Biology 2401 A & P I

Biology 2401 A & P I Chapter 2 Dr. Weis

Basic Chemistry : all living things are made of matter and require Energy to live

Concepts of Matter and Energy.....

Matter --> anything that occupies space and has mass

In chemistry, how building blocks are put together and interact.

Can exist in solid, liquid or gaseous state

Energy --> capacity to do work or put matter into motion

All living things are composed of matter and require Energy to grow and function.

A. Energy (E) Components

Total Energy = Potential Energy + Kinetic Energy

Potential E...stored energy, when released becomes Kinetic energy

Kinetic E ... energy of action or energy of work related to heat or thermal energy --

the hotter something is, the faster its molecules move

B. Energy Forms

* chemical energy --> E stored and released in chemical bonds

i.e. ATP --> ADP + P + E

* electrical energy --> movement of charged particles (e.g. electrons, ions)

i.e. electrical currents in nervous system & heart

use ions in solution called ELECTROLYTES

that move across cell membranes creating electrical events

* mechanical energy --> energy to move matter (e.g. movement of muscles)

* electromagnetic energy --> radiant E that travels in waves

i.e. X-rays, light, UV, radio, microwaves, infrared

Energy (E) Conversions

Remember that Energy can neither be created nor destroyed.

Conversions involve changing one form of energy to another (chemical---> mechanical)

process is inefficient, some E "lost" (unusable) to the environment as heat

Matter .....all matter is composed of fundamental substances called elements.

Four fundamental elements make up 96% of body and can be designated by a chemical shorthand called an Atomic Symbol.
The most common elements found in the body are listed below

1. carbon (C)

2. oxygen (O)

3. hydrogen (H)

4. nitrogen (N)

Other elements of importance are Sulphur (S), Phosphorus (P), Sodium (Na), Potassium (K),

Magnesium (Mg), Calcium (Ca), and Iron (Fe).

Each element is composed of particles called atoms which are unique to that element

and give it specific

physical properties....color, texture, freezing/boiling

chemical properties....bonding behavior

Chemical Level of Organization

I. Atoms and Molecules

Def: smallest chemical units of matter and can contain many subatomic particles

Three (3) fundamental subatomic particles:

1. Protons + electrical charge

2. Neutrons neutral

3. Electrons - electrical charge

Protons in the center of the atom ---> Nucleus

The number of protons is also known as the ATOMIC NUMBER.

A chemical element is a substance that has atoms with the same atomic number.

The number of protons = number of electrons, in the electrically neutral atom.

Neutrons in the center of the atom ---> Nucleus

do not change properties, except Atomic weight

can have a variable #, even with the same element. These are called isotopes and

refer to atoms whose nuclei contain different umbers of neutrons.

Radioisotopes emit subatomic particles (alpha, beta, gamma particles) that are used as tracers in Nuclear imaging.

Electrons circle around the nucleus and can occupy a series of energy levels (orbitals)

The outermost energy level will determine the atom's chemical properties

Each orbital has a maximum number of electrons it can contain

The atomic symbol is used for chemical shorthand to identify the element as well as its atomic number, mass number, and atomic weight.

1. Atomic number.... is equal to the number of protons in the nucleus.

This number is found as a subscript, to the left of the atomic symbol

H hydrogen 1 proton

O oxygen 6 protons

In an electrically neutral atom, the number of protons will equal the number of electrons

+ charges balance out - charges = neutral atom

Element Symbol	Element Name	# protons	# electrons	charge
He	helium	2 protons	2 electrons	neutral
K	potassium	19 protons	19 electrons	neutral
N	nitrogen	7 protons	7 electrons	neutral

to summarize.......

Atomic Number (AN) = protons (p)

Protons (p) = electrons (e) in the neutral atom

2. Mass number is the sum of its protons and neutrons

MN = p + n

electrons have a small mass, so are ignored

The mass number is found above the atomic number

He 2 = AN 4 = MN

MN = P + N

AN = P ..............MN = 2 + N

4 = 2 + N......4 - 2 = N, then 2 = N

C AN = 6, therefore 6 = protons

MN = 14 and MN = p + n

14 = 6 + n........14 - 6 = n.......8 = n

this is an example of an isotope of carbon

Isotopes of an element have the same atomic number, but vary in the number of neutrons.

Their atoms decompose spontaneously into more stable forms giving off

alpha, beta, or gamma radiation.

3. Atomic Weight

Very close to the mass number (protons + neutrons)

Will differ due to

1.) Isotopes (neutrons)

2.) Small weight difference between protons/neutrons

Let's take a look at some atoms from the periodic table

Oxygen : Atomic Number 8, therefore 8 protons

Mass Number 16 (protons + neutrons)

The Mole refers to every element that has a mass in grams equal to the atomic weight, will contain the same number of atoms.

