Chapter 2 the Chemical Basis of Life

Chapter 2 – The Chemical Basis of Life I. Biological function starts at the chemical level (Fig. 2.1)

A. Hollistic vs. Reductionist Approach

1. Reductionists try to explain the universe based on it’s smallest pieces

2. The holistic approach focuses on the big picture and all its parts working together.

3. It is likely that both approaches will be required to understand the natural world in its entirety.

B. Different approaches, but agree on the following

1. Each level in the hierarchy builds on the one below it (actin to muscle) and with each step upward, novel properties emerge that were not present at simpler levels of organization and new functions arise – EMERGENT PROPERTIES!!!

2. Structure and function are related at every level (A hammer is for nails and a screwdriver is for screws). STRUCTURE (FORM) – FUNCTION RELATIONSHIP II. Life requires about 25 chemical elements

A. MATTER – anything that occupies space and has mass, composed of atoms

B. An “Atom” is the name for the smallest CHEMICALLY UNBREAKABLE unit of matter. There are many types of atoms

C. CHEMICAL ELEMENT – pure substance with only one type of atom (a substance that cannot be broken down to other substances by ordinary chemical means (They are atom types))

D. 92 Naturally occurring elements

1. 25 are essential for life (Table page 19) – know the top 4 in order. 2. Trace elements – essential to at least some organisms (Iron needed by all organisms)

3. Goiter III.Elements may combine to form compounds

A. COMPOUND – substance containing 2 or more elements in a fixed ratio

B. ATOM – smallest unit of matter that still retains the properties of an element (diameter = ~0.1nm) – Define nano at this time

C. Different combinations of atoms determine the unique properties of each compound. (EMERGENT PROPERTIES)

D. Molecule – group of at least two or more atoms joined together by strong covalent bonds. H2 is a molecule, but it is also an element (one atom type). CO2 is a molecule, but it is a compound (two types of atoms here). IV. Atoms consist of protons, neutrons, and electrons

A. Subatomic particles (discuss EM force and other three briefly)

1. Protons and neutrons occupy the center region of the atom and form the atomic nucleus

2. electrons occupy the region surrounding the nucleus

3. Protons and neutrons are very similar particles with the same mass of 1 amu (1.67 x 10-27 Kg). Protons have a charge of Plus 1 and neutrons have a charge of zero (neutral).

4. Electrons have a much (2000x) smaller mass than p and n and a charge of -1. 5. Mostly empty space!! Nucleus diameter ~ 10-15m (1/1,000,000,000,000 mm) and atom diameter is ~10-10m (1/10,000,000th of a mm) (so nucleus diameter is ~10,000X smaller than atom!!) ANALOGY – nucleus is a golf ball, first shell would be at a distance of 1Km (0.62 miles), second shell at 4Km, third at 9Km, etc…!! or, if nucleus were this period (.) then first shell would be 50m away!! 99.9999999999999% empty

http://www.phrenopolis.com/perspective/atom/

6. We are just empty space and some dots of mass!! Why don’t we just pass right through things (like the floor right now!!!)

7. So how many water molecules do you think might be in ~18ml of water?

8. How big is this number?

a) If you had this many marshmallows, they would cover the earth six miles deep

b) If you had this many dollars, and you were spending 1 billion dollars per second, it would take you 19 million years to spend it all!

c) Avogadro's number of pennies placed side by side would stretch for more than a million light years (5.9 trillion miles)!

B. Differences in Elements

1. What makes an element an element? What makes carbon, carbon? Atomic Number = # of protons an atom has. The atomic number defines the elemental makeup. Carbon always has 6 protons. So any atom with 6 protons will be carbon – no exception!!

2. The mass number of an atom is determined by adding the number of protons and neutrons – why ignore electrons?? 3. Charge is determined by subtracting the number of positive protons from the number of negative electrons.

C. Isotopes

1. atoms with the same atomic number (same element), but different masses (different number of neutrons). Do isotopes have the same chemical properties?

2. Not all isotopes are happy (stable). The unhappy isotopes strive to become stable by releasing radioactivity (nuclear radiation) – particles and energy given off by the nucleus. V. Radioactive isotopes can be helpful and harmful!!

A. Basic research – we can label molecules like CO2 by using a radioactive carbon atom. Now we can follow the path of carbon (used to study photosynthesis).

B. Medical diagnosis – (Nuclear Medicine) – We can inject you with radioactive sugar and do a PET (positron emission tomography) scan to find tumors.

C. Dangers – radioactive emission can damage your own molecules (especially DNA) resulting in an increased incidence of cancer. (Discuss natural sources of radiation) VI. Electron arrangement determines the chemical properties of an atom

A. Electrons occur at certain energy levels called electron shells (Bohr Model).

B. Each shell holds a maximum number of electrons (go over Bohr rules).

C. Electrons in the outermost shell determine the chemical properties of the atom. Why do you think that would be?

D. Atoms are happy when their outermost (valence) shell is FULL!! (Like me with food). a) Most atoms alone do not have filled shells. This will cause them to react with other atoms to try and get some electrons – time to make some bonds!! VII. Chemical bonding – sticking atoms together – three major ways to do this

A. Ionic bonds - attractions between ions of opposite charge

1. What’s an ion? An atom or molecule with an electric charge caused by having an unequal number of protons and neutrons.

