Chapter 2 Atoms, Ions, and Molecules · Atoms, Ions, and Molecules • At its simplest level of...

Chapter 2 Atoms, Ions, and Molecules Sections Covered 2.1 Atomic Structure 2.2 Ions and Ionic Compounds 2.3 Covalent Bonding, Molecules, and Molecular Compounds 2.4 Molecular Structure and Properties of Water 2.5 Acidic and Basic Solutions, pH, and Buffers 2.7 Biological Macromolecules

Atomic Structure

Atoms, Ions, and Molecules

•  At its simplest level of organization, the human body is composed of chemical structures: –  Atoms –  Ions –  Molecules

•  A basic understanding of chemical concepts is necessary to understand physiological processes

Atoms, Ions, and Molecules: Matter, Atoms, Elements, and the Periodic Table

•  Human body is composed of matter –  Three forms:

•  solid (e.g., bone) •  liquid (e.g., blood) •  gas (e.g., oxygen)

•  Matter is composed of atoms –  Atom, smallest particle that exhibits the chemical properties of

an element

•  Elements organized into chart form in the periodic table of elements

•  Elements called major, lesser, or trace based on percentage by weight in the body –  Six major elements, over 98% total –  Six minor elements, less than 1% total

Periodic Table of Elements

1 3 1.008

6.941 Li

Na 22.99

Rb 85.47 55

Cs 132.9

Fr 223.0

4 Be 9.012 12

Mg 24.31

Ca 40.08

Sr 87.62 56

Ba 137.3

Ra 226.0

Sc 44.96

88.91 57

La 138.9

Ac 227.0

Ti 47.87

Zr 91.22 72

Hf 178.5

Rf 267.0

Nb 92.91

Db 268.0

Cr 52.00

Mo 95.94 74

Sg 271.0

Mn 54.94

Tc 98.00 75

Re 186.2

Bh 272.0

Fe 55.85

Ru 101.1 76

Os 190.2

Hs 270.0

Co 58.93

Rh 102.9 77

Ir 192.2

Mt 276.0

Ni 58.69

Pd 106.4 78

Pt 195.1

Ds 281.0

Cu 63.55

Ag 107.9 79

Au 197.0

Rg 274

Zn 65.38

Cd 112.4 80

Hg 200.6

277 Uub

B 5 10.81

Al 26.98

Ga 69.72 49

In 114.8 81

Tl 204.4

277 Uut

Si 28.09

Ge 72.64

Sn 118.7 82

Pb 207.2

Uuq 289.0

12.01 C 6

P 30.97

As 74.92

Sb 121.8 83

Bi 209.0 115

Uup 288.0

14.01 N 7

Se 78.96

Te 127.6 84

Po 209.0

293.0 Uuh

F 9 19.00

17 Cl 35.45

Br 79.90

126.9 85

At 210.0

292.0 Uus

2 He 4.003 10

Ne 20.18

Ar 39.95

Kr 83.80

Xe 131.3 86

Rn 222.0

294.0 Uuo In

Increasing electronegativity

1 H Atomic number

Element symbol Atomic mass number

58 59 60 61 62 63 64 65 66 67 68 69 70 71

Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu 140.1 140.9 144.2 145.0 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0

90 91 92 93 94 95 96 97 98 99 100 101 102 103

U Th Pa Np Pu Am Cm Bk Cf Es Fm Md No Lr 232.0 231.0 238.0 237.0 244.0 243.0 247.0 247.0 251.0 252.0 257.0 258.0 259.0 262.0

IIA IIIA IVA VA VIA VIIA VIIIA

Most Common Elements of the Human Body

Major elements (collectively compose more than 98% of body weight)

Lesser elements (collectively compose less than 1% of body weight)

P Ca N H C O Symbol % Body weight

Oxygen

Carbon

Hydrogen

Phosphorus

Calcium

Nitrogen

Symbol % Body weight

Sulfur

Potassium

Magnesium

Chlorine

Sodium

The components of an atom •  Atoms composed of three subatomic particles:

–  Protons •  mass of one atomic mass unit (amu) •  positive charge of one (+1)

–  Neutrons •  mass of one amu •  no charge

–  Electrons •  1/1800th mass of a proton or neutron •  negative charge of one (-1) •  located at varying distance from the nucleus in regions called orbitals

The periodic table •  Elements differ in number of subatomic particles •  Periodic table displays:

–  Chemical symbol •  unique to each element •  usually identified by first letter, or first letter plus an additional letter

–  e.g., C is carbon –  Atomic number

•  number of protons in an atom of the element •  located above symbol name •  elements arranged by anatomic number within rows

–  Average atomic mass •  mass of both protons and neutrons •  shown below the element’s symbol on the table

Determining the number of subatomic particles •  Proton number = atomic number •  Neutron number = atomic mass – atomic number

•  neutron number = (p + n) – p •  neutron number of Na = 23 – 11 = 12

•  Electrons number = proton number

Diagramming Atomic Structures •  An atom has shells of electrons

surrounding the nucleus –  Each shell with a given energy

level –  Each shell holding a limited

number of electrons –  Innermost shell two electrons,

second shell up to eight –  Shells close to the nucleus: must be

filled first

Shell model

Nucleus: Proton (+)

