chemistry comes alive: part a -...

PowerPoint® Lecture Slides prepared by Janice Meeking, Mount Royal College

C H A P T E R

Copyright © 2010 Pearson Education, Inc.

2 Chemistry Comes Alive: Part A


Matter

• Anything that has mass and occupies space

• States of matter:

1. Solid—definite shape and volume

2. Liquid—definite volume, changeable shape

3. Gas—changeable shape and volume


Energy

• Capacity to do work or put matter into motion

• Types of energy:

• Kinetic—energy in action

• Potential—stored (inactive) energy

PLAY Animation: Energy Concepts


Forms of Energy

• Chemical energy—stored in bonds of chemical substances

• Electrical energy—results from movement of charged particles

• Mechanical energy—directly involved in moving matter

• Radiant or electromagnetic energy—exhibits wavelike properties (i.e., visible light, ultraviolet light, and X-rays)


Energy Form Conversions

• Energy may be converted from one form to another

• Conversion is inefficient because some energy is “lost” as heat


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 one another


Composition of Matter

• Atoms

• Unique building blocks for each element

• Atomic symbol: one- or two-letter chemical shorthand for each element


Major Elements of the Human Body

• Oxygen (O)

• Carbon (C)

• Hydrogen (H)

• Nitrogen (N)

About 96% of body mass


Lesser Elements of the Human Body

• 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 of the Human Body

• < 0.01% of body mass:

• Part of enzymes, e.g., chromium (Cr), manganese (Mn), and zinc (Zn)


Atomic Structure

• Determined by numbers of subatomic particles

• Nucleus consists of neutrons and protons


Atomic Structure

• Neutrons

• No charge

• Mass = 1 atomic mass unit (amu)

• Protons

• Positive charge

• Mass = 1 amu


Atomic Structure

• Electrons

• Orbit nucleus

• Equal in number to protons in atom

• Negative charge

• 1/2000 the mass of a proton (0 amu)


Models of the Atom

• Orbital model: current model used by chemists

• Depicts probable regions of greatest electron density (an electron cloud)

• Useful for predicting chemical behavior of atoms


Models of the Atom

• Planetary model—oversimplified, outdated model

• Incorrectly depicts fixed circular electron paths

• Useful for illustrations (as in the text)

Copyright © 2010 Pearson Education, Inc. 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


Identifying Elements

• Atoms of different elements contain different numbers of subatomic particles

• Compare hydrogen, helium and lithium (next slide)


ProtonNeutronElectron

Helium (He)(2p+; 2n0; 2e–)

Lithium (Li)(3p+; 4n0; 3e–)

Hydrogen (H)(1p+; 0n0; 1e–)



• 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


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


Radioisotopes

• Valuable tools for biological research and medicine

• Cause damage to living tissue:

• Useful against localized cancers

• Radon from uranium decay causes lung cancer


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)


Mixtures

• Most matter exists as mixtures

• Two or more components physically intermixed

• Three types of mixtures

• Solutions

• Colloids

• Suspensions


Solutions

• Homogeneous mixtures

• Usually transparent, e.g., atmospheric air or seawater

• Solvent

• Present in greatest amount, usually a liquid

• Solute(s)

• Present in smaller amounts


Concentration of Solutions

• Expressed as• Percent, or parts per 100 parts

• Milligrams per deciliter (mg/dl)

• 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)


Colloids and Suspensions

• Colloids (emulsions)• Heterogeneous translucent mixtures, e.g.,

cytosol

• Large solute particles that do not settle out

• Undergo sol-gel transformations

• Suspensions:• Heterogeneous mixtures, e.g., blood

• Large visible solutes tend to settle out


Solution

Soluteparticles

Soluteparticles

Soluteparticles

Solute particles are verytiny, do not settle out or

scatter light.

ColloidSolute particles are larger

than in a solution and scatterlight; do not settle out.

SuspensionSolute particles are verylarge, settle out, and may

scatter light.

