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2

2.1 Basic Chemistry

• Matter is anything that takes up space and

has mass.

– The three states of matter are solid, liquid, and gas.

• All matter, living or nonliving, is made up of elements.

– Elements are substances that cannot be broken down into simpler substances by ordinary

chemical means.

3

• 92 naturally-occurring elements serve as building blocks of all matter.

• Other elements have been “human-made” and are not biologically important.

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

organisms

Fe Ca K S P Si Al Mg Na O N C H

Element

60

40

20

0

Pe

rce

nt b

y W

eig

ht

Earth’s crust

© Tom Mareschal/Alamy RF

2

4

Elements That Make up 95%

of Organisms (by weight)

– C Carbon

– H Hydrogen

– N Nitrogen

– O Oxygen

– P Phosphorus

– S Sulfur

5

Atomic Structure

• An atom is the smallest part of an element

that displays the properties of the element.

• Atoms are made up of subatomic particles.

– Protons-positively charged, found in nucleus

– Neutrons-uncharged, found in nucleus

– Electrons-negatively charged, move around nucleus

6

Helium (He)

Subatomic Particles

= proton

= neutron

= electron

b.a.

Particle

Proton

Neutron

Electron

Nucleus

Nucleus

Electron orbital

+1

0

–1

1

1

~0

Atomic Mass Location

c.

Electric

Charge


Figure 2.2

3

7

Atomic Mass = Number of Protons +

Number of Neutrons

Atomic Number = The Number of Protons

in the Nucleus

C6

mass number

atomic number

atomic symbol12


8

The Periodic Table

• Even though each element consists of a

different atom, certain chemical and physical properties recur (periodicity).

• The periodic table is used to group the elements according to these characteristics.

• The vertical columns are groups.

• The horizontal rows are periods.

9

The Periodic TableCopyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

Groups

I

1

1.008

H

2

4.003

He

II III IV V VI VII

3

6.941

Li

4

9.012

Be

5

10.81

B

6

12.01

C

7

14.01

N

8

16.00

O

9

19.00

F

10

20.18

Ne

22.99

Na

12

24.31

Mg

13

26.98

Al

14

28.09

Si

15

30.97

P

16

32.07

S

17

35.45

Cl

18

39.95

Ar

19

39.10

K

20

40.08

Ca

31

69.72

Ga

32

72.59

Ge

33

74.92

As

34

78.96

Se

35

79.90

Br

36

83.60

Kr

VIII

Periods

11

4

10

Isotopes

• Isotopes are atoms of the same element

with different numbers of neutrons.

• Radioactive isotopes emit various types of

energy as they decay.

C C C*12 13 14

6 6 6

*radioactive

11

• Uses of low level radiation

• a. The missing area in this thyroid scan indicates the presence of a tumor that does not take up the

radioactive iodine.

• b. A PET (positron emission tomography) scan reveals which portions of the brain are most

active (yellow and red colors).

12


larynx

thyroid gland

trachea

b.

a.

a: © Biomed Commun./Custom Medical Stock Photo; b(left): © Mazzlota et al./Photo Researchers, Inc; b(right): © Hank Morgan/Rainbow

5

13

• Uses of high levels of radiation

– Radiation kills bacteria and fungi. Irradiated peaches spoil less quickly and can be kept for

a longer length of time.

– Physicians use radiation therapy to kill cancer cells.


a. b.© Mark Kostich/Getty RF

14

Electrons

• In an electrically neutral atom, the positive charges of protons in the nucleus are balanced by negative charges of electrons.

• Electrons move around the nucleus in orbitals.

• Electrons move in energy levels (electron orbitals).

– First contains two electrons.

– Every one after that can hold eight electrons.

• Octet rule

15


electron

electron orbital

nucleusH

C

carbon12C

6

hydrogen1H1

Figure 2.6

6

16

N

7

nitrogen14N

O

Oxygen16O8


Figure 2.6

17

P S

Sulfur32S16

Phosphorus31P15


Figure 2.6

18


electron

electron orbital

nucleusH

C

O

N

P S

7

nitrogen14N

carbon12C

6

Sulfur32S16

Oxygen16O8

Phosphorus31P15

hydrogen1H1

Figure 2.6

7

19

2.2 Molecules and Compounds

• Molecules form when two or more of the same elements bond together (example: O2).

• Compounds form when two or more different elements bond together (H2O).