2. Ionic bond – two atoms of opposite charge held together by the EM force. (show how to draw this)

3. The word “salt” means ionic compound (it is not just table salt!). Ex. NaCl, KCl

B. Covalent bonds (stronger than ionic) – sharing of electrons to form molecules (show how to draw and compare to ionic drawing)Ex. H2, H2O, O2, CO2, etc…

1. Two atoms share one or more pairs of outershell (valence) electrons

2. Two or more atoms held together by a covalent bond form a molecule.

3. Double Bond – the sharing of two pairs of electrons between two atoms (stonger than a single bond)

C. Hydrogen Bonds – discussed below VIII. THE PROPERTIES OF WATER

A. Water is a polar molecule

1. Electronegativity – an atoms attraction for the shared electrons 2. Nonpolar molecules – have equal sharing of electrons between atoms – covalent bonds are balanced (H2, O2, CH4)

a) Nonpolar covalent bonds

3. Polar Molecules – An unequal sharing of electrons between atoms due to different attractions by each atom resulting in one part of molecule being slightly positive and the other part slightly negative. Ex. H2O, CH3COOH

a) polar covalent bonds

B. Overview: Water’s polarity leads to hydrogen bonding and other unusual properties

1. Hydrogen Bond – special type of bond that exists b/w an electronegative atom and a hydrogen atom that is bound to another electronegative atom.

2. Water can make 4 H-bonds (possibly to 4 other waters)

3. H-bonding and polarity of water makes it very special and unique (only substance to exist in all 3 states (solid, liquid, gas) on Earth).

C. Hydrogen bonds make liquid water COHESIVE

1. Cohesion – tendency of molecules to stick together – water is highly cohesive – important for life (ex. Water droplets/RainX, transportation thru plants)

2. Surface Tension - measure of how difficult it is to break or stretch the surface of a liquid – so strong it can support insects like a trampoline.

D. Adhesive

E. Water’s hydrogen bonds moderate temperature

1. HEAT - form of energy associated with the movement of the atoms and molecules

2. Temperature – measurement of the intensity of the heat (average speed of molecules). 3. Water can absorb and store a large amount of heat while only increasing a few degrees in temp.

a) Heat absorbed to first break H-bonds, then to make molecules move faster. (Ex. Long Island Temp.)

4. Evaporative cooling – when a substance evaporates, it takes the heat energy with it leaving the surface it left cooler. (sweating, alcohol bath) – molecules with greatest energy leave first.

F. Ice is less dense than liquid water

1. When H-bonds form between water molecules to make ice, the water molecules spread out.

2. As a liquid, water molecules are closer together

3. Results in less molecules per unit volume for the ice (less dense) – ice floats!

G. Water is a versatile solvent

1. Solution – a homogeneous (same throughout) mixture of two or more substances (Ex. salt water)

2. Solvent – the dissolving agent (liquid) (Ex. water)

3. Solute – that which is being dissolved (can be gas, liquid or solid) (Ex. the salt like NaCl)

4. Aqueous Solution – one in which water is the solvent

5. Water has been termed the “universal solvent” because of its polarity – it can interact with and dissolve MANY other charged species.

H. The Chemistry of Life is Sensitive to acidic and basic conditions

1. Most water in an aqueous solution exists as H2O. However, a very small amount breaks apart to form H+ (protons) and OH- ions.

2. Other chemicals can release protons (or grab them!) 3. A chemical that release H+ (protons) into solution is called an ACID

4. A chemical that accepts (removes) H+ from solution is called a base (or alkali).

5. We use a pH scale to measure the acidity (number of “free” protons) in a solution.

a) It uses a scale from 0 to 14 with 0 being the most acidic (many free protons) and 14 being the most basic (few free protons).

b) A pH=7 is considered neutral (same number of H+ and OH- ions in solution)

c) A change of one unit in the scale is equal to a 10- fold change in H+ concentration. Ex. Solution with pH=5 has 1000X higher H+ concentration as a solution with pH=8

d) Buffers – molecules that resist change in pH by grabbing and releasing protons. Different molecules for different pH values. (swimming pools, fish tanks, blood)

6. Strong acid vs. strong base

a) Both completely ionize

(1) strong acid = HCl (hydrochloric acid)

(a) HCl  H+ + Cl-

(2) Strong Base (ex: NaOH = sodium hydroxide)

(a) NaOH  Na+ + OH- 7. Acid precipitation threatens the environment

a) Sulfur and nitrogen compounds are emitted from the combustion of fuels and react with water in the atmosphere to form acids.

b) These acids fall back down to Earth in the form of rain, snow, dew or dry particles.

c) Organisms that are adapted to neutral pH will be harmed by acid rain IX. Chemical Reactions Rearrange Matter

A. Chemical reaction – process leading to chemical change in matter. 2H2 + O2  H2O

1. Identify the reactant and products of the above reaction

2. Equation is balanced: same # of H and O on either side of arrow.

3. Organisms can’t make water (that would be nice), but we do 100,000’s of other reactions to rearrange matter to our benefit.

a) C40H56 + O2 + 4H  2C20H30O

Beta-carotene  Vitamin A

4. Metabolism – all the chemical reactions occurring within our bodies or within any organism for that matter.