EH1008 : Biology for Public Health : Biomolecules and Metabolism

Transcription

1 EH1008 : Biology for Public Health : Biomolecules and Metabolism

2 Biochemistry: The chemistry of living things What has this got to do with Epidemiology & Public Health? Aims of 'Epidemiology & Public Health: '...prevention of disease, the prolonging of life and the promotion of health through the organised efforts of society.' Understanding of life and its underlying mechanisms is essential for the execution of these aims, including communication of this understanding to the general public 2

3 Biochemistry & Physiology: Essential for understanding important epidemiological issues: Obesity and type 2 diabetes Human Genome Project and Personalised Medicine Stem Cell Therapy 3

4 Some Key Concepts Matter (harder to define than you might think): anything that has mass and volume. Density = mass / volume Mass: the amount of matter in an object. Weight: the force exerted by gravity on an object of a certain mass. 4

5 Elements and atoms Element: the simplest types of matter that possess their own unique set of chemical properties. Composed of just one type of atom. To date, there are 118 known elements. 5

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7 Atoms Atom: smallest particle of an element that has chemical characteristics of that element. Atoms can be broken down into smaller components, but these components no longer have the properties of that element. 7

8 What are atoms made of? Nucleus: composed of neutrons, that have no electrical charge and protons, that have one positive charge. Surrounded by an electron cloud. Electrons have one negative charge. Mainly empty space! 8

9 Atom Number and Mass Number Atomic number = number of protons in an atom = number of electrons in an atom. Mass number = number of protons + neutrons in an atom. 9

10 Isotopes of elements Isotopes: element that have the same number of protons and electrons (i.e. chemical properties) but different numbers of neurons (which determine some physical properties). For example, isotopes of hydrogen: (protium) (1/6420 H, %) Atomic mass: is the average mass of naturally occurring isotopes 10

11 Radioactive Isotopes (Radioisotopes) Some isotopes (both natural and 'man-made') release energy in the form of radiation, eg. gamma rays. Radioactivity can be measured: this has numerous applications in research and in medicine: Examples of uses: carbon-dating: 14 C radiotherapy X-rays 11

12 The Mole Avogadro s Number: x (6022 with 20 zeros after it!). Mole (M): Avogadro's number of atoms, molecules or ions. Molar mass = the mass of one mole of a substance in grams = atomic mass. Examples: 1 M present in 1 g of hydrogen; 12 g of carbon and g of oxygen 12

13 Chemical Bonds Bonds between atoms within a molecule occur when the outermost electrons of an atom are shared with, or transferred to another atom. Ionic bonds involve transfer of electrons. If electrons are lost, the atom becomes a positively charged ion, or cation, eg. Na + If electrons are gained, the atom becomes a negatively charged ion, or anion, eg. Cl - In ionic bonding, cations and anions attract each other and stay close together, eg. NaCl 13

14 Covalent bonds One or more pairs of electrons are shared between atoms: Single covalent bond: 2 atoms share 1 pair of electrons Double covalent bond: 2 atoms share 4 electrons Nonpolar covalent: electrons shared equally because nuclei attract the electrons equally, eg. O 2 Polar covalent: electrons not shared equally because one nucleus attracts the electrons more than the other does, eg. H 2 O 14

15 Molecules, Compounds and Salts Molecule: 2 or more atoms chemically combined, eg. O 2 Compound: atoms of two or more different types chemically combined to form a molecule, eg. C 2 H 5 NO 2 Molecular mass: the sum of masses of individual atoms or ions within a molecule or a salt, eg. NaCl ( = 58.44*) H 2 O ( = 18.01*) * not whole numbers because most elements have naturally occurring isotopes 15

16 Intermolecular Forces Weak forces between molecules Caused by weak electrostatic (electrical charge) attractions between oppositely charged parts or molecules, or between ions and molecules Weaker than forces producing chemical bonding 16

