Human Anatomy and Physiology

Basic Chemistry Unit 2Remember Metabolism is just those chemical

reactions in our body…so we have to understand Basic Chemistry

Matter vs. Energy

Matter can store energy through atomic bonds

We understand that matter can release energy…and that moving matter really fast (energy) can create new matter (particle accelerator)

Matter…a practical working definition

Anything that has mass and occupies space

States of matter:1. Solid—definite shape and volume2. Liquid—definite volume, changeable shape3. Gas—changeable shape and volume

Energy can be converted from one form to another, but always loses heat between conversions

Potential

Kinetic

Chemical

Electrical

Mechanical

Radiant or Electromagnetic

Composition of Matter Elements

◦ Cannot be broken down by ordinary chemical means ◦ Each has unique properties:

Physical properties Are detectable with our senses, or are

measurable Chemical properties

How atoms interact (bond) with each other Atoms

◦ Unique building blocks for each element◦ Atomic symbol: one- or two-letter chemical shorthand

for each element Au Fe Mg K Cl F Sn Cu

Elements of the Human Body Oxygen (O) Carbon (C) Hydrogen (H) Nitrogen (N)

Minor elements: About 3.9% of body mass◦ Calcium (Ca), phosphorus (P), potassium (K), sulfur

(S), sodium (Na), chlorine (Cl), magnesium (Mg), iodine (I), and iron (Fe)

Trace elements: < 0.01% of body mass◦ Part of enzymes, e.g., chromium (Cr), manganese

(Mn), and zinc (Zn)

About 96% of body mass

These are major elements

Your guide to atoms. Make sure you can read it!

So the basic unit of matter is…the ATOM

Be able to describe and draw an atom…including:

Protons Neutrons Electrons Ions Isotopes

Atomic Structure Neutrons

No charge Mass = 1 atomic mass unit (amu)

Protons Positive charge Mass = 1 amu

Electrons◦ Orbit nucleus◦ Equal in number to protons in atom◦ Negative charge ◦ 1/2000 the mass of a proton (0 amu)

Figure 2.1

(a) Planetary model (b) Orbital model

Helium atom

2 protons (p+)2 neutrons (n0)2 electrons (e–)

Helium atom

2 protons (p+)2 neutrons (n0)2 electrons (e–)

Nucleus Nucleus

Proton Neutron Electroncloud

Electron

Figure 2.2

ProtonNeutronElectron

Helium (He)(2p+; 2n0; 2e–)

Lithium (Li)(3p+; 4n0; 3e–)

Hydrogen (H)(1p+; 0n0; 1e–)

Identifying Elements Atomic number = number of protons in

nucleus Mass number = mass of the protons and

neutrons◦ Mass numbers of atoms of an element are not all

identical◦ Isotopes are structural variations of elements that

differ in the number of neutrons they contain Atomic weight = average of mass numbers

of all isotopes

Figure 2.3

ProtonNeutronElectron

Deuterium (2H)(1p+; 1n0; 1e–)

Tritium (3H)(1p+; 2n0; 1e–)

Hydrogen (1H)(1p+; 0n0; 1e–)

Radioisotopes Spontaneous decay (radioactivity) Similar chemistry to stable isotopes Can be detected with scanners Valuable tools for biological research and

medicine Cause damage to living tissue:

◦ Useful against localized cancers◦ Radon from uranium decay causes lung cancer

Ionic Bonds Ions are formed by transfer of valence shell

electrons between atoms◦ Anions (– charge) have gained one or more

electrons◦ Cations (+ charge) have lost one or more

electrons Attraction of opposite charges results in an

ionic bond

Molecules and Compounds Most atoms combine chemically with other

atoms to form molecules and compounds◦ Molecule—two or more atoms bonded together

(e.g., H2 or C6H12O6) Compound—two or more different kinds of atoms

bonded together (e.g., C6H12O6) Elements- one or more of the same atom bonded

together

Chemical Bonds Electrons occupy up to seven electron shells

(energy levels) around nucleus Octet rule: Except for the first shell which is

full with two electrons, atoms interact in a manner to have eight electrons in their outermost energy level (valence shell)

Figure 2.5a

Helium (He)(2p+; 2n0; 2e–)

Neon (Ne)(10p+; 10n0; 10e–)

2e 2e8e

(a) Chemically inert elementsOutermost energy level (valence shell) complete

Figure 2.5b

2e4e

2e8e

1e

(b) Chemically reactive elementsOutermost energy level (valence shell) incomplete