The atomic weight of oxygen is 15.99

If we express this in grams 16.0 grams, then we will have one mole of oxygen.

If you will recall from chemistry, that one mole will contain

6.02 X 10 number of atoms. This is also known as Avogadro's number.

This will work for any atom, ion, molecule. For instance

Nitrogen (N)

Atomic number 7

Mass # 14

Atomic wt 14.007

........14g Nitrogen = 1 mole = 6.02 x 10 # of atoms

Potassium (K)

Atomic # 19

Mass # 39

Atomic Wt 39.098

...........39g Potassium = 1 mole = 6.02 x 10

C. Electrons and Energy Levels Remember in neutral Atoms that the # protons = # electrons

....electrically neutral + = -

proton number -----> are also the atomic #

Ok, so for Magnesium (Mg)

Atomic # 12

Mass # 24 THEN.........

# protons 12 (atomic number)

# neutrons 12 (Mass number - protons)

# electrons 12 (protons = electrons)

Remember Electrons have different Energy levels....ORBITALS

The maximum number of electrons in each energy level

is given by the formula

2n where n = 1, 2, 3 etc.

1st principal E level max 2 e (n = 1) 2(1) = 2x1 = 2

2nd " " max 8 e (n = 2) 2(2) = 2x4 = 8

3rd " " max 18e (n = 3) 2(3) = 2x9= 18

Each energy level can have a set of orbitals called s, p, d, f .

these orbitals have sublevels for electrons

s has (1)

p has (3)

d has (5)

f has (7)

and no more than 2 electrons can occupy each sublevel.

This means that s can have 2 electrons

p can have 6 electrons

d can have 10 electrons

and f can have 14 electrons

For all elements, their atoms want a full energy level to be happy and stable

The first energy level (1) can only have 2 electrons (2n ), then 2(1) = 2 x 1 = 2 {electrons}

therefore the s orbital fulfills it's needs first, and will only hold 2 electrons.

The second energy level can have 8 electrons. The s orbital isn't quite enough (only 2), but if we also

include the p level that can hold an additional 6 electrons, then we can meet our maximum

(2 + 6 = 8).

Chemistry people love short hand notation for things, so we can abbreviate the above discussion by saying

1s 2s 2p 3s 3p 3d etc. and using the superscript to let us know how many electrons we have

2 2 6 1

1s 2s 2p 3s (Remember, we fill the closest orbitals first before going on to the next energy level.)

Take a look at Hydrogen Atomic Number 1

protons = 1

electrons = 1

This electron will be in the first Energy level occupying the s orbital.

The shorthand notation for this is 1s .

Helium (an inert gas) has Atomic # 2

so, the electron number is 2. The first energy level can hold 2 electrons and the s orbital still

has one opening. So we fill the first energy level

(1s ) and the atom is happy and stable.

Oxygen (Atomic Number 8) will have 8 electrons to place in the orbitals of the energy levels

So it will look like this ............

2 2 4

1s 2s 2p if you add the superscripts

2 + 2 + 4 you will get 8 electrons. But remember that the p level can only hold 6 electrons

and it only has 4. This will make oxygen act in certain ways to get 2 more electrons to

complete the p orbital and have a full outer shell to be happy and stable.

Molecules and Compounds

Atoms will chemically combine with other atoms to achieve a full outer orbital shell.

In doing this, atoms can form MOLECULES.

Two or more atoms of the same element form a molecule of

that element.

2 hydrogen atoms form hydrogen gas

2 oxygen atoms form oxygen gas

two or more different atoms bind together to form a molecule of a compound

2 hydrogen atoms + 1 oxygen atom = water

In chemistry a solution is defined as a homogenous mixture of two or more components

1. solvent..... dissolving medium (usually a liquid)

2. solute ..... what is being dissolved

The concentration of a solution can be expressed as

a. percentage (%)

b. molarity (M)

II. Chemical Bonds

The energy relationship between electrons of reacting atoms

An Atom with a full outer shell is stable

Helium These are the inert gases, do not react with one another or combine with others.

Neon

Argon

Atoms w/ unfilled outer energy levels (valence shell) are unstable

to be stable------ they will need to :

1. share

2. gain electrons

3. loose

In order to form molecules : a chemical structure containing more than one atom

Form chemical compounds : combination of atoms of different elements

Types of Bonds

I. SHARED

A. Covalent bonds

1. Strong, electrons tie them together

2. electrically neutral, electrons spend equal time

around each atoms nucleus.