• When a chemical reaction occurs, energy may be given off or absorbed because of the energy present in bonds.

20

Ionic Bonding

• Ions form when electrons are transferred

from one atom to another.

• For example:

– Na, with one electron in its 3rd orbital, tends to be an electron donor.

• Becomes positive after giving up one electron

– Cl, with seven electrons in its 3rd orbital, tends to be an electron acceptor.

• Becomes negative after gaining one electron

21

Ionic Bonding

• After the transfer of electrons between Na

and Cl:

– Both the Na and Cl ions have eight electrons in their outer orbitals.

– Ions now have opposite electrical charges.

• Ionic compounds are held together by an attraction between oppositely charged ions

called an ionic bond.

8

22

• Each atom now has 8 electrons in its

outermost orbital

• Electron transfer creates charge imbalance

• Charge imbalance

creates ionic bond


sodiumion (Na+)

sodium chloride (NaCl)

sodium atom (Na) chlorine atom (Cl)

Na Cl

+

Na

–

Cl

chlorideion (Cl+)

Figure 2.7a

23

b.

Na+ Cl–


b (salting food): © PM Images/Getty RF; b (crystals): © Evelyn Jo Johnson

24

Covalent Bonding

• A covalent bond results when atoms share

electrons in such a way that each atom has an octet of electrons in the outer orbital.– An atom may share electrons with one or more atoms

• After sharing electrons, each atom has a

completed outer orbital.

• For example, two hydrogen atoms can share their single electron.

9

25

Covalent Bonding

• A single covalent bond results from sharing

one pair of electrons


H H H H

MolecularFormula

StructuralFormula

Electron Model

a. Hydrogen gas

H2

Figure 2.7a

26

Covalent Bonding

• A double covalent bond results from

sharing two pairs of electrons

O OO O

b. Oxygen gas

O2


Figure 2.7b

27

Shape of Molecules

• Structural formulas make it seem as if

molecules are one-dimensional.

• Molecules have a three-dimensional shape

that determines their biological function.

• Molecules with two atoms are linear.

• Molecules such as methane, with five atoms, have a tetrahedral shape.

10

28

Shape of MoleculesCopyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

c. Methane

HH

H

H

C

H

C H

H

H CH4

Figure 2.7c

29

Shape of Molecules

Space-filling ModelBall-and-stick Model

H

H

H

H109°

hydrogen

carbon

covalent bond

d. Methane–continued


Figure 2.7d

30

Nonpolar and Polar Covalent

Bonds

• If the sharing of electrons between two

atoms is fairly equal, a nonpolar covalent bond results.

• As in water, the sharing of electrons between

oxygen and each hydrogen is unequal, resulting in polar covalent bonds.

• Electronegativity is the attraction of an

atom for electrons in a covalent bond.

11

31

Nonpolar Covalent Bonds

• The sharing

between two

atoms that is

mostly equal


c. Methane

HH

H

H

C H C H

H

H2H H H H

H

MolecularFormula

StructuralFormula

Electron Model

a. Hydrogen gas

CH4

Figure 2.7

32


Space-filling ModelBall-and-stick Model

c. Methane

H

H

H

H109°

HH

H

H

C H C H

H

O OO O O2

H2H H H H

H

Molecular

Formula

Structural

FormulaElectron Model

a. Hydrogengas

b. Oxygen gas

CH4

hydrogen

carbon

covalent bond

d. Methane–continuedFigure 2.7

33

Polar Covalent Bonds

• The sharing between two atoms is unequal,

the covalent bond is described as polar

• Oxygen is more electronegative than hydrogen


δδδδ+ δδδδ+

δδδδ–

HH H

Hydrogens are partially positive.

O

Ball-and-stick ModelElectron Model Space-filling Model

H

O

O

H H

Oxygen attracts the shared

electrons and is partially negative.

a. W ater (H2O)

104.5°

12

34

Hydrogen Bonding

• Polarity within a water molecule causes the

hydrogen atoms in one molecule to be attracted to the oxygen atoms in other water

molecules.

• The attraction between partially (-) oxygen

and partially (+) hydrogen results in a

hydrogen bond.

• Bond is weak individually but strong collectively.