Hydrogen (H)(1p+; 0n0; 1e–)

Carbon (C)(6p+; 6n0; 6e–)

1e

Oxygen (O)(8p+; 8n0; 8e–) Sodium (Na)

(11p+; 12n0; 11e–)

2e6e

Be able to draw and predict bonding for the following atoms:◦ Carbon◦ Hydrogen◦ Calcium◦ Oxygen◦ Fluorine◦ Lithium◦ Sodium◦ Chlorine◦ Helium

Valence Electrons and Bonding

Ionic

Covalent

Hydrogen

Three important Bonds

Figure 2.6a-b

Sodium atom (Na)(11p+; 12n0; 11e–)

Chlorine atom (Cl)(17p+; 18n0; 17e–)

Sodium ion (Na+) Chloride ion (Cl–)

Sodium chloride (NaCl)

+ –

(a) Sodium gains stability by losing one electron, and chlorine becomes stable by gaining one electron.

(b) After electron transfer, the oppositely charged ions formed attract each other.

Formation of an Ionic Bond Ionic compounds form crystals instead of

individual molecules◦ NaCl (sodium chloride)

Figure 2.6c

CI–

Na+

(c) Large numbers of Na+ and Cl– ions associate to form salt (NaCl) crystals.

Covalent Bonds Formed by sharing of two or more valence

shell electrons Allows each atom to fill its valence shell at

least part of the time

Figure 2.7a

+

Hydrogenatoms

Carbonatom

Molecule ofmethane gas (CH4)

Structuralformulashows singlebonds.

(a) Formation of four single covalent bonds: carbon shares four electron pairs with four hydrogen atoms.

or

Resulting moleculesReacting atoms

Figure 2.7b

or

Oxygenatom

Oxygenatom

Molecule ofoxygen gas (O2)

Structuralformulashowsdouble bond.(b) Formation of a double covalent bond: Two

oxygen atoms share two electron pairs.

Resulting moleculesReacting atoms

+

Figure 2.7c

+ or

Nitrogenatom

Nitrogenatom

Molecule ofnitrogen gas (N2)

Structuralformulashowstriple bond.(c) Formation of a triple covalent bond: Two

nitrogen atoms share three electron pairs.

Resulting moleculesReacting atoms

Covalent Bonds Sharing of electrons may be equal or

unequal◦ Equal sharing produces electrically balanced

nonpolar molecules CO2

Figure 2.8a

Covalent Bonds Unequal sharing by atoms with different

electron-attracting abilities produces polar molecules◦ H2O

Atoms with six or seven valence shell electrons are electronegative, e.g., oxygen

Atoms with one or two valence shell electrons are electropositive, e.g., sodium

Figure 2.8b

Figure 2.9

Hydrogen Bonds Attractive force between electropositive

hydrogen of one molecule and an electronegative atom of another molecule◦ Common between dipoles such as water◦ Also act as intramolecular bonds, holding a large

molecule in a three-dimensional shape◦ http://www.youtube.com/watch?v=lkl5cbfqFRM

(a) The slightly positive ends (+) of the watermolecules become aligned with the slightlynegative ends (–) of other water molecules.

+

–

–

–– –

+

+

+

+

+

Hydrogen bond(indicated bydotted line)

Figure 2.10a

Figure 2.10b

(b) A water strider can walk on a pond because of the highsurface tension of water, a result of the combinedstrength of its hydrogen bonds.

Chemical vs Physical reactions

Chemical reactions Physical reactions

Chemical bonds are formed, rearranged or broken

Synthesis, Decomposition, Exchange, Redox

Chemical Equations◦ Reactants◦ Products◦ Molecular formulas ◦ Balance atoms

No chemical bonds are changed

Most matter exists as MIXTURES◦ Solution

Aqueous/ homeogeneous Solute/ solvent http://www.youtube.com/watch?

v=3G472AA3SEs◦ Colloid

Medium solute/ gel/ foam heterogeneous

◦ Suspension Large solute/ heterogeneous

Concentration of Solutions Expressed as

◦ Percent, or parts per 100 parts◦ Milligrams per deciliter (mg/dl)◦ Ppm: parts per million◦ Molarity, or moles per liter (M)

1 mole = the atomic weight of an element or molecular weight (sum of atomic weights) of a compound in grams

1 mole of any substance contains 6.02 1023 molecules (Avogadro’s number) What is the molecular weight of a gram of glucose

(C6H12O6)? To make a 1M (one molar) solution of glucose, how much

glucose would you weigh to mix with water to make one liter?

