1. 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 nucleus is made up of protons (p + ) and neutrons (n o ) and is surrounded by rings of orbiting electrons (e - ) Subatomic particle Relative massRelative charge Proton1+1 Electron0 Neutron10

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

5.  What is the difference between these two atoms?

6.  Atoms in which the number of neutrons may differ  12 C and 14 C 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.  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 14 C decays, it becomes 14 N.  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

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

10.  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

11.  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

14.  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 Ca 2+ (used for muscle contraction), Na + and K + (nerve and muscle function), Fe 2+ and Fe 3+ (in hemoglobin) and H+ (required for synthesis of ATP)

15.  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

16.  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 (-)

17. Formation of NaClFormation of MgF 2

18.  These result in a lattice of ions rather than individual molecules, so we refer to MgF 2 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

19.  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.

20. Formation of H 2 Formation of NH 3

21. Formation of O 2 – a double bond

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

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

26.  There are two types of covalent bonds  Atoms that have the same E n will have an ΔE n of zero.  These atoms will attract the shared electrons equally, and so the distribution of electrons is uniform  These are nonpolar covalent bonds

27.  Covalent bonds that have two different elements will have different E n 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

28.  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

29.  Linear or

30.  Bent

31.  Pyramidal

32.  Tetrahedral

33.  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 E n values

34. Methane: CH 4 VSEPR diagram: Polar bonds? Yes Overall dipole? No Methane is: NON-POLAR

35. Ammonia: NH 3 VSEPR diagram: Polar bonds? Yes Overall dipole? Yes Ammonia is: POLAR

36. Water: H 2 O VSEPR diagram: Polar bonds? Yes Overall dipole? Yes Water is: POLAR

37. Carbon dioxide: CO 2 VSEPR diagram: Polar bonds? Yes Overall dipole? No Carbon dioxide is: NON-POLAR

38.  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.

39.  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.

40. 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?

41.  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

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

43.  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

44.  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

45.  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 SiO 2 or the cellulose tubes of paper) ◦ This helps to explain the meniscus inside a tube

51.  Water has outstanding solvent properties  Used to be called the ‘universal solvent’, but this is not a good name, since not everything dissolves in water  The polar nature of water allows any other polar substance or any charged particle to dissolve easily  The δ - will attract the δ + end of solutes, and this attraction will remain once the solute is dissolved.  The same is true for ionic substances – the cation will be attracted to the δ - end of water, and the anion will be attracted to the δ + of water.

53.  Water has a high specific heat capacity  This is a measure of the amount of heat energy required to increase the temperature of a 1g of a substance by 1℃.  c water = 4.18 J/g‧℃  This is high compared to other substances: c copper = 0.385 J/g‧℃c air = 1.00 J/g‧℃ c glass = 0.735 J/g‧℃c iron = 0.450 J/g‧℃  A metal pan absorbs heat energy quickly and loses it quickly. This makes metals useful for cooking.  Water takes more energy to heat up – thus the time it takes to boil water in a pot.

54.  Moderation of climate  This property of water also helps to moderate temperature changes in cells

55.  Water has a high latent heat of vaporization and fusion.  Latent heat is the energy absorbed or released by a substance during a change of state.  L f water = 334 J/g  L v water = 2260 J/g Special Properties of water

56. Latent heat

57.  Evaporative cooling relies on L v of water. Latent heat of vaporization

58.  Tender fruit farmers take advantage of the latent heat of fusion of ice when there is a chance of frost  On an evening when there is frost in the forecast, they spray water over their fruit, causing ice to form as the temperature drops below 0°C.  How does this help to protect the fruit? Latent heat of fusion

59.  Water’s density decreases as it changes from liquid to solid.  This is because the distance between molecules in a crystal lattice (as ice) on average further than when in a liquid. Special Properties of water

60.  Ionization occurs when 2 water molecules break apart into a hydronium ion H 3 O + and a hydroxide ion OH - 2H 2 O ⇌ H 3 O + + OH -  A substance that releases H + ions in solution is an acid.  A substance that releases OH - ions in solution is a base. Acids and Bases

61.  In pure water at 25 ℃, [H 3 O + ] = 1.0 x 10 -7 mol/L  pH is a measure of the concentration of hydronium ions [H 3 O + ]  pH = -log [H 3 O + ]  When an acid is combined with a base, a neutralization reaction occurs, with a salt and water as products. HCl + NaOH ⇌ NaCl + H 2 O Acids and Bases

62.  Strong acids and bases ionize or dissociate completely when dissolved in water.  A strong acid produces a high number of excess H + ions which combine with water to produce H 3 O + ions.  Weak acids and bases ionize only partly.  Only about 1.3% of weak acid’s molecules (such as acetic acid) will ionize in water, so contribute far fewer H + to the solution. Strong vs Weak acids and bases

63.  Hydrochloric acid 100% HCl → H + + Cl - Strong acids

64.  Acetic acid 1.3% CH 3 COOH ⇄ CH 3 COO - + H + Weak acids

65.  Biological systems tend to dislike significant shifts in pH. Our environment is full of weak acids and bases that can easily shift biological systems from their optimum pH. Buffer systems prevent these shifts.  Example: blood has an ideal blood pH of ~ 7.4  Below pH 6.8 or above pH 7.8, death occurs  CO 2 and H + are produced during cell respiration  When blood pH rises, carbonic acid dissociates to form bicarbonate and H +. H 2 C0 3 ⇌ HC0 3 - + H +  When blood pH drops, bicarbonate ions bind H + to form carbonic acid. HC0 3 - + H + ⇌ H 2 C0 3 Acid-Base Buffers