ExampleMineral water

ExampleGelatin

ExampleBlood


Mixtures vs. Compounds

• Mixtures• No chemical bonding between components

• Can be separated physically, such as by straining or filtering

• Heterogeneous or homogeneous

• Compounds• Can be separated only by breaking bonds

• All are homogeneous


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)


Chemically Inert Elements

• Stable and unreactive

• Outermost energy level fully occupied or contains eight electrons

Copyright © 2010 Pearson Education, Inc. Figure 2.5a

Helium (He)(2p+; 2n0; 2e–)

Neon (Ne)(10p+; 10n0; 10e–)

2e 2e8e

(a) Chemically inert elements

Outermost energy level (valence shell) complete


Chemically Reactive Elements

• Outermost energy level not fully occupied by electrons

• Tend to gain, lose, or share electrons (form bonds) with other atoms to achieve stability

Copyright © 2010 Pearson Education, Inc. 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


Types of Chemical Bonds

• Ionic

• Covalent

• Hydrogen


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

Copyright © 2010 Pearson Education, Inc. 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, andchlorine becomes stable by gaining one electron.

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


Formation of an Ionic Bond

• Ionic compounds form crystals instead of individual molecules

• NaCl (sodium chloride)

Copyright © 2010 Pearson Education, Inc. Figure 2.6c

CI–

Na+

(c) Large numbers of Na+ and Cl– ionsassociate 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


+

Hydrogenatoms

Carbonatom

Molecule ofmethane gas (CH4)

Structuralformulashows singlebonds.

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

or

Resulting moleculesReacting atoms


or

Oxygenatom

Oxygenatom

Molecule ofoxygen gas (O2)

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

oxygen atoms share two electron pairs.


+


+ or

Nitrogenatom

Nitrogenatom

Molecule ofnitrogen gas (N2)

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

nitrogen atoms share three electron pairs.



Covalent Bonds

• Sharing of electrons may be equal or unequal

• Equal sharing produces electrically balanced nonpolar molecules

• CO2


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


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

PLAY Animation: Hydrogen Bonds


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

δ+

δ–

δ–

δ–δ– δ–

δ+

δ+

δ+

δ+

δ+

Hydrogen bond(indicated bydotted line)

Figure 2.10a


(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 Reactions

• Occur when chemical bonds are formed, rearranged, or broken

• Represented as chemical equations

• Chemical equations contain:

• Molecular formula for each reactant and product

• Relative amounts of reactants and products, which should balance


Examples of Chemical Equations

H + H → H2 (hydrogen gas)

4H + C → CH4 (methane)

(reactants) (product)


Patterns of Chemical Reactions

• Synthesis (combination) reactions

• Decomposition reactions

• Exchange reactions


Synthesis Reactions

• A + B → AB

• Always involve bond formation

• Anabolic


ExampleAmino acids are joined together toform a protein molecule.

(a) Synthesis reactionsSmaller particles are bonded

together to form larger,more complex molecules.

Amino acidmolecules

Proteinmolecule


Decomposition Reactions

• AB → A + B

• Reverse synthesis reactions

• Involve breaking of bonds

• Catabolic


ExampleGlycogen is broken down to releaseglucose units.

Bonds are broken in largermolecules, resulting in smaller,

less complex molecules.

(b) Decomposition reactions

Glucosemolecules

Glycogen


Exchange Reactions

• AB + C → AC + B

• Also called displacement reactions

• Bonds are both made and broken


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) Reactions

• Decomposition reactions: Reactions in which fuel is broken down for energy

• Also called exchange reactions because electrons are exchanged or shared differently

• Electron donors lose electrons and are oxidized

• Electron acceptors receive electrons and become reduced


Chemical Reactions

• All chemical reactions are either exergonic or endergonic

• Exergonic reactions—release energy

• Catabolic reactions

• Endergonic reactions—products contain more potential energy than did reactants

• Anabolic reactions


Chemical Reactions

• All chemical reactions are theoretically reversible

• A + B → AB

• AB → A + B

• Chemical equilibrium occurs if neither a forward nor reverse reaction is dominant