3. stable framework (i.e. Carbon)

4. sharing electrons (e) ======> Covalent bonds

sharing one pair of e ......single covalent bond

sharing two pairs of e .......double covalent bond

The best example of this is water.

Remember, hydrogen has one electron in its outer orbital and oxygen has 6.

Hydrogen could gain an electron and be happy and full.

Oxygen needs 2 electrons to be full, happy and stable.

Soooooo, 2 atoms of hydrogen can each share their electron with oxygen.

Each hydrogen has "2" electrons and the oxygen has it's "8".

But, oxygen isn't very nice, and doesn't share equally. In fact, it hogs the hydrogen's

electron, which creates a little Atension@ and makes the molecule polar

(having a positive charge at one end and negative charge at the other ).

These are slight charges, but do affect how the molecule orients itself.

II. Hydrogen Bonds

Hydrogen bonds occur when polar covalent bonds are formed

Since the atom no longer has the electron equally, this makes it slightly positive.

This slightly positive charge is attracted to a slightly negative charge and
forms a weak bond that can easily be broken, but it still can alter shapes of molecules such as proteins, DNA, and water.

III. IONIC Bonds

Ionic Bonds are the third type of chemical bond

It is a bond in which one atom has strong affinity for the electron an the other atom loses its hold

entirely. These create ions, (charged particles).

They are stable since their outer shell is full.

They do NOT share electrons but are now attracted to each other by the electrical charge.

An atom that gains an electron becomes negatively charged and is called an anion.

A minus - sign in the upper corner is used to denote an anion

Sodium has 1 electron in its outer shell. Chloride has 7.

[Rembember the rule of "8" for most valence shells]

Chloride has a strong attraction and will take the electron from sodium to fill its outer shell.
Sodium doesn't mind since getting rid of that valence electron will mean it too has a full and complete (stable) outer shell.

This means that chloride has 8 electrons, but still only has seven protons, therefore having an extra

negative charge. The chemical shorthand is

Cl - and is called chloride ion (anion).

Sodium has now lost it's electron, but still has protons ( + charge ) in the nucleus that are not

balanced so this makes the sodium +, commonly written as......

Na+ and is called sodium ion ( which is a cation)

These ions will still be attracted to each other by their opposite charge, even though they do not

share the electron in question.

Na+ Cl - ---> NaCl (a salt)

Most Ionic compounds will fall into the chemical category called SALTS.

III. Chemical Notation.......Chemical shorthand

1. Abbreviation of an element indicates 1 atom

2. Number proceeding abbrev indicates more

than one atom (2 O...two atoms of oxygen)

3. Subscript------molecule w/ that # of atoms

H one molecule with 2 atoms of hydrogen

4. Superscript indicates ionic charge

+ - ++

Na Cl Ca

IV. Chemical Reactions

a chemical reaction neither creates nor destroys, only will rearrange

and the number of atoms will be the same.

a chemical reaction will have ::

a. Reactants

b. Products

c. Arrow indicates direction

types of chemical reactions :

a). Decomposition.....break down molecule into smaller fragments. Energy is released

when bond is broken.

AB -----> A + B

b). Synthesis.....assemble larger molecule from smaller fragments

A + B -----> AB requires Energy

c). Exchange or displacement

AB + CD -----> AC + BD

in this reaction, break old bonds and make new ones.

Can be exergonic , releasing Energy or endergonic, absorbing Energy.

Oxidation/Reduction {REDOX} reactions that involve electron transfer

(e.g. cellular respiration and the Krebs Cycle)

d). Reversible Reactions

A + B <=====> AB

will reach an equilibrium when the two rates are in balance.

The rate of the chemical reaction depends on

1. Temperature

2. particle size

3. concentration

4. enzymes

5. energy requirements

BIOCHEMISTRY

V. Inorganic compounds

do not contain Carbon or Hydrogen atoms as primary structure.

are usually small molecules held together by ionic bonds.

Examples........Acids, Bases, salts, water

A. WATER 2/3 body weight

60% - 80% of volume of living cells

properties due to hydrogen bonding and polar covalent bonds

1. Temperature

takes a lot of Energy to break bonds to make a liquid into a vapor (vaporization)

Absorbs and releases large amount of heat before changing temperature and therefore prevents sudden changes in temperature.

2. An effective solvent

acts as transport medium

causes dissociation of ionic bonds and the ions in solution can carry electrical current (electrolytes).

This is important in muscle and nerve conduction.

3. Chemical reactant, involved in hydrolysis and dehydration synthesis

4. pH (hydrogen ion concentration) of body fluids

In one liter of pure water, a small number of molecules will dissociate giving.........