Figure 2.4

Solution

Soluteparticles

Soluteparticles

Soluteparticles

Solute particles are verytiny, do not settle out orscatter light.

ColloidSolute particles are largerthan in a solution and scatterlight; do not settle out.

SuspensionSolute particles are verylarge, settle out, and mayscatter light.

ExampleMineral water

ExampleGelatin

ExampleBlood

Chemical Reactions can be expressed as EquationsH + H H2 (hydrogen gas)

4H + C CH4 (methane)

Bonds are always broken, made, rearrangedAtoms are always accounted for so formulas must be

balanced to work Chemical equilibrium occurs if neither a forward nor

reverse reaction is dominant Many biological reactions are essentially irreversible

due to Energy requirements and/or◦ Removal of products

(reactants) (product)

Synthesis◦ Anabolic = forms bonds/ Endergonic (absorbs energy)◦ A + B → AB Found in rapidly growing tissues

Decomposition◦ Catabolic = breaks bonds/ Exergonic (releases energy)◦ AB → A + B Found in metabolic reactions/ cell respiration

Exchange or Displacement◦ Involves synthesis and decomposition/ atoms rearranged◦ AB + C → AC + B AB + CD → AD + CB

Oxidation-Reduction (Redox) reactions◦ Both decomposition and exchange reactions◦ Electrons move between reactants◦ Reactant that loses electron is an electron donor/ oxidized◦ Reactant that gains electron is an electron acceptor/ reduced

LEO the lion goes GER

Patterns of Chemical Reactions

Rate of Chemical Reactions Rate of reaction is influenced by:

◦ temperature rate◦ particle size rate ◦ concentration of reactant rate

Catalysts: rate without being chemically changed◦ Enzymes are biological catalysts

Figure 2.11a

ExampleAmino acids are joined together toform a protein molecule.

(a) Synthesis reactionsSmaller particles are bondedtogether to form larger,more complex molecules.

Amino acidmolecules

Proteinmolecule

Figure 2.11b

ExampleGlycogen is broken down to releaseglucose units.

Bonds are broken in largermolecules, resulting in smaller,less complex molecules.

(b) Decomposition reactions

Glucosemolecules

Glycogen

Figure 2.11c

ExampleATP transfers its terminal phosphategroup to glucose to form glucose-phosphate.

Bonds are both made and broken(also called displacement reactions).

(c) Exchange reactions

Glucose Adenosine triphosphate (ATP)

Adenosine diphosphate (ADP)Glucosephosphate

+

+

Oxidation Reduction (Redox)

LEO the lion goes GER

Common Example of Redox Reactions in Body:

Aerobic Cell Respiration

Glucose is oxidized to carbon dioxide as it loses hydrogen atoms oxygen is reduced to water as it accepts the hydrogen atoms

Classes of Compounds Inorganic compounds

Water, salts, and many acids and bases Do not have to contain carbon

Organic compounds Carbohydrates, fats, proteins, and nucleic acids Must contain carbon, usually large, and are

covalently bonded

Water 60%–80% of the volume of living cells Most important inorganic compound in

living organisms because of its properties

Properties of Water High heat capacity

◦ Absorbs and releases heat with little temperature change

◦ Prevents sudden changes in temperature High heat of vaporization

◦ Evaporation requires large amounts of heat◦ Useful cooling mechanism

Polar solvent properties◦ Dissolves and dissociates ionic substances◦ Forms hydration layers around large charged

molecules, e.g., proteins (colloid formation)◦ Body’s major transport medium

Figure 2.12

Water molecule

Ions in solutionSalt crystal

–

+

+

Properties of Water Reactivity

◦ A necessary part of hydrolysis and dehydration synthesis reactions

Cushioning◦ Protects certain organs from physical trauma,

e.g., cerebrospinal fluid

Figure 2.14

+

Glucose Fructose

Water isreleased

Monomers linked by covalent bond

Monomers linked by covalent bond

Water isconsumed

Sucrose

(a) Dehydration synthesisMonomers are joined by removal of OH from one monomer

and removal of H from the other at the site of bond formation.

+

(b) HydrolysisMonomers are released by the addition of a water molecule, adding OH to one monomer and H to the other.

