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Lecture Notes for BIOL111: Molecules of Life

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Lecture 1: Energy, Elements, Atoms & Bonding Lecture 2: Water, Concentration, Equilibrium Constants, pH Lecture 3: Molecule Shape, Functional Groups & Isomerism Lecture 4: Nucleic Acids Lecture 5: Amino Acids Lecture 6: Polypeptides & Protein Folding Lecture 7: Protein Structure & Function I ...

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  • August 9, 2021
  • 23
  • 2020/2021
  • Lecture notes
  • Emma shawcross, andrew lewis
  • All classes
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amal2001
Lecture 1: Energy, Elements, Atoms & Bonding

Energy
- kinetic: movement, heat
- potential: energy possessed due to position
 the higher up you are, the more potential energy you have
 chemical energy is form of potential energy
- metabolism: sum of all building up & breaking down chemical reactions in the body
- anabolism: building up molecules
- catabolism: breaking down molecules

Elements
 substance that can’t be broken down to other substances in chemical reactions
- about 25/92 are essential to life with C, H, N & O making up 96% of living matter
 17/92 essential to plant life
- most of organism’s remaining weight made of P, S, Ca & K with trace elements like Fe & I
 phosphorous found in DNA and bones
 sulphur found in amino acids where it forms disulphide bridges essential for protein structure
 calcium found in bones & teeth
 potassium is key for cell signalling in nervous system
 iron found in haemoglobin
 iodine key for thyroid hormones

Atoms
- smallest unit of matter still retaining an element’s properties
- made up of negative electrons, positive protons and neutral neutrons
 electrons move very fast around nucleus in clouds
 no. of protons = no. of electrons
 nucleus is extremely tiny
- chemistry of atom determined by no. of electrons
- atomic mass is measured in daltons
 mass of protons & neutrons = 1.66x10-24
- atomic number: no. of protons
- atomic mass: sum mass of all subatomic particles




atom with cloud of simplified model
electrons




Isotopes
- chemistry is identical but masses are different due to same number of protons and electrons but
different number of neutrons
- unstable isotopes are radioactive
 half-life: amount of time taken for radioactivity to decrease by half

, - s orbitals are spherical
 1s orbital closest to nucleus
- p orbitals have a dumbbell shape
 there are 3 p orbitals at each
Atomic Orbitals
energy level that hold 6 electrons
- region in atom that can hold up to two electrons with opposite spins
with 2 electrons in each one
- filled orbitals more stable than unfilled
- elements try lose or gain electrons to fill orbitals and become more stable
- valence electrons: electrons directly involved in chemical reactions to form bonds



Compounds
- substance made of two or more elements in fixed ratio
 table salt (NaCl) made of equal numbers of Na and Cl atoms
Covalent Bonds
- sharing of electrons between two atoms
- non-polar covalent bond = equal sharing of electrons e.g. H2
- polar covalent bonds = inequal sharing of electrons e.g. O-H bond in H2O
 electronegativity: measure of the tendency of an atom to attract a bonding pair of electrons
 polar covalent bond stronger than regular covalent bonds
 in polar covalent bond, element that’s more electronegative is given a partial negative charge
(δ−) and less electronegative element given partial positive charge (δ+)

Hydrogen Bonds
- molecules with polar bonds
 can only form in a H-F bond, H-O bond or H-N bond as they are the most electronegative
elements
- weak compared to covalent & ionic bonds however when paired with one of them, they’re very
strong together
 e.g. DNA structure held by H bonds

Ionic Bonds
- when two atoms with different attractions for valence electrons come together forming a crystal
 very strong ~700kJ/mole
sodium:
1s2, 2s2, 2p6, 3s1  1s2, 2s2, 2p6, 3s1
chlorine:
1s2, 2s2, 2p6, 3s2, 3p5  1s2, 2s2, 2p6, 3s2, 3p6




= Na+
= Cl -

Van Der Waal’s Forces
- interaction between electron clouds forming δ− and δ+ which are attracted to each other
- very weak ~0.5kJ/mole but collectively strong
- short range

, Lecture 2: Water, Concentration, Equilibrium Constants, pH



Formation of Water Molecule

- oxygen more electronegative than hydrogen so hydrogen ends
up with a δ+ and oxygen with a δ−
 due to this, water molecules can interact with each other &
form H bonds
 unequal electron sharing leads to V shape
- two H atoms form single polar covalent bonds with O atom



Electron Orbitals in Water
3
1S 2S 2P 1 sp hybrid
O s p orbitals
H - both hydrogens share their 1s electron with oxygen’s 2p sub shells
 forms covalent bond
- electronic configuration of water is sp3 hybridised
H  sp3 hybridised is the mixing of one s orbital and three p orbitals having nearly
the same energy to form four equal orbitals
 sp3 hybridised orbitals repel each other and are directed to four corners of a
tetrahedral

Water Molecule Shape
- has a V shape
- electrons repel as much as they can forming tetrahedral shape with 109.5°
 however, lone pair is closer to central atom nucleus so repels more than bonding pair reducing
bond angle to 104.5°

Water & Hydrogen Bonding
- higher boiling point:
 water is liquid at room temperature and has a very high boiling point compared to molecules with
similar molecular masses
 H bonds in water require a lot of energy to break
- higher cohesion & adhesion:
 water molecules can H bond with each and with surfaces
 important for plants as water can travel from the roots to the leaves against gravity
- higher heat of vaporisation:
 energy required to convert 1g of water from liquid to gas at 25℃
 high for water (2424J) so lots of energy needed to go from liquid to gas
 when water turns to gas, highest energy molecules leave as gases cooling the rest of the water
molecules e.g. sweat cooling

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