Biochem first year notes PDF

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Biochem first year lecture notes...

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Atoms, Ions and Molecules Atoms consist of a tiny nucleus of protons and neutrons; this is surrounded by a large volume of space apart from the electrons. Electrons have a negative charge that balances the positive charge of the protons. The majority of the mass of an atom is in the nucleus. (99.9%) Isotopes of an element have a different number of neutrons in their nuclei. Atomic number – number of protons Mass number – number of protons and neutrons Number of protons = number of electrons All the atoms of a particular element must have the same number of protons in the nucleus, but the number of neutrons may differ. Isotopes have the same atomic number but different mass numbers 

Mixture – more than one type of substance, e.g air Elements – cannot be split into simpler compounds, e.g gold Compounds – a specific combination of elements that can be broken down, e.g NaCl Organic compound – contains carbon and usually hydrogen, e.g methane, glucose Elements can be atomic (Na, Fe, Ar) or molecular (O₂, S₈) Compounds – molecular (CO₂, H₂O) or ionic (NaCl, CaO) The Law of Conservation of Mass – in a chemical reaction, matter cannot be created or destroyed Total mass of reactants = total mass of products The Law of Constant Composition – a compound always contains the same elements united together in the same fixed proportions by mass Relative atomic mass – comparing the mass of atoms to the mass of carbon 12 atoms. Symbol = (Ar) Relative molecular mass (Mr) – sum of relative atomic masses of all the atoms in one molecule of a covalent compound Moles (units = Mol) The amount of an element that contains 6.022 x 10²³ atoms (Avagadro Constant) = one mole of that element 12g of Carbon¯¹² contains 6.022 x 10²³ atoms = one mole Mass of one mole of a substance = molar mass (units – g/mol¯¹) Molar mass of carbon = 12g/mol¯¹ Amount of substance (mol) = mass of substance (g) ÷ molar mass of substance (g/mol) Stoichiometry – relationship between the amounts of reactants and products in a chemical reaction

Percentage yield = (actual mass of product ÷ calculated mass of product) X 100 100% yield – quantitative yield Limiting reactant – added in amounts that are not stoichiometric. Molar concentration (Molarity – M) – amount of solute (in moles) per litre of solution M = mol dm¯³ = moles of solute (mol) ÷ litres of solution (dm³) 1L = 1dm³ 1ml = 0.001dm³ Concentration = moles ÷ volume Dilutions To determine the molarity of a diluted solution: M₁ X V₁ = M₂ X V₂

(1 = concentrate, 2 = dilution)

Finding Molarity: 1. Calculate relative molecular mass (Mr) 2. Actual mass divided by Mr = number of moles 3. Number of moles divided by volume 4. Answer = concentration with units (M) – molarity Finding Mass: 1. Convert volume to dm³ 2. Concentration x volume, to get the number of moles 3. Number of moles x Mr (relative molecular mass) = mass Mol = grams ÷ grams/mol Atomic Structure Metal atoms lose electrons to form cations Non-metal atoms gain electrons to form anions Supplying energy to any gaseous atoms can make them lose electrons and become ionized The energy levels of electrons in atoms are arranged into shells which are numbered 1,2,3… this is called the electronic configuration. Atomic orbitals – mathematical descriptions or probability distributions of electrons in the space surrounding the nucleus The outer most shell has ‘valence’ electrons; this means they are not a complete ring. They have the lowest ionisation energies and are involved in bonding. Inner shell electrons have higher ionisation energies Subshells – subsets of the principal energy levels of electrons in atoms These are labelled s, p, d, f (stemming from sharp, principle, diffuse and fundamental, in reference to the lines they describe)

Electrons in the same subshell do not shield each other very much from the attraction of the nucleus. Subshell energy levels can be arranged into subsets called orbitals. S-subshell = 1 orbital P-subshell = 3 orbitals D-subshell = 5 orbitals F-subshell = 7 orbitals  No more than two electrons can be in the same orbital Quantum numbers S – l = 0, ml = 0, ms = ± ½ (spherical shape) P – l = 1, ml = -1, 0, +1, ms = ± ½ (two lobes) D – l = 2, ml = -2, -1, 0, +1, +2, ms = ± ½ (four lobes/ two lobes with a middle ring) Example electronic configurations: Oxygen (element number 8 = 8 electrons) 1s², 2s², 2p⁴, the electron valence is 6e¯ (2+4=6) Phosphorus (element number 15 = 15 electrons) 1s², 2s², 2p⁶, 3s², 3p³, the electron valence is 5e¯ Full electron shells: 1s², 2s², 2p⁶, 3s², 3p⁶, 3d¹⁰, 4s², 4p⁶, 4d¹⁰, 4f¹⁴ Example values of the four quantum numbers for an electron in a: 2p orbital: n=2, l = 1, ml = -1, 0, +1, ms = ± ½ 1s orbital: n=1, l = 0, ml = 0, ms = ± ½ 3d orbital: n=3, l = 2, ml = -2, -1, 0, +1, +2, ms = ± ½ (ms – the spin, two electrons in the same orbital have opposite spins) Noble gas molecules are monatomic – their electron shells are full, therefore the atoms do not combine with one another Octet rule – elements try to gain noble gas configuration Ionic bonding – electrons are transferred from one atom to another, the ions also pack together to form a crystal lattice Coordination number – (of an ion in a crystal ionic compound) how many equidistant oppositely charged ions The type of crystal structure adopted by an ionic compound is governed by the numbers, shapes and sizes of cations and anions. Loss of electrons = cations Gain of electrons = anions Anions may be formed by molecules acting as acids and losing protons Cations may be formed by molecules acting as bases and gaining protons