(c) Example reactionsDehydration synthesis of sucrose and its breakdown by hydrolysis

Monomer 1 Monomer 2

Monomer 1 Monomer 2

+

Salts Ionic compounds that dissociate in water Contain cations other than H+ and anions

other than OH–

Ions (electrolytes) conduct electrical currents in solution

Ions play specialized roles in body functions (e.g., sodium, potassium, calcium, and iron)

pH…its all about the ions

Acids and Bases Both are electrolytes Acids are proton (hydrogen ion) donors

(release H+ in solution) HCl H+ + Cl–

Bases are proton acceptors (take up H+ from solution)◦ NaOH Na+ + OH–

OH– accepts an available proton (H+) OH– + H+ H2O

Bicarbonate ion (HCO3–) and ammonia (NH3)

are important bases in the body

pH: Acid-Base Concentration pH = the negative logarithm of [H+] in

moles per liter Neutral solutions:

◦ Pure water is pH neutral (contains equal numbers of H+ and OH–)

◦ pH of pure water = pH 7: [H+] = 10 –7 M◦ All neutral solutions are pH 7

pH: Acid-Base Concentration Acidic solutions

◦ [H+], pH ◦ Acidic pH: 0–6.99◦ pH scale is logarithmic: a pH 5 solution has 10

times more H+ than a pH 6 solution Alkaline solutions

◦ [H+], pH◦ Alkaline (basic) pH: 7.01–14

Common Substances and their pH value

Figure 2.13

Concentration(moles/liter)

[OH–]100 10–14

10–1 10–13

10–2 10–12

10–3 10–11

10–4 10–10

10–5 10–9

10–6 10–8

10–7 10–7

10–8 10–6

10–9 10–5

10–10 10–4

10–11 10–3

10–12 10–2

10–13 10–1

[H+] pH Examples1M Sodiumhydroxide (pH=14)

Oven cleaner, lye(pH=13.5)

Household ammonia(pH=10.5–11.5)

Neutral

Household bleach(pH=9.5)

Egg white (pH=8)

Blood (pH=7.4)

Milk (pH=6.3–6.6)

Black coffee (pH=5)

Wine (pH=2.5–3.5)

Lemon juice; gastricjuice (pH=2)

1M Hydrochloricacid (pH=0)10–14 100

14

13

12

11

10

9

8

7

6

5

4

3

2

1

0

Acid-Base Homeostasis pH change interferes with cell function and

may damage living tissue Slight change in pH can be fatal pH is regulated by kidneys, lungs, and

buffers

Buffers Mixture of compounds that resist pH

changes Convert strong (completely dissociated)

acids or bases into weak (slightly dissociated) ones◦ Carbonic acid-bicarbonate system

Carbon base◦ Electroneutral! …makes

carbon a small sharer

Functional groups

Monomers and polymers

Organic Compounds

Carbohydrates

Monosaccharides: CH2O

Disaccharides

FunctionIs Fuel, Easy, AvailableFuel for Our cells

Polysaccharides

Carbohydrates Sugars and starches Contain C, H, and O [(CH20)n] Three classes

◦ Monosaccharides◦ Disaccharides◦ Polysaccharides

Functions◦ Major source of cellular fuel (e.g., glucose)◦ Oxidation reduction reactions◦ Structural molecules (e.g., ribose sugar in RNA or

glycocalyx)

Figure 2.15a

ExampleHexose sugars (the hexoses shown here are isomers)

ExamplePentose sugars

Glucose Fructose Galactose Deoxyribose Ribose

(a) MonosaccharidesMonomers of carbohydrates

Simple sugars containing three to seven C atoms (CH20)n

Figure 2.15b

ExampleSucrose, maltose, and lactose(these disaccharides are isomers)

Glucose Fructose Glucose Glucose GlucoseSucrose Maltose Lactose

Galactose

(b) DisaccharidesConsist of two linked monosaccharides

Double sugarsToo large to pass through cell membranes

Figure 2.15c

ExampleThis polysaccharide is a simplified representation of glycogen, a polysaccharide formed from glucose units.