H 2O <=====> H + + OH-

Equal numbers of hydrogen ions and hydroxyl ions were found, .0000001 moles of each.

This can be expressed as

[H+ ] = 1 X 10(-7 power) moles per liter

The pH is the negative exponent of the hydrogen ion concentration expressed in moles per liter.

Or pH = - log [H + ] {remember this definition}

Therefore the pH of pure water is - (-7) or 7 and is considered neutral.

The pH scale ranges from 0 to 14.

pH of 0 means [H +] = 10 or 1.0

pH of 6 means [H+ ] = 10 to -6 power or .000001

pH of 11 means [H+ ] = 10 tp -11 power or .00000000001

The pH of blood ranges from 7.35 (venous) to 7.4 (arterial).

A low pH corresponds to high [H +] or acidosis

A high pH corresponds to low [H +] or alkalosis

To prevent acidosis or alkalosis the body uses

1. Acid/Base buffer systems (seconds)

use weak acids and weak bases

a. Bicarbonate Chemical Buffer system

b. Phosphate Chemical Buffer system

c. Protein Buffer system

2. Changes in respiration and CO2 (minutes)

3. Renal/urine excretion (hours)

B. Acids, Bases, Salts

Acids release H+ and shift pH down, making the solution more acidic because there is more free H+

are considered proton donor. A strong acid will completely dissociate in a one way reaction.

For example

HCl ------> H+ + Cl-

other acids H2CO3- (carbonic acid) are weak acids

and help to control pH by controlling free H+

Bases are solutes that removes H+

liberate OH ions

a strong base will completely dissociate in a one way reaction,

NaOH -------> Na+ + OH-

other bases HCO3- (bicarbonate) are weak bases

Salts inorganic compound created by the reaction of a strong acid with a strong base

held together by ionic bonds and dissociate in water.

These ions in solution are called electrolytes. Electrolytes create an electrical current when they move across cell membranes.

Common salts : NaCl, KCl, Ca2CO3

Electrolytes : Na+, Cl-, K+, Ca++

Buffers are compounds that maintain pH within normal limits by removing or replacing hydrogen ions

primarily use weak acids and weak bases, that fail to completely dissociate.

H2CO3 <====== > HCO3- + H+

VI. Organic Compounds contain Carbon, Oxygen, Hydrogen

Carbon is important because ::

1. shares electrons forming covalent bonds

2. 4 bonds

4 classes of Organic Compounds (Macromolecules)

1. Carbohydrates

2. Lipids

3. Proteins

4. Nucleic Acids

Carbohydrates (CH2O)

Energy source.....cellular fuel, provide some structure for DNA, cell membranes

CH2O will vary in size and therefore solubility, as larger molecules are less soluble in water.

Sugars and starches

C : H : O in a 1 : 2 : 1 ratio

Sugars ......

Monosaccharide one sugar (glucose)

3-7 C, chain or ring

Disaccharide means two monosaccharides (sucrose, lactose, maltose)

Polysaccharide three or more monosaccharide

1. glycogen

2. starches (plants), these larger molecules are used for storage

Lipids

involved in Energy reserve, structure and temperature

insoluble in water, but readily dissolve in other lipids

examples are fats, oils, waxes

Classes

1. Fatty Acids R - COOH

the R group is composed of carbon chains of variable length

can be saturated (four single covalent bonds for the carbon chain)

or unsaturated C == C

polyunsaturated refers to multiple C == C

2. Glycerides

glycerol + Fatty Acid

glycerol (a sugar alcohol) has three sites for Fatty Acids to bind.

Can make mono-, di-, triglycerides, the neutral fats.

3. Prostaglandins (PG)

belong to a larger group of lipids called EICOSANOIDS

a five carbon ring with 2 short chain fatty acids

there are 6 types : A, B, C, D, E, F with various degrees of saturation (1,2,3)

involved in uterine contraction

control of inflammation

regulation of body temp

regulation of stomach acid

regulation of gi motility

example.... PGF 2 alpha (uterine contraction)

4. Steroids found in diet and body synthesis by the liver

four ringed structure made of hydrocarbon, fat soluble, very little oxygen

based on cholesterol which is involved in

cell membranes

steroid hormones : testosterone, estrogen, corticosteroids

bile salts

adrenal cortical hormones

5. Phospholipids/Glycolipids

synthesized from fatty acids and glycerol, with an added phosphate

(phospholipid)or carbohydrate (glycolipid) group.