Neutron (no charge)

8 electrons

8 protons 8 neutrons

Energy shell

Electron shells: Electron (–)

Atoms, Ions, and Molecules: Isotopes

•  Isotopes are different atoms of the same element –  Have same number of protons and electrons –  Have different numbers of neutrons –  Exhibit essentially identical chemical characteristics –  One usually predominant

•  Carbon exists in three isotopes: –  carbon-12, with 6 neutrons

•  most prevalent type –  carbon-13, with 7 neutrons

–  carbon-14 , with 8 neutrons

Carbon–13 Carbon–14 6 protons 7 neutrons 6 electrons

6 protons 8 neutrons 6 electrons

Carbon–12

6 electrons

6 protons 6 neutrons

Atoms, Ions, and Molecules: Chemical Stability and the Octet Rule

•  Periodic table organized into columns based on number of electrons in outer shell (valence shell) –  Column one, with hydrogen, lithium, sodium, potassium

•  all with one electron in their outer shell –  Each consecutive column with one additional electron in outer shell –  Elements in column VIIIA with a full valence shell

•  results in chemical stability •  helium, neon, etc., chemically inert (noble gases)

Atoms, Ions, and Molecules: Chemical Stability and the Octet Rule

•  Elements tend to lose, gain, or share electrons to obtain complete outer shells with eight electrons –  Known as the octet rule

1 2 3 4 5 6 7 8

IA IIA IIIA IVA VA VIA VIIA VIIIA

Ar Cl Si Al

Number of valence electrons

Ions and Ionic Compounds

•  Chemical compounds –  Stable associations between two or more elements combined in a

fixed ratio –  Classified as ionic or molecular

•  Ionic compounds are structures composed of ions held together in a lattice of ionic bonds

Ions and Ionic Compounds: Ions

•  Ions –  Are groups of atoms with a positive or a negative charge –  Are produced from the loss or gain of an electron or

electrons –  Are used very commonly in the body with significant

physiological functions •  e.g., Na+ for electrical signals in neurons •  e.g., Ca2+ for blood clotting and muscle contraction •  e.g., Cl- in stomach acid, and many more

Losing electrons and the formation of cations •  Sodium can reach stability by donating an electron

–  Now satisfies the octet rule –  Now has 11 protons and 10 electrons –  Charge is +1

•  Ions with positive charge called cations

Gaining electrons and the formation of anions •  Chlorine can reach stability by gaining an electron

–  Now satisfies the octet rule –  Now has 17 protons and 18 electrons –  Charge is -1

•  Ions with negative charge called anions

Ions and Ionic Compounds: Ionic Bonds

•  Cations and anions may bind together in ionic bonds –  Salts formed –  For example, table salt (NaCl)

•  Each sodium atom donates one outer shell electron to a chlorine atom

•  Sodium and chlorine ions are held together by ionic bonds in a lattice crystal structure

•  This is an ionic compound

Formation of and Ionic Bond Involving Sodium and Chloride

+ = Na 11p

Cl 17p

11p Cl–

Cl– Na+ Cl–

Cl– Na+ Na+

Cl– Na+ Cl–

(d) Lattice salt crystal of NaCl (c) Sodium ion (Na+) (a) Sodium atom (Na) (b) Chlorine atom (Cl) Chloride ion (Cl–)

Covalent Bonding, Molecules, and Molecular Compounds

•  Sharing of electrons between atoms results in a covalently bonded molecule

•  Most molecules are composed of two or more different elements –  Termed molecular compounds

•  examples include carbon dioxide (CO2) but not molecular oxygen (O2)

Covalent Bonding, Molecules, and Molecular Compounds: Covalent Bonds

•  A covalent bond is formed when atoms share electrons –  Occurs when both atoms require electrons –  Occurs with atoms that have four to seven electrons in their outer shell

•  Four elements of the human body form covalent bonds most commonly: –  oxygen (O) –  carbon (C) –  hydrogen (H) –  nitrogen (N)

Single, double, and triple covalent bonds •  Single covalent bond

–  One pair of electrons shared •  e.g., between two hydrogen atoms

•  Double covalent bond –  Two pairs of electrons shared

•  e.g., between two oxygen atoms

•  Triple covalent bond –  Three pairs of electrons shared

•  e.g., between two nitrogen atoms

Hydrogen gas (H2) Single bond

Double bond

Triple bond

Nitrogen gas (N2)

Oxygen gas (O2)

Single covalent bond

Double covalent bond

Triple covalent bond

(c) 26

Single, double, and triple covalent bonds (continued) •  Carbon needs four electrons to satisfy the octet rule

–  Can be obtained in multiple different ways

H H C C

H H H H

H H O O O C C

C O O C H H H C O

C H H H

Methane (CH4) Carbon dioxide (CO2) Ethanol (C2H5OH)

(c) (b) (a)

Carbon Skeleton Formation •  Carbon can bond in straight chains, branched chains,

or rings –  Carbon present where lines meet at an angle; additional atoms hydrogen