(c) PolysaccharidesLong branching chains (polymers) of linked monosaccharides

Glycogen

Polymers of simple sugars, e.g., starch and glycogen/ Not very soluble

Lipids: energy storage, structure, insulation

Contain C, H, O (less than in carbohydrates), and sometimes P, Insoluble in water

Triglycerides Steroids

Eicosanoids Lipoproteins

◦ HDL and LDL

Triglycerides Neutral fats—solid fats and liquid oils Composed of three fatty acids bonded to a

glycerol molecule Main functions

◦ Energy storage◦ Insulation◦ Protection

Figure 2.16a

Glycerol

+

3 fatty acid chains Triglyceride,or neutral fat

3 watermolecules

(a) Triglyceride formation Three fatty acid chains are bound to glycerol by

dehydration synthesis

Saturation of Fatty Acids Saturated fatty acids

◦Single bonds between C atoms; maximum number of H

◦Solid animal fats, e.g., butter Unsaturated fatty acids

◦One or more double bonds between C atoms

◦Reduced number of H atoms ◦Plant oils, e.g., olive oil Trans Fats Omega 3 Fatty acids

Omega 6

Figure 2.16b

Phosphorus-containing

group (polar“head”)

ExamplePhosphatidylcholine

Glycerolbackbone

2 fatty acid chains(nonpolar “tail”)

Polar“head”

Nonpolar“tail”

(schematicphospholipid)

(b) “Typical” structure of a phospholipid molecule Two fatty acid chains and a phosphorus-containing group are

attached to the glycerol backbone.

Phospholipids Modified triglycerides:

◦ Glycerol + two fatty acids and a phosphorus (P)-containing group

“Head” and “tail” regions have different properties

Important in cell membrane structure

Steroids Steroids—basically flat, interlocking four-ring

structure Cholesterol:

◦ vitamin D, steroid hormones, and bile salts◦ Steroid hormones include estrogen, progesterone,

testosterone, adrenocortical hormones (cortisol and aldosterone)

◦ Role in membrane: Limits movement (fluidity) of fatty acid chains Stabilizes membrane Lowers freezing point so allows functional membrane in

colder temperatures

Figure 2.16c

ExampleCholesterol (cholesterol is thebasis for all steroids formed in the body)

(c) Simplified structure of a steroid

Four interlocking hydrocarbon rings form a steroid.

Other Lipids in the Body Other fat-soluble vitamins

◦ Vitamins A, E, and K Lipoproteins

◦ Transport fats in the blood◦ HDL: carries cholesterol from cells to liver◦ LDL: carries cholesterol to cells

Eicosanoids◦ Many different types such as prostaglandins and

leukotrienes◦ Derived from a fatty acid (arachidonic acid) in cell

membranes◦ Acts as paracrines

Protein

Proteome makes up 10-30% of cell mass. While proteins make the basic structure of body, they also have many other vital roles of life.

Figure 2.17

(a) Generalized structure of all amino acids.

(b) Glycine is the simplest amino acid.

(c) Aspartic acid (an acidic amino acid) has an acid group (—COOH) in the R group.

(d) Lysine (a basic amino acid) has an amine group (–NH2) in the R group.

(e) Cysteine (a basic amino acid) has a sulfhydryl (–SH) group in the R group, which suggests that this amino acid is likely to participate in intramolecular bonding.

Aminegroup

Acidgroup

Polymers of amino acids (20 types)Joined by peptide bonds

Contain C, H, O, N, and sometimes S and P

Figure 2.18

Amino acid Amino acid Dipeptide

Dehydration synthesis:The acid group of one amino acid is bonded to the amine group of the next, with loss of a water molecule.

Hydrolysis: Peptide bonds linking amino acids together are broken when water is added to the bond.

+

Peptidebond

Figure 2.19a

(a) Primary structure: The sequence of amino acids forms the polypeptide chain.

Amino acid Amino acid Amino acid Amino acid Amino acid

Figure 2.19b

a-Helix: The primary chain is coiledto form a spiral structure, which isstabilized by hydrogen bonds.

b-Sheet: The primary chain “zig-zags” backand forth forming a “pleated” sheet. Adjacentstrands are held together by hydrogen bonds.

(b) Secondary structure:The primary chain forms spirals (a-helices) and sheets (b-sheets).

Figure 2.19c

Tertiary structure of prealbumin(transthyretin), a protein thattransports the thyroid hormonethyroxine in serum and cerebro-spinal fluid.

(c) Tertiary structure: Superimposed on secondary structure. a-Helices and/or b-sheets are folded up to form a compact globular molecule held together by intramolecular bonds.

Figure 2.19d

Quaternary structure ofa functional prealbuminmolecule. Two identicalprealbumin subunitsjoin head to tail to formthe dimer.

(d) Quaternary structure: Two or more polypeptide chains, each with its own tertiary structure, combine to form a functional protein.