35

Hydrogen BondingCopyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

δδδδ+

δδδδ+

δδδδ–

H

HO

b. Hydrogen bonding between water molecules

hydrogenbond

Figure 2.9b

36


δδδδ+ δδδδ+

δδδδ–

δδδδ+

δδδδ+

δδδδ–

HH H

Hydrogens are partially positive.

O

Ball-and-stick ModelElectron Model Space-filling Model

H

O

HH

O

O

H H

b. Hydrogen bonding between water molecules

Oxygen attracts the shared

electrons and is partially negative.

hydrogen

bond

a. W ater (H2O)

104.5°

Figure 2.9

13

37

2.3 Chemistry of Water

• The first cell(s) evolved in water.

• Organisms are composed of 70–90% water.

• Water is a polar molecule.

• Water molecules form hydrogen bonds

which cause them to cling to one another.

– Water is liquid at temperatures typical of the Earth’s surface due to hydrogen bonding.

38

Properties of Water

• Water has a high heat capacity

– A calorie is the amount of heat energy needed

to raise the temperature of 1 g of water 1°C.

– The hydrogen bonds that link water molecules help water absorb heat without a great change in

temperature.

– Because the temperature of water rises and falls slowly, organisms are better able to maintain their normal internal temperatures.

39

Properties of Water

• Water has a high heat of vaporization

– Converting 1 g of the hottest water to a gas requires an input of 540 calories of heat energy.

– Gives animals in a hot environment an efficient way to release excess body heat.

– Also helps moderate temperatures along coasts.

14

40

Properties of Water

• Water is a solvent

– Due to its polarity, water facilitates chemical

reactions, both outside and within living systems.

– It dissolves many chemical substances.

– A solution contains dissolved substances, which

are then called solutes.

– Hydrophilic molecules attract water.

– Hydrophobic molecules do not attract water.

41


H

H

H H HH H

H H H

H

δδδδ+

δδδδ+

δδδδ+

δδδδ+

H HO

O

O OO O

δδδδ–

δδδδ–

Cl–Na+

42

Properties of Water

• Water molecules are cohesive and adhesive

– Water molecules cling together because of hydrogen bonding (cohesion).

– Water’s positive and negative poles allow it to adhere to polar surfaces (adhesion).

– Water is an excellent transport system, both outside and within living organisms.

• For example, blood transports dissolved and suspended substances throughout the body.

15

43

Properties of Water

• Water has a high surface tension

– The stronger the force between molecules in a liquid, the greater the surface tension.

– This allows some insects to walk on the surface of a pond or lake.

44

Properties of Water

• Frozen water (ice) is less dense than liquid water

– As liquid water cools, the molecules come closer together (densest at 4°C).

– Water expands as it freezes because a crystal lattice forms with hydrogen bonds farther apart.

– Ice floats on liquid water because it is less dense.

– Bodies of water freeze from the top down.

45


liquid water

ice lattice

Figure 2.11

16

46

Acids and Bases

• When water ionizes, it releases an equal number of hydrogen ions (H+) and hydroxide

ions (OH-)


H+ +OH H

water hydrogen

ion

hydroxide

ion

OH–

47

Acids and Bases

• Acidic Solutions (High H+ Concentrations)

– Acids are substances that that release hydrogen ions (H+) when dissociated in water.

• An example:

HCl H+ + Cl-

48

Acids and Bases

• Basic Solutions (Low H+ Concentrations)

– Bases are substances that dissociate in water, release hydroxide ions (OH-) or take up

hydrogen ions (H+)

• An example:

NaOH Na+ + OH-

17

49

Acids and Bases

• The pH scale indicates the acidity or

alkalinity of a solution.

– Scale ranges from 0 – 14.

– A pH below 7 is acidic. [H+] > [OH-]

– A pH above 7 is alkaline. [OH-] > [H+]

– A pH of 7 is neutral. [H+] = [OH-]

50


pu

re w

ate

r, tea

rs

51

Acids and Bases

• Buffers and pH

– A buffer is a chemical or combination of

chemicals that keep pH within normal limits.

– Bicarbonate ions (HCO3-) and carbonic acid

(H2CO3) found in human blood buffers the pH to

7.4.

18

52

Buffers and pH

If hydrogen ions (H+) are added to blood, this reaction occurs:

H+ + HCO3

- H2CO

3

If hydroxide ions (OH-) are added to blood, this reaction occurs:

OH- + H2CO

3HCO

3- + H

2O

These reactions prevent any significant change in blood pH.