Phospholipids have a phosphorus group and two fatty acid chains which cause

specific arrangement of their polar and non-polar ends

Phospholipids important in cell membranes and their arrangement in water to

form micelles.....fatty acid tails to the inside, and glycerol and phosphate heads to the outside

C. Proteins

Chains of amino acids linked by peptide bonds

Consist of Carbon, Hydrogen, Oxygen, Nitrogen

Make up 20% of total body weight.

Several functions:

(I) Fibrous

1. Structure/Support

2. Movement/Contraction (muscles)

3. External Defense by keratin of the skin

(II) Globular

4. Transport --- minerals, hormones, hemoglobin

5. Metabolic.......enzymes

6. Defense Internal by immune system (immunoglobulins)

Amino acids (AA) are the building blocks of proteins

Consist of a carbon with four bonds to ::

a. Hydrogen

b. amino group (NH2)

c. COOH (acid)

d. Side chain (R group) makes each amino acid different and unique in

chemical behavior because of different number and arrangement of atoms.

There are 20 amino acids. Some are manufactured by the body, others are required in the diet.

The peptide bond is between

COOH of one AA and NH of another AA

2 amino acids........dipeptide

3 amino acids.........tripeptide

many amino acids.......Short chain peptides and long chain polypeptide or PROTEIN

The shape of proteins are based on structure and will determine its function.

1. Primary Structure is the sequence or order of amino acids. Coded by the DNA and copied

and interpreted by the RNA.

2. Secondary structure

hydrogen bonding between the NH and CO groups to form alpha helix (coil)

twisting and folding upon themselves

some chains can link together by their hydrogen bonds in a side to side manner (not coiled)

and this is called a beta pleated sheet.

3. Tertiary Structure

coiling and folding due to the interactions of R groups into a ball-like or globular molecules

4. Quaternary structure

interaction between polypeptide (protein) chains

The structure of a protein dictates its function ::

Fibrous proteins ....linear, strand-like, stable, insoluble in water

function : support, strength

ex : collagen, keratin, elastin, muscle protein

Globular proteins.....compact and spherical chemically active @ certain sites, water soluble

ex : antibodies, hormones, enzymes

Enzymes....globular proteins with active sites

biological catalysts increase speed of reactions by lowering the activation energy

highly specific and control single chemical reaction

named for the type of reaction they catalyze. Usually end in - ase

may need co-factors (Vitamins)

have specific pH and temperature for for optimal function

If protein shape is altered, especially globular proteins, then the functional properties will change.

This primarily involves the tertiary and quaternary structures and hydrogen bonding.

These can be changed by increases in

* temperature

* pH

Denaturation occurs when bonds are broken :

proteins unfold an loose their 3-D shape.

May be reversible, unless the temperature or pH change is extreme.

D. Nucleic Acids

Two classes of nucleic acids

1. DNA (deoxyribonucleic acid)

2. RNA (ribonucleic acid)

Basic functions are to store and process information at the molecular level.

DNA in the nucleus of Eukaryotic cells and is free floating in Prokaryotic cells

codes for enzymes and protein synthesis thereby regulating cellular metabolism,

Cell structure and shape

RNA helps to manufacture proteins. Three types : rRNA, tRNA, and mRNA

Nucleic Acids are a series of nucleotides, composed of C, O, H, N, P

These nucleotides consist of

purine is a two ringed structure found in DNA and RNA

Nitrogen bases are Adenine (A) and Guanine (G)

Pyrimadine are single ringed structures

Nitrogen bases are Cytosine (C) found in DNA and RNA

Thymine (T) found in DNA only

Uracil (U) found in RNA only as it takes the place of Thymine

RNA is single chain

distributed throughout cytoplasm and nucleus

contains ribose sugar and bases A, U, C, G

DNA consists of 2 nucleotide chains attracted by hydrogen bond of opposing nitrogen bases

creating a double helix described as a twisted ladder

complimentary base pairing :

A = T (DNA) A = U (RNA)

G = C (DNA, RNA)

High Energy compounds are molecules whose structure includes high energy bonds.

The bonds store energy formed by phosphorylation, therefore when the bonds are broken,

Energy is released.

Examples of High Energy Compounds include

1. Adenosine Triphosphate ATP

2. Guanosine Triphosphate GTP

3. Uridine Triphosphate UTP

Adenosine monophosphate AMP

a nucleotide (adenine, ribose sugar, and has one phosphate group).

To that, two more phosphate groups are added to form ATP.

AMP + P and ENERGY = ADP

ADP + P and ENERGY = ATP

Function of ATP ::

1. Energy for transport of substances across cell membrane

2. Energy for mechanical work (muscle contraction)

3. Energy requiring chemical reactions (endergonic)