Fibrous and Globular Proteins Fibrous (structural) proteins

◦ Strandlike, water insoluble, and stable ◦ Examples: keratin, elastin, collagen, and certain

contractile fibers Globular (functional) proteins

◦ Compact, spherical, water-soluble and sensitive to environmental changes

◦ Specific functional regions (active sites) ◦ Examples: antibodies, hormones, molecular

chaperones, and enzymes

Protein Denaturation Shape change and disruption of active sites

due to environmental changes (e.g., decreased pH or increased temperature)

Reversible in most cases, if normal conditions are restored

Irreversible if extreme changes damage the structure beyond repair (e.g., cooking an egg)

Molecular Chaperones (Chaperonins) Ensure quick and

accurate folding and association of proteins

Assist translocation of proteins and ions across membranes

Promote breakdown of damaged or denatured proteins

Help trigger the immune response

Produced in response to stressful stimuli, e.g., O2 deprivation

Enzymes

Enzymes are biological catalysts that do not initiate a reaction, but only speed up a reaction. Note the change in configuration as enzyme binds to substrate .

Enzymes speed up reactions by reducing required activation energy. How do enzymes do this?

Figure 2.20

Activationenergy required

Less activationenergy required

WITHOUT ENZYME WITH ENZYME

Reactants

Product Product

Reactants

Biological catalystsLower the activation energy, increase the speed of a reaction (millions of reactions per minute!)

Characteristics of Enzymes Often named for the reaction they catalyze;

usually end in -ase (e.g., hydrolases, oxidases)

Chain reactions/ inactive forms Either pure protein or functional enzymes

(holoenzymes)◦Apoenzyme (protein portion) ◦Cofactor (metal ion) or coenzyme (a

vitamin) Copper/ iron B Vitamins

Holoenzyme

Figure 2.21

Substrates (S)e.g., amino acids

Enzyme (E)Enzyme-substratecomplex (E-S)

Enzyme (E)

Product (P)e.g., dipeptide

Energy isabsorbed;bond isformed.

Water isreleased. Peptide

bond

Substrates bindat active site.Enzyme changesshape to holdsubstrates inproper position.

Internalrearrangementsleading tocatalysis occur.

Product isreleased. Enzymereturns to originalshape and isavailable to catalyzeanother reaction.

Active site

+ H2O

1 23

ProteosomesSuch a strange discovery: a protein that acts like an enzyme. Mad Cow disease is the result of a similar type molecule called a prion.

Nucleic Acids DNA and RNA

◦ Largest molecules in the body Contain C, O, H, N, and P Building block = nucleotide, composed of N-

containing base, a pentose sugar, and a phosphate group

Nucleic Acids

DNA (Deoxyribonucleic acid)

RNA(Ribonucleic Acid)

Found in nucleus Genome Replicates (interphase)

and transcribes Monomer of deoxyribose

and 4 complementary bases: adenine/ guanine/ cytosine/ thymine

Double helix Watson and Crick

Produced in nucleus but functions in cytoplasm

Transcription/ translation (protein synthesis)

Different types: mRNA, tRNA, rRNA

Single strand/ straight or folded

Ribose sugar and substitute uracil for thymine

Figure 2.22

DeoxyribosesugarPhosphate

Sugar-phosphatebackbone

Adenine nucleotideHydrogenbond

Thymine nucleotide

PhosphateSugar:

Deoxyribose PhosphateSugarThymine (T)Base:

Adenine (A)

Adenine (A)

Thymine (T)

Cytosine (C)

Guanine (G)

(b)

(a)

(c) Computer-generated image of a DNA molecule

Nucleotides: Monomers of Nucleic Acids

Figure 2.23

Adenosine triphosphate (ATP)

Adenosine diphosphate (ADP)

Adenosine monophosphate (AMP)

Adenosine

Adenine

Ribose

Phosphate groups

High-energy phosphatebonds can be hydrolyzedto release energy.

Adenine-containing RNA nucleotide with two additional phosphate groups

Adenosine Triphosphate

Function of ATP Phosphorylation:

◦ Terminal phosphates are enzymatically transferred to and energize other molecules

◦ Such “primed” molecules perform cellular work (life processes) using the phosphate bond energy

Figure 2.24

Solute

Membraneprotein

Relaxed smoothmuscle cell

Contracted smoothmuscle cell

+

+

+

Transport work: ATP phosphorylates transportproteins, activating them to transport solutes(ions, for example) across cell membranes.

Mechanical work: ATP phosphorylates contractile proteins in muscle cells so the cells can shorten.

Chemical work: ATP phosphorylates key reactants, providing energy to drive energy-absorbing chemical reactions.

(a)

(b)

(c)