Van der Waals forces – noble gases are monotomic (full electron shells) therefore the atoms do not combine with each other, at very high pressure the noble gases condense to liquids in which the atoms are held together by Van der Waals forces. The level of forces increases with the number of electrons in an atom. Covalent bonds – two nuclei sharing electrons, the attraction of the nuclei to the electrons is greater than repulsions between the nuclei/nuclei and electrons/electrons Hydrogen, the halogens, oxygen and nitrogen molecules are diatomic (two atoms in one molecule) The molecule shares electrons between both atoms Single bond = sharing two electrons (one from each atom) Double bond = sharing four and triple bond = sharing six It is sometimes, in the case of carbon for example, energetically favourable to promote electrons from 2s to 2pz to give 4 half-filled orbitals. Electrons can only be promoted in the same shell The greater the overlap between atomic orbitals – the stronger the covalent bond Long bond length = weak covalent bond Short bond length = strong covalent bond Homonuclear molecules (both have the same pull on the electrons) electrons are equally shared in the sigma bond between atoms Heteronuclear molecules – electrons are shared unequally, eg. Polar molecule – one atom is more electronegative and therefore attracts more of the electron cloud, leading to unequal sharing of the electron bond pair Metals consist of a lattice structure of mutually repelling positive ions, held together by their attraction for a ‘sea’ of mobile electrons. (The ones not involved in the covalent bond are mobile) Electronegativity difference ΔEN ΔEN > 1.7 ionic bond ΔEN < 0.4 covalent bond 0.4 < ΔEN < 1.7 polar covalent bond Bond length – average distance between the nuclei of the two atoms, aka equilibrium distance Bond energy – needed to separate the atoms in the molecules of one mole of compound A dative/coordinate bond is a covalent bond in which one of the atoms has supplied both of the electrons being shared Lewis acid – a molecule/ion that can accept a pair of electrons to form a covalent bond Pauling electronegativity index (Np) – measure of how strongly an atom in a compound attracts electrons in a bond

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The greater the difference in the electronegativities of two atoms, the more polar the covalent bond between them

Intermolecular forces Van der Waals – weak, short-range forces between atoms/molecules due to the attraction between dipoles, these forces arise from the movement of electrons in relation to the nuclei which produces weak instantaneous dipoles that attract each other The electron clouds distort (polarise) due to electrostatic repulsions Dipole-dipole attractions – between the positive end of one permanently polar molecule, and the negative end of one permanently polar molecule, this interaction lowers the potential energy Dipole moment = charge x separation Ion-dipole interactions – many ionic compounds dissolve well in water and other polar solvents: energy released by the formation of ion-dipole interactions, compensates for the lattice energy consumed in overcoming the electrostatic forces between the cations and anions Cations and anions attract polar water molecules Water molecules get between the ions and weaken their attractions Water molecules completely surround the cations and anions Hydrated ions leave the crystal and mix with the water  

Forming covalent bonds is an exothermic reaction (heat is given out) Breaking covalent bonds is endothermic

Metal + non metal reaction = ionic compound (exothermic) Metallic bonding – valence electrons are delocalised over all atoms, a metal is a lattice of positive ions in a sea of electrons. These electrons are free to travel through the metallic lattice Metals are good conductors of electricity and heat – the nuclei vibrate and pass vibrations through the lattice to the other side The bigger the molecule, the higher the boiling point Hydrogen bonds are formed only with small electronegative elements: F, O, N, Cl A hydrogen bond is an extremely strong dipole-dipole interaction  Covalent bonds are shorter than hydrogen bonds Dative bonding – like covalent, but one atom supplies both electrons, This is often done to help other molecules become more stable, and in joining the new molecule is overall more stable, this lowers the overall energy Bases are proton acceptors Acids are proton donators VSEPR – valence shell electron pair repulsion: electron pairs have the same negative charges, they try to get as far away from each other as possible (electrostatic repulsion)