53

2.4 Organic Molecules

• Organic molecules always include:

– carbon (C) and hydrogen (H)

– Those with only (H) and (C) are called hydrocarbons

H C C C C C C C C H

H

H

H

H

H

H

H

H

H

H

H

H

H

H

H

H


54


• The chemistry of carbon accounts for the formation of great variety of organic molecules.

• Carbon atoms contain four valence electrons.

• A carbon atom may share electrons with another carbon atom or other atoms in order to achieve eight electrons.

• Satisfying the octet rule

19

55


• Functional groups are a specific

combination of bonded atoms that always react in the same way.

• The more common functional groups are listed in Table 2.1.

56

57


• Macromolecules contain many molecules

joined together.

– Monomers: Simple organic molecules that exist individually

– Polymers: Large organic molecules form by

combining monomers

20

58


• Polymers in cells and their monomers

Polymer Monomer

carbohydrate (e.g., starch) monosaccharide

protein amino acid

nucleic acid nucleotide


59


• Cells use common reactions to join monomers.

– In a dehydration reaction an -OH and -H are removed as a water molecule.

– In a hydrolysis reaction, components of water

are added.

60

monomer monomer

monomer monomer

monomer monomer

OH H

OH H

b.

a.

monomer monomer

dehydrationreaction

hydrolysisreaction

H2O

H2O


Figure 2.13

21

61

2.5 Carbohydrates

• Carbohydrates function for quick fuel and

short-term energy storage in organisms.

– Play a structural role in woody plants, bacteria and insects

– On cell surfaces, involved in cell-to-cell recognition

62

Simple Carbohydrates

• Monosaccharides are sugars with 3 - 7

carbon atoms.

• Pentose refers to a 5-carbon sugar

• Hexose refers to a 6-carbon sugarCopyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

OH

OH

H

H

H

H

HO OH

HO

C

C

C

CC

4

5

6

3 2

1

OH

OH

H

H

H

HO OH

HO O

CH2OHCH2OH

C6H12O6Figure 2.14

63

Simple Carbohydrates

• Disaccharides contain two monosaccharides

joined by the dehydration reaction.

– Examples – maltose, sucrose, lactose


OH HO

HH

O+ +

O O O O

+ +

CH2OH

glucose C6H12O6 glucose C6H12O6

monosaccharide monosaccharide

CH2OH

dehydration reaction

hydrolysis reaction

CH2OH CH2OH

maltose C12H22O11

disaccharide

H2O

water

water

Figure 2.15

22

64

Polysaccharides

• Polysaccharides such as starch,

glycogen, and cellulose are long polymers that contain many glucose subunits.

65

Starch and Glycogen

• Starch is the storage form of glucose in plants.

– May contain up to 4,000 glucose units

– Fewer side branches than glycogen

• Glycogen is the storage form of glucose in animals.

– Liver stores glucose as glycogen

– In between meals, the liver releases glucose

stored in glycogen

66

O

O O

HHH

H

OH H

OH

O

O

HHH

H

OH

OH

O

O

HHH

H

OH

OH

H H

O

O

HHH

H

OH

OH

H

starchgranule

cell wall

potato cells

nonbranched

branched

CH2OH CH2OH CH2OH CH2OH


© Jeremy Burgess/SPL/Photo Researchers, Inc.Figure 2.16

23

67

O

O O

HH

H

H

OH H

OH

O

O

HH

H

H

OH

OH

O

O

HH

H

H

OH

OH

H H

O

O

HH

H

H

OH

OH

H

glycogen

granule

liver cells

CH2OH CH2OH CH2OH CH2OH


© Don W. Fawcett/Photo Researchers, Inc.Figure 2.17

68

Cellulose

• Some polysaccharides function as structural components of cells.

• Cellulose is found in the cell walls of plants.

– Accounts for the strong nature of the cell walls

– Has different chemical linkage than starch or glycogen

• Prevents us from digesting foods with cellulose

• Chiton, found in the exoskeleton of crabs, is another structural polysaccharide.

