |
• transition metals on active site form weak bonds with substrates • unpaired d electrons allow intermediate complexes to form • substrate covalent bonds weaken • substrate more susceptible to attack by reactant molecules • alternate pathway with lower Ea |
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State transition metal catalysts used in 3 processes |
• [iron] in Haber process (ammonia) • [platinum] in Ostwald (nitric acid) • [platinum], [rhodium] & [palladium] in catalytic converters |
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Describe the circumstances affecting position of dynamic equilibrium |
● Rate of forward reaction = rate of backward. • Conc. of products/reactants constant, not equal. • Forward & backward continue, but at same speed. ● Position if equilibrium changed by: • altering conc. of reactant/product • changing pressure if moles differ • altering temp. (+ heat: endo; - heat: exothermic) |
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Describe implications of magnitude of equilibrium constant (K) |
Magnitude of K: equilibrium position • K > 1 ➡ products favoured • K = 1 ➡ equal split (50/50) • K < 1 ➡ reactants favoured |
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What states of reactants/products are in equilibrium expression? |
• solids not included (activity = 1) • in heterogeneous equilibria, liquids aren't included (considered solvents) |
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What affects the equilibrium constant? |
Only temperature: ▪ endothermic reaction: + temp. ➡ products favoured ➡ [products]:[reactants] increased ➡ K increased ▪ exothermic: + temp. ➡ reactants favoured ➡ [products]:[reactants] decreased ➡ K decreased Conc./pressure/catalysts don't: aA + bB ↔ cC + dD • increase A ➡ used to increase C/D, but not all is used up ➡ A higher; B lower; C & D proportionally higher ➡ same ratio ➡ "K" constant |
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Describe the partition coefficient |
• liquids that don't mix: immiscible • polar solutes dissolve in polar solvents (water); non-polar (hydrocarbons) in non-polar • add substance more/less soluble in 2 immiscible liquids & shake • when settled, conc. in each layer is constant, but there is interchange between liquid layers (equilibrium) ▪ "K" is called patition coefficient |
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State uses for partition coefficients |
• extract & purify product from reaction mixtute (solvent extraction) • uses separating funnel • e.g. coffee/water/dichloromethane ▪ caffeine more soluble in dichloro than water, can be removed via evaporation • method improved by extracting multiple times or using several small volumes of extracting solvent |
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Describe the water equilibrium (ionic product of water) |
H2O(l) + H2O(l) ↔ H3O+(aq) + OH- (aq) • (H3O+): hydronium ion (hydrated protons). (H-): hydride ion • H+ & OH- ➡ water conducts • equal H+ & OH- ➡ pH 7 • endothermic (-ve) ➡ equilibrium lies to the left • ionic product (Kw): 1×10-14 • amphoteric: acts as acid & base |
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Describe the pH scale |
• logarithmic scale: pH change of 1 unit ➡ [H+] change of 10 ▪ dilute strong acid by 10; pH change of 1 unit ▪ dilute weak acid by 10; pH change of 0.5 unit • in chemistry, "p" = -log • pH × pOH = 14 • [OH-] × [H+] = 1 ×10-14 |
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Compare strong & weak acids/bases |
• strong acid is completely ionised in aqueous solution (HCl, HNO3, H2SO4). Stronger acids have more than 1 ionisable proton (polyprotic) • weak acid is partially ionised, forming an equilibrium (H2SO3) • strong base: NaOH, KOH • weak base: ammonia/amminrs |
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Define Brønsted-Lowry acids/bases |
• Brønsted-Lowry acid donates H+ • Brønsted-Lowry base accepts H+ |
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Define Lewis acid/bases |
• Lewis base: can donate unpaired electrons (forming covalent bond) • Lewis acid: can accept unpaired electrons (forming covalent bond) |
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Describe conjugate acids & bases |
• whatever is left of acid after proton is donated - conjugate base • whatever is left of base after proton is accepted - conjugate acid • stronger acid/base ➡ weaker conjugate base/acid |
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Describe the acid dissociation constant (Ka) |
• for every H+ ion produced, an A- ion must be produced also • in solution of weak acid, only a small proportion of original acid molecules dissociate into ions • smaller Ka ➡ stronger base |
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Explain buffer solutions |
• Buffer: pH remains approx. constant when small amounts of acid/base/water are added • acid buffer: weak acid + salt • basic buffer: weak base + salt
e.g. ethanoic acid/sodium ethanoate ▪ + acid: H+ reacts with CH3COOH- ➡ more CH3COOH ➡ [H+] same ▪ + base: H+ reacts with OH- ➡ more CH3COOH dissociates ➡ [H+] same |
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Discuss use of indicators |