17 Example of Intermolecular Forces: The positively charged H of one molecule is attracted to the negatively charged O, N or F of another molecule Hydrogen Bonds Eg., in H 2 O the positively charged hydrogen atoms of one H 2 O molecule bond with the negatively charged oxygen atoms of other H 2 O molecules Important role in determining the shape of complex molecules and how such molecules interact with ions. 17

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19 Synthetic Reactions: Anabolism two or more reactants combine to form a new, larger product. chemical bonds (covalent bonds) formed. energy stored in these bonds. used in growth, maintenance, repair & energy storage. dehydration: a synthetic reaction in which H 2 0 is formed. synthetic reactions are responsible for forming the biological chemicals (biomolecules) characteristic of life: proteins, lipids, carbohydrates and nucleic acids. 19

20 Example of a synthetic reaction -more on amino acids, peptides and proteins in Lecture 3 20

21 Monomers and polymers Synthetic reactions are involved information of long chains of repeated units, known as polymers. The smaller units from which polymers are formed are called monomers. Glucose and amino acids are examples of monomers. Glycogen and starch; and polypeptides and proteins are polymers of these molecules. 21

22 Monomers and polymers, example: Glucose in plants Starch Salivary amylase (in animals) Maltose 22

23 Decomposition Reactions: Catabolism Reactant broken down into two or more smaller products. Chemical bonds are broken; energy is released. Hydrolysis: H 2 O is split into two parts that form part of the reaction products. 23

24 Example of a decomposition reaction 24

25 Metabolism The sum of all of the catabolic and anabolic reactions taking place in a living thing. The set of chemical reactions occurring in living organisms that are used to maintain life. 25

26 Types of reaction Reversible reactions: can start from the reactants and proceed to the products, or the other way around. Equilibrium: rate of formation of the products equals the rate of formation of the products. Example: CO 2 and H + formation in plasma CO 2 + H 2 O H 2 CO 3 H + + HCO 3-26

27 Chemical Reactions & Energy Energy is the capacity to do work All living things need a source or sources of energy in order to exist. For most life on Earth, the primary source of energy is the Sun; this energy is trapped by photosynthetic reactions of plants, which are then consumed by animals, etc. 27

28 The ordered state of living things is maintained at the cost of a constant input of energy. No energy Input Death Energy Intake = Food Intake If the food taken in was just 'burnt', the energy it contains would be released just as heat. Although heat production is important for the control of temperature in animals, the energy contained in food must be coupled to biological processes in order for work to take place. 28

29 ATP and Potential Energy Hydrolysis of ATP is a key mechanism coupling energy to work in the body. 29

30 Heat Energy When chemical bond are broken and energy is released, only some of that energy is used to make ATP. Suprisingly inefficient. more than 99% efficient less than 20% efficient? The remaining energy is released as heat. Used to maintain body temperature in mammals. 30

31 If Heat Output > Energy Intake NEGATIVE ENERGY BALANCE e.g. Starvation or disease (fever) If Heat Output < Energy Intake POSITIVE ENERGY BALANCE e.g. Growth, Pregnancy, Energy Storage getting fat! Obesity is a growing health concern in the Western World 31

32 Speed of Chemical Reactions The speed (rate) of a chemical reaction depends on many factors: Temperature Concentration of reactants. Catalysts 32

33 Temperature: affects speed of reaction. Increase in temperature means increase of kinetic energy (in living things, motion of molecules). Molecules move faster, collide harder and more often. Makes it more likely that they will react. 33

34 Concentration of reactants. As concentration of reactants increases, reaction speed increases. For example, a decrease of O 2 in cells can cause death as rate of aerobic (O 2 -dependent) chemical reactions decreases. 34

35 Catalysts: substances that increase the speed of chemical reactions without being permanently changed or depleted Enzymes: proteinaceous, biological catalysts that decrease the activation energy necessary for reaction to begin. Living things contain many types of enzyme. Activation Energy: minimum energy reactants must have to start a chemical reaction. -more in Lecture 3, as all enzymes are proteins 35

36 Activation Energy and Enzymes 36