69


O

OH

OH HH

H

H

H

H

OHOH

H

H

H

H H H

OH

O

H H H

OH

O

OH

O

O

H

OH H

H OH

O

O O

O

OH

OH HH

H

H

H

H

OHOH

H

H

H

H H H

OH

O

H H H

OH

O

OH

O

O

H

OH H

H OH

O

O O

O

OH

OH HH

H

H

H

H

OHOH

H

H

H

H H H

OH

O

H H H

OH

O

OH

O

O

H

OH H

H OH

O

O O

glucosemolecules

microfibrils

cellulose fibers

CH2OH CH2OH

CH2OH

CH2OH

CH2OH

CH2OH

CH2OH

CH2OH

CH2OH

CH2OH

CH2OH

CH2OH

cellulosefiber

plantcell wall

© Science Source/J.D. Litvay/Visuals UnlimitedFigure 2.18

24

70

2.6 Lipids

• Lipids contain more energy per gram

than other biological molecules.

• Types

– Fats and oils used for energy storage

– Phospholipds from membranes

– Steroids include sex hormones

71

2.6 Lipids

• Lipids are diverse in structure and function.

• Lipids have one common characteristic –they do not dissolve in water (hydrophobic).

72

Fats and Oils

• Fats– Usually of animal origin

– Solid at room temperature

– Store energy, insulate against heat loss, form protective cushion

• Oils– Usually of plant origin

– Liquid at room temperature

25

73

Fats and Oils

• A fat molecule is also known as a triglyceride

or neutral fat.

• A triglyceride consists of

– One glycerol backbone

– Three fatty acids

74


+

C OHH

H

C OHH

C

C

OHH

H

H

H

C OH

H

C OH

C OH

H

C

H

C

H

C

H

H

CO

C

H

H

H C

HO

H

C

H

C

H

C

H

H

C

C

C

H

H

H

O

C

H

H H H

C

H

C

H

C

H

H

CO

HOC

H

H

C

H

H

H C

H

H H H

C

H

C

H

C

H

H

C

O

C

H

H

C

H

H

H

C

H

H H H

C

H

C

H

C

H

H

CO

H C

H

H H H

C

H

C

H

C

H

H

H

HO

HO

+ 3 H2O

glycerol 3 fatty acids fat molecule

dehydration reaction

hydrolysis reaction

3 watermolecules

Fats and Oils

Figure 2.19

75

• Emulsification

– Fat droplets disperse in water.

– Emulsifiers contain molecules with a polar

and nonpolar end.

26

76

Saturated, Unsaturated, and

Trans-Fatty Acids

• A fatty acid is a hydrocarbon chain that

ends with the acidic group —COOH.

• Saturated fatty acids have no double

bonds between carbon atoms.

• Unsaturated fatty acids have one or more

double bonds between carbon atoms.

77

Comparison of Saturated,

Unsaturated and Trans-Fats


C

H

C

H

C

H

C

H

C

H

C

H

H H

Saturated

(butter)

Unsaturated cis fats

(oils)

Unsaturated trans-fats

(hydrogenated oils)

78

Phospholipids

• Phospholipids are comprised of two fatty acids and a phosphate group

• The phosphate group is polar so the molecules are not electrically neutral.

• The phosphate group forms a polar head (hydrophilic) while the rest of the molecule is a nonpolar (hydrophobic) tail.

27

79

PhsopholipidsCopyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

1CH

2 –O

–O

R–

O–

P–

O–

3CH

2

2CH–

O

O

O O CC

CH2

CH2

CH2

CH2

CH2

CH2

CH2

CH

CH

CH2

CH2

CH2

CH2

CH2

CH2

CH2

CH2

CH2

CH2

CH2

CH2

CH2

CH2

CH

CH

CH2

CH2

CH2

CH2

CH2

CH2

CH2

CH2

CH2

CH3

CH3

inside cell

outside cell

Fa

tty a

cid

s

a. Plasma membrane of a cell

Nonpolar Tails

Polar Head

glycerol

phosphate

b. Phospholipid structure

• Spontaneously form a bilayer in which the

hydrophilic heads face outward toward watery solutions and the tails

form the hydrophobic interior

Figure 2.21

80

Steroids

• Steroids have a backbone of four fused

carbon rings

– Examples: Cholesterol, Testosterone, Estrogen


HO

CH3

OH

O

CH3

CH3

OH

a. Testosterone b. EstrogenFigure 2.22

81

2.7 Proteins• Proteins are polymers composed of amino acid

monomers

• Amino acids

– Amino group (-NH2)

– Acidic group (-COOH)

