Biochemistry

1. Biochemistry Chemical Fundamentals

2. Biology is the study of living things • All living matter is ultimately composed of chemical substances • Matter is anything that has mass and takes up space

3. The Bohr-Rutherford Model of the Atom • The nucleus is made up of protons (p+) and neutrons (no) and is surrounded by rings of orbiting electrons (e-)

4. Standard atomic notation • Also called isotope notation X = element symbol A = atomic mass = # protons + # neutrons Z = atomic number = # protons

5. Isotopes • What is the difference between these two atoms?

6. Isotopes • Atoms in which the number of neutrons may differ • 12C and 14C are two isotopes of carbon • In nature, these isotopes differ in abundance • The relative abundance of isotopes is taken into account to produce the atomic mass you see on periodic tables • m carbon= 12.011 amu

7. Isotopes

8. Radioisotopes • The nucleus on some isotopes spontaneously break apart or decay. • The matter and energy given off in this decay process causes these isotopes to be radioactive. • This results in the formation of new elements • When 14C decays, it becomes 14N. • The length of time it takes for a radioactive substance to decay by half is called the half-life. • Radioisotopes can be both useful and dangerous • Radiation can cause mutations in DNA, so need to be handled with care in order to limit exposure

10. Radioisotopes • radioisotopes are used in medical imaging • Injected isotopes localize in specific tissues and release radiation outwards • this radiation is detected by special cameras

11. Radioisotopes • Radioisotopes are also useful in tracing molecules in biochemical pathways (a complex series of reactions in a cell). • Molecules which contain a lot of nitrogen (amino acids for instance), can be ‘tagged’ with a radioisotope of nitrogen

12. Radioisotopes • Are also useful for finding the absolute age of rocks, fossils, or ancient specimens unearthed by archaeologists or palaeontologists • Radiometric dating relies on the half-life of radioisotopes • While an organism is alive, it is taking in carbon and incorporating it into its tissues – all isotopes in their relative amounts. • When it dies, it stops taking in carbon • By measuring the amount of parent isotope vs. daughter isotope, the half-life can be used to calculate how long it has been since the organism stopped taking in the parent isotope

15. Ions • Ions are elements that have gained or lost electrons • Ions are commonly found dissolved in water, such as in the cytoplasm or plasma of the blood • Elements in the same family tend to form the same type of ion (e.g.: Na+, Li+, K+, Rb+) • Some important ions are Ca2+ (used for muscle contraction), Na+ and K+ (nerve and muscle function), Fe2+ and Fe3+(in hemoglobin) and H+ (required for synthesis of ATP)

16. Chemical Bonding • Electrons orbit the nucleus of an atom at a great distance compared to the size of the particles • Analogy: If an apple represented the size of an atom’s nucleus and it was placed at the center of the earth’s core, the valence electrons would be orbiting close to the surface of the earth’s crust • The valence electrons therefore are the part of the atoms that interact in chemical reactions to form compounds

17. Ionic Bonds • Form between a metal and a nonmetal • Metal tends to lose electrons which are transferred to the nonmetal • Metals form a cation (+) and nonmetals form an anion (-)

18. Ionic Bonds

19. Ionic Bonds Formation of NaCl Formation of MgF2

20. Ionic substances • These result in a lattice of ions rather than individual molecules, so we refer to MgF2 and NaCl as formula units, not molecules. • Properties of ionic substances: • Crystalline solids at room temperature • Hard and brittle • High melting and boiling points • Conduct electricity when in liquid form • Most are soluble in water

21. Covalent Bonds • Form between two nonmetals • Electrons are shared rather than transferred • Macromolecules and organic molecules are covalent molecules using covalent bonds, such as lipids, carbohydrates, proteins and nucleic acids.

22. Covalent bonds

23. Covalent Bonds Formation of H2 Formation of NH3

24. Covalent Bonds Formation of O2 – a double bond

25. Electronegativity • Linus Pauling developed the concept of electronegativity (En) • It is a measure of how strongly an atom attracts electrons to itself • Fluorine has the highest En value, and Pauling assigned it an arbitrary value of 4.1 • Elements to the left and below fluorine have decreasing En values

27. Electronegativity difference • In bond formation, it is useful to look at the electronegativity difference (ΔEn) • When Pauling looked at a range of bonds and their ΔEn values, a pattern was noticed • Bonds with an ΔEn between 1.7 – 4.1 tended to exhibit ionic characteristics • Bonds with an ΔEn below 1.7 tended to exhibit covalent characteristics HBr En (hydrogen) = 2.1 LiF En (lithium) = 1.0 En (bromine) = 2.8 En (fluorine) = 4.1 ΔEn = 2.8 – 2.1 = 0.7 ΔEn = 4.1 – 1.0 = 3.1