• used in acid/base titrations - change colour at end point • usually weak acids (conjugate base) • pH range of indicator must coincide with point at which pH of titre rapidly changes (half-way down vertical) |
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Define bond enthalpy & mean bond enthalpy |
• bond enthalpy: specific bind within specific molecule (O=O only in O2) • mean bond enthalpy: average for particular bond measured over many types of molecule (C-H in lots) |
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Define standard enthalpy of formation |
Enthalpy change when 1 mole of substance is formed from its elements in their standard states |
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State the 2nd law of thermodynamics |
In any spontaneous process, overall degree of disorder (entropy) must increase |
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Define electron affinity |
E(g) + e- ➡ E- (g) Energy required to add 1 mole of electrons to 1 mole of gaseous atoms (reverse ionisation energy) |
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What is the total entropy change of a system at equilibrium? |
Total entropy change = 0 Entropy of any substance at 0K = 0 |
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What is the effect of exo-/endothermic reactions on entropy? |
Exothermic ➡ increased entropy Endothermic ➡ decreased entropy |
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Discuss circumstances indicating feasibility of reactions |
• Change in entropy: positive for feasible reaction • Gibbs free energy: negative for feasible reaction (just feasible at 0) ▪ at equilibrium, change in G = 0 |
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Define empirical formula & state how it is determined. |
Shows simplest whole number ratio of atoms in a compound. Determined by elemental microanalysis |
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How are molecular orbitals formed? |
Overlap between 2 atomic orbitals where a maximum of 2 electrons can be found |
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Describe the energy of molecular orbitals |
• one of the MO (bonding) has lower energy than the atomic orbitals ▪ lower energy = covalent bond of high electron density. Electron is attracted to both nuclei (stable) • other MO (anti-bonding) has higher ▪ low electron density, cancelling each other out (less stable) |
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Describe difference between sigma and pi bonds |
• sigma: electrons along internuclear axis (end-on overlaps) • pi: p-orbitals can also overlap side-on (in double/triple bonds) • sigma stronger than pi - second bond in C=C has less energy than 1st |
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Describe hybridisation |
• electron from 2s promoted to empty 2p orbital • 2s orbital mixes with 3 × 2p orbitals, forming 4 degenerate sp3 hybrid orbitals - used in alkanes • alkenes use sp2 hybrids: 2s orbital mixes with two 2p orbitals • alkynes use sp hybrids: 2s orbital mixes with one 2p |
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Discuss the bonding continuum |
Non-polar covalent • bonding electrons shared equally Polar covalent • bonding electrons shared unequal •greater difference in electronegativity ➡ more polar ➡ greater ionic character Ionic bonding • bonding electrons transferred from one atom to another (X+:Y-) |
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Describe colour in organic compounds |
• electrons fill bonding before anti • HOMO: electrons with highest E • LUMO: lowest energy, unoccupied • lower gap between HOMO and LUMO allows colour to be emitted • conjugated molecules lowers gap: light absorbed in visible region |
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Define homolytic and heterolytic bond fission |
Homolytic: products are the same - free radicals formed (unpaired electrons - very reactive) Heterolytic: products are different - positive and negative ions formed |
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Define carbocation ion |
• intermediate in organic processes • very unstable: ▪ doesn't have share of 8 electrons ▪ only makes three bonds ▪ electrophile (+ve charge) • tertiary carbocat most stable |
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State Markovnikov's Rule |
Hydrogen atom of small molecule will attach to the C of the double bond that is already bonded to the most H atoms |
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How are halogens (X-X) and hydrogen halides (H-X) added to molecules? |
• Halogens (X-X): cyclic intermediate formed - repulsion of double bond polarises X-X, breaking heterolytically • Hydrogen halides (H-X): carbocat ion formed - pi electrons form bond between C & electrophile |
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Define elimination |
Atoms are removed from one organic molecule, and a double bond forms between two C atoms. • e.g. dehydration (alcohol ➡ alkene) catalysed by conc. sulfuric acid • e.g. removal of atoms on C atoms of a haloalkane - catalyst: strong alkali dissolved in alcohol (e.g. NaOH in ethanol) |