– R group varies

H N C C

amino acid

OH

H H

R

O

N C C

amino acid

acidic groupamino group

O

H R

H

OH

H


Figure 2.24

28

82


H

CH3

CC

O

O–

H

CH

CH3

CC

O

O–

H

CH2

SH

CC

O

O–

H

CH2

CC

O

O–

H3N+

H3N+

H3N+

H3C

H3N+

Figure 2.23

83

2.7 Proteins

• Proteins perform many functions

– Structural proteins give support (keratin, collagen)

– Enzymes speed up chemical reactions

– Hormones are chemical messengers

– Actin and myosin move cells and muscles

– Some proteins transport molecules in blood

– Antibodies protect cells

– Channels allow substances to cross membranes

84

Peptides

• Peptides

– A polypeptide is a single chain of amino acids.

– A peptide bond joins two amino acids.

dehydration reactionH2O

water

H

H

R

H N C C N CC

H

H

RO

peptide bond

dipeptide

H N C C

amino acid

OH

H H

R

O

N C C

amino acid

acidic groupamino group

O

H R

H

OH

O

OH

H hydrolysis reaction


Figure 2.24

29

85

Levels of Protein Organization

• Proteins have up to 4 levels of structural

organization.

– Primary structure is the linear sequence of the amino acids.

– Secondary structure occurs when the protein

takes on a certain orientation in space

• Two types include

– Alpha helix

– Beta sheet

86


αααα (alpha) helix

COO–amino acidpeptide bond

hydrogen bond

C NCH

R

C

CH

R

C N

C

CH

R

C

N

C

CH

R

N

C

CH

R

N

CH

R

N

C

N

CH

R

CH

hydrogen bond

ββββ(beta) pleated sheet

Figure 2.25

87


– The tertiary structure is the final three-dimensional shape.

• Maintained by various types of bonding between R groups

• Covalent, ionic, hydrogen bonding, disulfide bonding

– Quaternary structure is found in proteins with

multiple polypeptide chains.

• Separate polypeptide chains are arranged to give this highest structure

30

88



disulfide bond


Figure 2.25

89



COO–amino acidpeptide bond

disulfide bond

hydrogen bond

CN

CH

R

C

CH

R

CN

C

CH

R

C

N

C

CH

R

N

C

CH

R

N

CH

R

N

C

N

CH

R

CH

hydrogen bond


Figure 2.25

90


• The final shape of a protein is very important

to its function.

• A protein is denatured when it loses

structure and function.

– Occurs when proteins are exposed to extreme

heat or pH

31

91

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2.8 Nucleic Acids

• The two types of nucleic acids are

– DNA (deoxyribonucleic acid)

• Stores genetic information in the cell and in the organism

• DNA replicates to transmit its information when a cell divides or organism reproduces

– RNA (ribonucleic acid)

93

Structure of DNA and RNA

• Both DNA and RNA are polymers of

nucleotides

– Every nucleotide is a molecular complex of

• Phosphate

• Pentose sugar (ribose or deoxyribose)

• Nitrogen-containing base

– DNA contains: Adenine (A), Thymine (T), Guanine (G) and Cytosine (C)

– In RNA, uracil (U) replaces thymine

32

94


O

4'

5'

3' 2'

1'

–O P O

O

O–

phosphate

nitrogen-containing

base

pentose sugar

Nucleotide structure

C

S

C

Figure 2.26

95


• The nucleotides form a linear molecule called a strand.

• DNA is a double helix of two strands.

• The two strands are held together by hydrogen

bonds.

• Rungs of the ladder are formed by complementary paired bases.

– Adenine (A) always pairs with thymine (T)

– Cytosine (C) always pairs with guanine (G)

96


A

A

TT G

G

C S

S

S

P

P

P

P

S

P

S

P

S

S

P

S

P

P

C

A

A

A

T

T

T

G

G C

C

S

S

P

a. b. c.

one nucleotide

a: © Radius Images/Alamy RF

33

97


• RNA is single-stranded.

– Several types are involved in carrying

information from DNA to make proteins.

• ATP (Adenosine Triphosphate)

– ATP is a high-energy molecule.

– ATP undergoes hydrolysis and energy is

released.

– ATP is the energy “currency” of the cell.

98

• Last two phosphate bonds are unstable

and easily broken.

• Hydrolization forms ADP (adenosine

diphosphate).

• ATP can be rebuilt.– Add P to ADP to make ATP


+ + energyPPP P P P

H2O

Figure 2.28

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