EH1008 : Biology for Public Health : Biomolecules

and Metabolism

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

Biochemistry:The chemistry of living things

2

• Obesity and type 2 diabetes

• Human Genome Project and ‘Personalised Medicine’• http://www.genome.gov/

• Stem Cell Therapy

Biochemistry & Physiology:Essential for understandingimportant epidemiological

issues:

3

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.

Some Key Concepts

4

Elements and atomsElement: 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

6

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

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!http://www.youtube.com/watch?v=lP57gEWcisY 8

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

Isotopes of elementsIsotopes: 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:

Atomic mass:is the average mass of naturally occurring isotopes10

(protium) (1/6420 H, 0.0156%)

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: 14C

– radiotherapy

– X-rays11

The MoleAvogadro’s Number: 6.022 x 1023 (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

Chemical BondsBondsbetween atoms within a moleculeoccur 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 chargedion, 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

Covalent bondsOne 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. O2

Polar covalent:electrons not shared equally because one nucleus attracts the electrons more than the other does, eg. H2O 14

Molecules, Compounds and SaltsMolecule: 2 or more atoms chemically combined, eg. O2

Compound: atoms of two or more different types chemically combined to form a molecule, eg. C2H5NO2

Molecular mass: the sum of masses of individual atoms or ions within a molecule or a salt,

eg. NaCl (22.99 + 35.45 = 58.44*)

H2O (1.01 + 1.01 + 16 = 18.01*)* not whole numbers because most elements have naturally occurring isotopes

15

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

Example of Intermolecular Forces:Hydrogen Bonds

• The positively charged H of one molecule is attracted to the negatively charged O, N or F of another molecule

• Eg., in H2O the positively charged hydrogen atoms of one H2O molecule bond with the negatively charged oxygen atoms of other H2O molecules

• Important role in determining the shape of complex molecules andhow such molecules interact with ions. 17

18

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 H20 is formed.

• synthetic reactions are responsible for forming the biological chemicals (biomolecules) characteristic of life:

proteins, lipids, carbohydrates and nucleic acids.19

Example of a synthetic reaction

-more on amino acids, peptides and proteins in Lecture 3

20

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

Monomers and polymers, example:

Glucose

Starch

Maltose

Salivary amylase

in plants

(in animals)

22

Decomposition Reactions: Catabolism

• Reactant broken down into two or more smaller

products.

• Chemical bonds are broken; energy is released.

• Hydrolysis: H2O is split into two parts that form part

of the reaction products.

23

Example of a decomposition reaction

24

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

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: CO2 and H+ formation in plasma

CO2 + H2O H2CO3 H+ + HCO3-

26

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

• 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 coupledto biological processes in order for work to take place.

28

ATP and Potential Energy

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

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

If Heat Output > Energy Intake

NEGATIVE ENERGY BALANCEe.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

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.

• Concentration of reactants.

As concentration of reactants increases, reaction speed increases.

For example, a decrease of O2 in cells can cause death as rate of aerobic (O2 -dependent) chemical reactions decreases.

34

• 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 proteins35

Activation Energy and Enzymes

36