• Many biological reactions are essentially irreversible due to

• Energy requirements

• Removal of products


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

PowerPoint® Lecture Slides prepared by Janice Meeking, Mount Royal College

C H A P T E R


2 Chemistry Comes Alive: Part B


Classes of Compounds

• Inorganic compounds

• Water, salts, and many acids and bases

• Do not contain carbon

• Organic compounds

• Carbohydrates, fats, proteins, and nucleic acids

• 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


Properties of Water

• 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


Water molecule

Ions in solutionSalt crystal

δ–

δ+

δ+


Properties of Water

• Reactivity

• A necessary part of hydrolysis and dehydration synthesis reactions

• Cushioning

• Protects certain organs from physical trauma, e.g., cerebrospinal fluid


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)


Acids and Bases

• Both are electrolytes

• Acids are proton (hydrogen ion) donors (release H+ in solution)

• HCl → H+ + Cl–


Acids and Bases

• 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


Acid-Base Concentration

• Acid solutions contain [H+]

• As [H+] increases, acidity increases

• Alkaline solutions contain bases (e.g., OH–)

• As [H+] decreases (or as [OH–] increases), alkalinity increases


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


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+] pHExamples

1M 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


Organic Compounds

• Contain carbon (except CO2 and CO, which are inorganic)

• Unique to living systems

• Include carbohydrates, lipids, proteins, and nucleic acids


Organic Compounds

• Many are polymers—chains of similar units (monomers or building blocks)

• Synthesized by dehydration synthesis

• Broken down by hydrolysis reactions


+

Glucose Fructose

Water isreleased

Monomers linked by covalent bond

Monomers linked by covalent bond

Water isconsumed

Sucrose

(a) Dehydration synthesis

Monomers are joined by removal of OH from one monomerand removal of H from the other at the site of bond formation.

+

(b) Hydrolysis

Monomers are released by the addition of a water molecule, adding OH to one monomer and H to the other.

(c) Example reactions

Dehydration synthesis of sucrose and its breakdown by hydrolysis

Monomer 1 Monomer 2

Monomer 1 Monomer 2

+


Carbohydrates

• Sugars and starches

• Contain C, H, and O [(CH20)n]

• Three classes

• Monosaccharides

• Disaccharides

• Polysaccharides


Carbohydrates

• Functions

• Major source of cellular fuel (e.g., glucose)

• Structural molecules (e.g., ribose sugar in RNA)


Monosaccharides

• Simple sugars containing three to seven C atoms

• (CH20)n


ExampleHexose sugars (the hexoses shown

here are isomers)

ExamplePentose sugars

Glucose Fructose Galactose Deoxyribose Ribose

(a) MonosaccharidesMonomers of carbohydrates


Disaccharides

• Double sugars

• Too large to pass through cell membranes


PLAY Animation: Disaccharides

ExampleSucrose, maltose, and lactose

(these disaccharides are isomers)

Glucose Fructose Glucose Glucose GlucoseSucrose Maltose Lactose

Galactose

(b) DisaccharidesConsist of two linked monosaccharides


Polysaccharides

• Polymers of simple sugars, e.g., starch and glycogen

• Not very soluble


PLAY Animation: Polysaccharides

ExampleThis polysaccharide is a simplified representation of

glycogen, a polysaccharide formed from glucose units.

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

Glycogen


Lipids

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

• Insoluble in water

• Main types:• Neutral fats or triglycerides

• Phospholipids

• Steroids

• EicosanoidsPLAY Animation: Fats


Triglycerides

• Neutral fats—solid fats and liquid oils

• Composed of three fatty acids bonded to a glycerol molecule

• Main functions

• Energy storage

• Insulation

• Protection


Glycerol

+

3 fatty acid chains Triglyceride,or neutral fat

3 watermolecules

(a) Triglyceride formationThree 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


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


Phosphorus-containing

group (polar“head”)

ExamplePhosphatidylcholine

Glycerolbackbone

2 fatty acid chains(nonpolar “tail”)

Polar“head”

Nonpolar“tail”

(schematicphospholipid)

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

attached to the glycerol backbone.


Steroids

• Steroids—interlocking four-ring structure

• Cholesterol, vitamin D, steroid hormones, and bile salts


ExampleCholesterol (cholesterol is the

basis for all steroids formed in the body)

(c) Simplified structure of a steroid

Four interlocking hydrocarbon rings form a steroid.