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Describe free radical substitution |
• stability of alkanes means a lot of energy required for substitution • reactions produce mixture of products, so aren't useful for specific haloalkane production • takes place in 3 steps: ▪ initiation - energy breaks bond homolytically ▪ propagation - radicals formed to maintain reaction ▪ termination - radicals removed & full molecules produced |
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Describe electrophilic substitution |
• Nucleophiles (e.g. benzene and its compounds) react like this ▪ electrophiles attack pi electron cloud & groups subtituted ▪ benzene's pi ring makes it very stable, so doesn't do addition ○ Benzene + NO2+ ○ Benzene + CH3Cl ○ Benzene + FeCl3 |
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Discuss nucleophilic substitution |
• nucleophiles (halide ions, cyanide, hydroxide, water & ammonia) attack electrophiles • SN1: two step mechanism, carbocat intermediate, one species in rds, tertiary haloalkanes react like this ▪ C-X bond polarised ▪ nucleophile attacks carbocat •SN2: 1 step mechanism, transition state intermediate, 2 species in rds, primary/secondary react like thid ▪ C-X partially broken; C-Nu is partially formed ▪ transition state as carbocat would be unstable |
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Define condensation |
Two molecules combine to form larger molecule + small biproduct (usually water) |
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Describe the processes of esterification |
• carboxylic acid + alcohol ➡ ester + water • acid chloride + alcohol ➡ ester + HCl |
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How are amides formed? |
Condensation reaction of amine + carboxylic acid ➡ amide + water |
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Define hydrolysis |
Splitting up of a molecule using water (reverse condensation) |
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Discuss oxidation of alcohols |
• primary alcohol ➡ aldehyde ➡ carboxylic acid • secondary alcohol ➡ ketone • tertiary alcohol doesn't oxidise |
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State agents for oxidation and reduction |
Oxidation: hot copper (II) oxide/acidified potassium dichromate Reduction: lithium aluminium hydride |
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Define stereoisomers |
Order of bonding in atoms is the same, but arrangement of atoms in space is different - they are non superimposable. |
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Describe geometric isomerism |
• arises when there is restricted rotation (C=C or saturated C ring) • molecule must have 2 different groups attached to each C of C=C • "cis" isomers: in same side of bond • "trans": in opposite sides of bond |
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State different properties of cis and trans isomers |
• cis: higher bp than trans (more polar). Lower mp and densities, as it is "U" shape while trans is linear |
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Describe optical isomerism |
• molecule contains chiral carbon (asymmetric; 4 different groups) • optical isomers are mirror images • two optical isomers: enantiomers • when drawing, simply swap two groups and the 3D lines attached |
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Describe the different properties of optical isomers |
Identical in every property except effect on plane-polarised light • polarisers allow light only vibrating in a single plan to pass through them • Plane polarised light passed through solution of optical isomer A rotates by angle X • solution of B rotates by angle -X |
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Define a racemic mixture |
• equimolar mixture (same concentration), which is optically inactive |
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Define agonist and antagonist |
• agonist: mimics natural response • antagonist: blocks natural response |
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What is the active part of a drug called? |
Pharmocore |
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What is an analgesic? |
Drug used to relieve pain |
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What is the difference between medicine and a drug? |
• drug: substance altering body's biochemical processes • medicine: drug with beneficial effect |
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Describe mass spectroscopy |
• quantitative analysis provides information about molecular masses • molecules changed to +ve ions, which are deflected according to their mass/charge ratio • Ions break to fragments, most likely on either side of functional groups • largest mass peak: molecular ion |
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Discuss IR spectroscopy |
• provides information about functional groups • when organic molecules absorb IR, energy causes bonds to vibrate ▪ wavelength of absorbed IR depends on type of atoms making up the bonds ¤ stiff bond: shorter wavelength ¤ loose bond: long wavelength • spectrometer measures absorbed wavelength, corresponding to peak ✅ IR obtained for sample in any state |