Lone pairs of electrons are less tightly held therefore they occupy more space and squeeze other bonds closer together, lone pairs are closer to the central atom Number of electron pairs around central atom: 2 – linear 3 – planar trigonal 4 – tetrahedral 5 – trigonal bipyramid 6 – octahedral 7 – pentagonal bipyramidal Op+ Hybridisation – atoms can mix their orbitals within the same shell Rate of reaction – change of concentration of a reactant or product with time ρ = density, units g/mL Radioactive Isotopes – depends on stability of nucleus – ratio of neutron/proton to determine overall charge Heavy nuclei have an increasing number of protons Heavy nuclei – undergo α decay and α radiation (shift two elements to the left – lose 2 protons and 2 neutrons) Light and heavy nuclei – β decay and β radiation (shift 1 element to the right – gain a proton and lose an electron) Excited nuclei and excess energy produces gamma γ radiation α rays – typically produce helium nucleus, 2-4cm travel in air, 0.05mm tissue depth, shielded by paper and clothing, typical source Radium 226 β rays – typically produce an electron, travel 200-300cm in air, 4-5mm tissue depth, shielded by heavy clothing/lab coats, carbon-14 γ rays – electromagnetic radiation, 500m travel in air, >50cm tissue depth, shielded by lead/concrete, technetium 99m Units: Bequerel (measure of activity of radiation) 1Bq = 1 disintegrations/s Gray (measure of absorbed radiation) 1Gy = 1J/Kg Sievert (measure of biological damage) 1Sv = 1Gy x constant (1 for β, γ and X-rays, 10 for high energy protons and neutrons, 20 for α) Half-life – time taken to decay half original concentration Synthetic isotopes – conversion of stable isotopes into radioactive, fast-moving particles (transmutation) Energy can be generated by fusion and fission Β radiation Applications: Gastrointestinal tract diagnosis, pancreatic cancer detection, treatment of bone and prostate cancer

γ radiation applications: imaging skeleton, brain, liver etc X-ray applications: Tumour detection, diagnosis of myocardial infarction, treatment of brain cancer (Energy) state of matter –can be completely described by pressure (pascals), volume (cubic metre), temperature (Kelvin) and amount (no of moles) Sublimation – direct change from solid to gas Deposition – direct change from gas to solid Energy – the capacity to do work 1st law of thermodynamics – matter has energy content; heat is equivalent to mechanical energy (kinetic/potential) Energy cannot be created nor destroyed, only transferred from one form to another Heat – the transfer of energy as a result of a temperature difference between the system and surroundings A change in heat/work that increases the energy of the system is defined as a positive change Heat (Q) – positive change means in (absorbed by system) – endothermic negative change means out (absorbed by surroundings) – exothermic Work (W) – positive means done on system negative means done by system Measuring heat changes – calorimeter (a reaction vessel surrounded by a water bath containing a thermometer) Q (heat supplied) = mass of water x heat capacity of water x temperature change Q = M x C x ΔT Constant pressure (eg plunger allowed to move) Q = ΔH Change in enthalpy Constant volume (eg stopper cannot move) Q = ΔU Change in internal energy ΔH = ΔU + PΔV PΔV (Work done by piston) The measured heat change for a reaction at atmospheric pressure = reaction enthalpy

Enthalpy – a state function (a change in enthalpy depends only on the difference between the stating state and final state, not the route taken) ΔH (change in enthalpy) = the same whether the reaction is carried out in one or several steps The standard enthalpy of formation is the enthalpy change when 1 mol of a pure substance is formed from its elements. Each element must be in the physical and chemical form which is most stable at normal atmospheric pressure and a specified temperature (usually 25°C).

Hess's Law – the standard enthalpy of a reaction = the sum of the standard enthalpies of the reactions into which the overall reaction may be divided, even if this division is artificial Standard enthalpy = at 1 atmosphere CO(g) + ½ O2(g) → CO2(g) 1. 2.