Eicosanoids

• Many different ones

• Derived from a fatty acid (arachidonic acid) in cell membranes

• Prostaglandins


Other Lipids in the Body

• Other fat-soluble vitamins

• Vitamins A, E, and K

• Lipoproteins

• Transport fats in the blood


Proteins

• Polymers of amino acids (20 types)

• Joined by peptide bonds

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


(a) Generalizedstructure of allamino acids.

(b) Glycineis 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 likelyto participate in

intramolecular bonding.

Aminegroup

Acidgroup


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


Structural Levels of Proteins

PLAY Animation: Introduction to Protein Structure


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

Amino acid Amino acid Amino acid Amino acid Amino acid

PLAY Animation: Primary Structure


α-Helix: The primary chain is coiledto form a spiral structure, which is

stabilized by hydrogen bonds.

β-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 (α-helices) and sheets (β-sheets).

PLAY Animation: Secondary Structure


Tertiary structure of prealbumin(transthyretin), a protein that

transports the thyroid hormonethyroxine in serum and cerebro-

spinal fluid.

(c) Tertiary structure:Superimposed on secondary structure. α-Helices and/or β-sheets are

folded up to form a compact globular molecule held together byintramolecular bonds.

PLAY Animation: Tertiary Structure

Copyright © 2010 Pearson Education, Inc. Figure 2.19d

Quaternary structure ofa functional prealbuminmolecule. Two identical

prealbumin subunitsjoin head to tail to form

the dimer.

(d) Quaternary structure:Two or more polypeptide chains, each with its own tertiary structure,

combine to form a functional protein.

PLAY Animation: Quaternary Structure


Fibrous and Globular Proteins

• Fibrous (structural) proteins

• Strandlike, water insoluble, and stable

• Examples: keratin, elastin, collagen, and certain contractile fibers


Fibrous and Globular Proteins

• 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

• Biological catalysts

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


Activationenergy required

Less activationenergy required

WITHOUT ENZYME WITH ENZYME

Reactants

Product Product

Reactants

PLAY Animation: Enzymes


Characteristics of Enzymes

• Often named for the reaction they catalyze; usually end in -ase (e.g., hydrolases, oxidases)

• Some functional enzymes (holoenzymes) consist of:

• Apoenzyme (protein)

• Cofactor (metal ion) or coenzyme (a vitamin)


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 in

proper position.

Internalrearrangements

leading tocatalysis occur.

Product isreleased. Enzymereturns to original

shape and isavailable to catalyzeanother reaction.

Active site

+ H2O

1 23



Enzyme (E)


Energy isabsorbed;

bond isformed.

Water isreleased.



proper position.



Active site

+ H2O

1 2



Enzyme (E)


Enzyme (E)

Product (P)e.g., dipeptide

Energy isabsorbed;

bond isformed.

Water isreleased. Peptide

bond



proper position.



Product isreleased. Enzymereturns to original

shape and isavailable to catalyzeanother reaction.

Active site

+ H2O

1 23


Summary of Enzyme Action

PLAY Animation: How Enzymes Work


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


Deoxyribonucleic Acid (DNA)

• Four bases:

• adenine (A), guanine (G), cytosine (C), and thymine (T)

• Double-stranded helical molecule in the cell nucleus

• Provides instructions for protein synthesis

• Replicates before cell division, ensuring genetic continuity


Deoxyribosesugar

PhosphateSugar-phosphate

backbone

Adenine nucleotideHydrogen

bond

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


Ribonucleic Acid (RNA)

• Four bases: • adenine (A), guanine (G), cytosine (C), and

uracil (U)

• Single-stranded molecule mostly active outside the nucleus

• Three varieties of RNA carry out the DNA orders for protein synthesis• messenger RNA, transfer RNA, and ribosomal

RNAPLAY Animation: DNA and RNA


Adenosine Triphosphate (ATP)

• Adenine-containing RNA nucleotide with two additional phosphate groups


Adenosine triphosphate (ATP)

Adenosine diphosphate (ADP)

Adenosine monophosphate (AMP)

Adenosine

Adenine

Ribose

Phosphate groups

High-energy phosphatebonds can be hydrolyzed

to release energy.


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


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)

chemistry comes alive: part a -...

Documents