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Discuss proton NMR spectroscopy |
• information about number of H atoms and their environments • spinning H nucleus acts as magnet, which will line up if placed in magnetic field • if energy of correct frequency is applied, H nucleus can be "flipped" • energy removed, H flips back and emitted energy is measured • TMS (tetramethylsilane): standard against which absorptions are measured. |
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State the (n + 1) rule |
n = number of H atoms attached to adjacent carbon n + 1 = number of peacks in cluster |
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State 4 characteristics of a primary standard |
• High purity - ensures mass weighed is composed entirely of substance • Stable in air & solution - so it isn't used up reacting with chemicals • Readily soluble - solutions of high concentrations are prepared • Large formula mass - minimises uncertainty of mass weighed |
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List suitable primary standards |
• acid: hydrated oxalic acid • base: anhydrous sodium carbonate • oxidising: potassium dichromate • reducing: sodium oxalate • complexing: hydrated salt of EDTA ❌ sodium hydroxide - unstable & impure (absorbs water) |
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Compare end point and equivalence point |
End point: excess of solution is added, when chemist sees colour Equivalence point: exact point when reaction is just complete |
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Why are back titrations used? |
• if there isn't a suitable indicator • if end point isn't clear in direct titre • if reaction is too slow • if analyte is insoluble |
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Provide examples of dessicating agent and dehydrating agent |
Dessicant: conc. sulfuric acid (less useful, carbonises) silica gel Dehydrating: phosphoric acid |
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Describe the process of colorimetry |
• uses the relationship between concentration & intensity of colour • solution of unknown conc. placed in colorimeter, absorbance noted • absorbance compared to calibration graph created from known concentrations • colorimeter = light source ➡ coloured filter ➡ cuvette ➡ light sensor ➡ display ▪ filter: complementary colour |
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State 7 processes for synthesis/analysis of compounds |
• refluxing • recrystallisation • vacuum filtration • distillation • solvent extraction • chromatography • mixed melting point analysis |
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Describe the process of refluxing |
• hydrolysis mixture + anti-bumping granules in round bottomed flask • mixture heated (granules prevent it from boiling violently) • condenser connects flask mouth to cold water tap • vapours produced when flask is heated are condensed into liquids |
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Describe recrystallisation |
Used to purify an impure solid (e.g. benzoic acid) • dissolve solid in small volume • heat mixture • add more solvent if necessary • hot mixture cools to form crystals (less soluble at lower temperatures) • insoluble impurites filtered away |
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Describe vacuum filtration |
• carry out faster filtration under reduced pressure • Buchner glass funnel used • pure recrystallised solid collects on filter paper • washed with cold solvent to remove soluble impurities |
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Describe distillation |
Heating 1+ liquids until boiling, then collecting & cooling vapours ○ used to purify by separation ○ used to identify based off bp • reaction mixture slowly heated • temp. of resultant vapour noted • vapour condenses in conical flask • distillate purified by solvent extraction |
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Describe melting point analysis |
Used to confirm identity of substance • mix a little unknown with pure compound • compare the melting points • impure substances lower mp and broaden temp. range of melting |
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Describe thin layer chromatography |
Common technique used to separate mixtures of substances. • stationary phase: water held on finely ground silica/aluminium • mobile phase: suitable solvent(s) ▪ solvent flows through stationary phase, carrying components - travel at different rates. ▪ solvents dissolve in spot on base ▪ compound carried up plate ▪ distance travelled noted ○ 1 spot = pure; 2 = impure |
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State carriers used in gas chromatography |
Helium or methane used in mobile phase. Light, so easily carry the vapourised liquid |
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Compare accuracy and precision |
Accurate measurements are in close agreement with true value; precise measurements are in close agreement with each other |
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