28. Trends in ΔEn

29. Polar and Nonpolar covalent bonds • There are two types of covalent bonds • Atoms that have the same En will have an ΔEn of zero. • These atoms will attract the shared electrons equally, and so the distribution of electrons is uniform • These are nonpolar covalent bonds

30. Polar and Nonpolar covalent bonds • Covalent bonds that have two different elements will have different En values and so the electron distribution will be non-uniform • These bonds are called polar covalent, since one end of the bond will be slightly electronegative (δ-)since the electrons are attracted more to the atom at that end

31. Molecular Shape • Valence Shell Electron Pair Repulsion theory (VSEPR) allows us to predict the 3-D shape of a molecule • VSEPR theory states that bond pairs of electrons repel one another, and lone pairs of electrons take up more space than bond pairs • There are four basic shapes which are common in organic molecules: • Linear • Bent or V-shaped • Tetrahedral • Pyramidal

32. Molecular shape • Linear or

33. Molecular shape • Bent

34. Molecular shape • Pyramidal

35. Molecular shape • Tetrahedral

36. Polarity of covalent molecules • How do we determine if a molecule is polar or nonpolar? • A polar molecule has an uneven distribution of electrons. This occurs when • There is at least one polar bond • The shape of the molecule is asymmetrical • Or the shape is symmetrical but the atoms surrounding a central atom have different En values

37. Polarity of Molecules - Examples Methane: CH4 VSEPR diagram: Polar bonds? Yes Overall dipole? No Methane is: NON-POLAR

38. Polarity of Molecules - Examples Ammonia: NH3 VSEPR diagram: Polar bonds? Yes Overall dipole? Yes Ammonia is: POLAR

39. Polarity of Molecules - Examples Water: H2O VSEPR diagram: Polar bonds? Yes Overall dipole? Yes Water is: POLAR

40. Polarity of Molecules - Examples Carbon dioxide: CO2 VSEPR diagram: Polar bonds? Yes Overall dipole? No Carbon dioxide is: NON-POLAR

41. Intermolecular Forces • The particle theory states that there are forces between particles, and the forces increase as the particles get closer. • These are the intermolecular forces • Compared to covalent and ionic bonds, they are very weak – but when there are many, they add up to a significant force • Collectively they are called van der Waals forces, but there are three different forces. • These forces have an effect on the boiling point and the solubility of substances.

42. London Dispersion Forces • London dispersion forces (LDF) occur when the protons in one atom or molecule attract the electrons in a neighbouring atom or molecule. • Since all particles have protons and electrons, all substances have LDF • Larger molecules have more protons and electrons, and so have greater London dispersion forces.

43. London Dispersion Forces When comparing the boiling points of hydrocarbons (non-polar molecules), we see that the boiling point increases as the number of carbons increases. Why is this?

44. Hydrogen bonding • Occurs only in polar molecules that have hydrogen and at least one of the following atoms: N, O or F. • These highly electronegative atoms have lone pairs of electrons which are attracted to the hydrogen atoms in neighbouring molecules. • These hydrogen atoms are essentially a proton

45. Dipole-dipole forces • Polar substances have a slightly electronegative end and a slightly electropositive end. • Dipole-dipole forces occur when oppositely charged poles momentarily attract one another

46. Water • Water is not an organic molecule but is essential for life on this planet • All cells are surrounded inside and out with water – anything that interacts with a cell must first be dissolved in water • Physical properties: • colourless and transparent • liquid at room temperature • density = 1.0 g/mL • m.p. = 0℃ b.p = 100℃ • water has LD, D-D forces, and H-bonding

47. Special Properties of Water • Water has cohesive properties – the high number of intermolecular forces causes water molecules to ‘stick’ together Examples: • surface tension – beading of water • water striders – too light to break surface tension • transpiration in plants – transport in xylem tubes

48. Special Properties of Water • Water has adhesive properties – it’s polar nature causes it to stick to other substances Examples: • capillary action – water ‘climbs’ up small diameter tubes, or ‘bleeds’ through the microscopic pores and channels in paper or other porous substances • this is due to the hydrogen bonding interactions between the water and the surface of the tube (either SiO2 or the cellulose tubes of paper) • This helps to explain the meniscus inside a tube

49. Cohesive-Adhesive Properties of Water

50. Cohesive-Adhesive Properties of Water