C(s) + O2(g) → CO2(g) C(s) + O2(g) → (3.) CO(g) + ½ O2(g) → CO2(g)

ΔH 1 = ΔH 2 + ΔH 3 Energy change is the same whichever route is taken When elements are in elemental form – enthalpy change = 0 Negative value for kj mol-1 = exothermic process The Second Law: During spontaneous change, energy and matter become more disordered Here energy is the release of enthalpy Matter is the increase in entropy Gibbs Free Energy ΔG = ΔH – TΔS When ΔG < 0, the process is spontaneous ΔH = enthalpy change ΔS = entropy change ΔG = Gibbs free energy change (allows you to decide whether a process is spontaneous T = temperature at which reaction is taking place ΔS will be positive Ideal Gas Law: PV = nRT P = pressure

V = volume n = number of moles R = gas constant (8.314 JK-1 mol-1) T = temperature Assumptions ideal gas law is based on: – gases consist of small particles which have negligible volume – gas molecules are in constant rapid motion, undergoing elastic collisions with each other and their container – molecules do not interact with each other except during collisions Carbon – electron configuration 1s2 2s2 2p2 (4 valences) covalent and polar covalent bonds sp3 hybridisation – promoting a 2s electron into an empty 2p orbital gives 4 unpaired electrons (2s1 2p3) and can form methane. The s and p orbitals are mixed to form a hybrid sp2 hybridisation – 1 electron is promoted to the p shell, the s and two of the p electrons hybridise together leaving one in the p shell – a double bond is formed, eg ethane sp hybridisation – 1 s orbital and 1 p orbital are mixed, results in carbon triple bonds, eg ethyne Benzene – planar molecule, each carbon is sp2 hybridised, 6 electrons in pure p orbitals become delocalised over the ring Most organic compounds are not soluble with water Organic compounds with oxygen or nitrogen are somewhat soluble – the electronegative atoms form hydrogen bonds with water Major classes or organic compounds: Aromatic – contain cyclic conjugated double bond system Aliphatic – all other compounds without a ring system Electrophile – partially positive (is attracted to electrons) Nucleophile – partially negative (is attracted to protons in the nucleus) Oxidation – when carbon atoms are oxidised they often form additional bonds to oxygen Reduction – when carbon atoms are reduced they often form additional bonds to hydrogen Substitution – exchange of one atom in a molecule with another atom/group of atoms Elimination – a single reactant is split into two products Addition – two reactants add together to form a single product

Homolytic bond cleavage of a single bond means a 1 electron transfer to each atom and the formation of radicals Heterolytic bond cleavage of a single bond means a two electron transfer and usually the formation of ions Protein structure DNA encodes the protein as a linear sequence of nucleotides  RNA  Protein (linear sequence of amino acids folds to an elaborate functional structure)

Amino acid structure: central carbon, a hydrogen, an amine group, a carboxyl group, a variable R-group Zwitterion – means the real life chemical form Angles around the central carbon influences folding in tertiary structure Alpha carbon is the central carbon – can be in two different forms (mirror images of each other) D- and LAmino acids in proteins all have L-configuration D-amino acids exist but not in proteins encoded by DNA, eg they exist in bacterial cell walls 20 possible amino acids Glycine – simplest, R group = H Alanine, R = CH3 Serine, R = CH2OH (has a polarity) Cysteine, R = CH2SH Dipeptide – two amino acids joined by a peptide bond Polypeptide – multiple amino acid residues joined by peptide bonds The nitrogen in the bond has a lone pair of electrons that can partially delocalise and the N adopts a double bond characteristic, as such the oxygen and nitrogen can’t rotate Amino acid residues – carbonyl group instead of carboxyl group, NH instead of NH3+, they are what’s left when the amino acids link Peptide backbone is good at hydrogen bonding Alpha helix formed of hydrogen bonds, made of n  n + 4 links Eg #1 links to #5, #2 links to #6 This forms a very stable cage structure Amino acids display side chains on the outside of the helix and are able to interact Helix formers – glutamate, glutamine, alanine, histidine Helix destabiliser (likely to break up helices if too frequent) – tyrosine Helix breaker – proline (has a side chain with a phi bond that cannot rotate) Beta sheets: made of strands of protein Anti-parallel beta-sheet have strands that run parallel but point in different directions CN NC This allows hydrogen bonds to form to knit the sheet together between the amine and carbonyl Some carbonyl groups point away to allow others to attach Eg silk fibroin, has long stretches of repeating Gly and Ala residues, it will not stretch longitudinally but very flexible laterally

Parallel beta-sheet Hydrogen bonds form diagonally between amine group and carbonyls NC NC Hydrogen bonds are strongest when molecules bond in straight lines – therefore parallel sheets are slightly weaker Structural motifs are arrangements that occur regularly, including: All alpha – alpha helix, U turn, alpha helix, eg myoglobin All beta – beta sheet, U turn, beta sheet, eg IgG Alpha/Beta – the two alternate, eg triose phosphate isomerase The repeat of alpha and beta forms a beta-barrel Tertiary structure depends on packing of secondary structures and side chain interactions Hydrophobic side chains – valine, leucine, phenylalanine (largely hydrocarbon) Hydrophilic – aspartate (COO-) lysine (NH3+), serine (OH) Hydrophobic effect – most important determinant of tertiary structure, Hydrophobic collapse – phase 1 of folding for tertiary structure, an entropic effect (the energetic driver fo...