C C C C C C C C C C C C C C C C C C

C C C C

Straight chain Branched chain Ring

(a) (b) (c)

CH3 H3C

Nonpolar and polar covalent bonds •  Atoms in a covalent bond may share electrons equally

or unequally –  How they share is determined by electronegativity

•  the relative attraction of each atom for electrons •  high electronegativity = electrons spend more time orbiting the nucleus •  determined by the number of protons in the nucleus and the proximity of

valence electrons –  Two atoms of the same element have same electronegativities –  Share electrons equally in a nonpolar covalent bond

Nonpolar and polar covalent bonds (continued) •  Atoms in a covalent bond may share electrons equally

or unequally –  Atoms with different electronegativity share electrons unequally –  This results in a polar covalent bond –  Exception is the bond between carbon and hydrogen, considered

nonpolar

Nonpolar and polar covalent bonds (continued) •  More electronegative atom develops a partial

negative charge –  Less electronegative atom develops a partial positive charge –  In a bond between oxygen and hydrogen, oxygen is slightly negative,

hydrogen slightly positive

Covalent Bonding, Molecules, and Molecular Compounds: Intermolecular Attractions

•  Intermolecular attractions –  Weak chemical attractions between molecules –  Collectively important in maintaining the shape of complex molecules

such as DNA and protein –  One type, the hydrogen bond

•  forms between polar molecules •  attraction between partially positive hydrogen atom and a partially

negative atom •  individually weak, collectively strong •  influences how water molecules behave

Molecular Structure and Properties of Water

Molecular Structure of Water and the Properties of Water: Molecular Structure

•  Water –  Composes two-thirds of the

human body by weight –  Polar molecule composed of

one oxygen atom bonded to two hydrogen atoms

–  Oxygen atom with two partial negative charges

–  Hydrogen with a single positive charge

–  Can form four hydrogen bonds with adjacent molecules

•  central to water’s properties

Water (H2O)

Hydrogen bonds

δ+ δ+

δ– δ–

(a) (b)

Water is a polar molecule due to unequal sharing of electrons.

Hydrogen bonds form between water molecules.

δ+ δ–

Molecular Structure of Water and the Properties of Water: Properties

Cohesion, surface tension, and adhesion •  Cohesion

–  The attraction between water molecules due to hydrogen bonding

•  Surface tension –  The inward pulling of cohesive forces at the surface of water –  Causes moist sacs of air in the lungs to tend to collapse

•  surfactant (mixture of lipids and proteins) helps prevent this

•  Adhesion –  The attraction between water molecules and a substance other

than water

High specific heat and high heat of vaporization •  Temperature

–  The measure of kinetic energy of atoms or molecules within a substance

•  Specific heat –  The amount of energy required to increase the temperature of 1 gram of

a substance by 1 degree Celsius –  Water’s value extremely high due to energy needed to break hydrogen

bonds –  Contributes to body temperature constancy

High specific heat and high heat of vaporization (continued)

•  Heat of vaporization –  The heat required for the release of molecules from a liquid phase into

a gaseous phase for 1 gram of a substance –  Water’s value very high due to hydrogen bonding –  Why sweating is an effective means of cooling the body

•  Excess heat dissipated as water evaporates

Molecular Structure of Water and the Properties of Water: The Universal Solvent

•  Water is the solvent of the body •  Substances that dissolve in water are called solutes •  Water called the universal solvent because most

substances dissolve in it

•  Some polar molecules and other charged substances dissolve (disperse) within water –  Substances termed hydrophilic, “water-loving”

•  Nonpolar molecules do not dissolve within water –  Substances termed hydrophobic, “water-fearing”

Acidic and Basic Solutions, pH, and Buffers

Acidic and Basic Solutions, pH, and Buffers: pH, Neutralization, and the Action of Buffers

•  The pH is a measure of H+

–  The relative amount of H+ in a solution –  Expressed as a number between 0 and 14 –  The inverse of the log for a given H+ concentration

•  greater H+ = lower pH value

•  The pH of plain water is 7 –  Water dissociates to produce 1/10,000,000 ions per liter

•  Moving from one increment to another is a tenfold change

–  E.g., a pH of 6 has 10 times greater concentration of H+ than pure water

Interpreting the pH scale •  Solutions with equal concentrations of H+ and OH-

–  Are neutral –  Have a pH of 7

•  Solutions with greater H+ than OH- –  Are acidic –  Have a pH < 7

•  Solutions with greater OH- than H+ –  Are basic (alkaline) –  Have a pH >7

0 1 2 3 4 5 6 8 9

Examples pH Value

Sodium hydroxide (NaOH): 14

Household bleach: 12

Household ammonia: 10.5–11

Antacid: 10.5

Seawater: 8

Human blood: 7.4 Pure water: 7 Milk, saliva: 6.3–6.6 Urine: 6

Tomato juice: 4.7

Grapefruit juice: 3 Wine: 2.4–3.5 Lemon juice, stomach acid: 2–3

Hydrochloric acid (HCl): 1

10–1

10–2 H+

H+ Decreasing

Increasing

H+ Concentration

H+ > OH–

Neutral

Acidic

10–3

10–4

10–5

10–6

10–7

10–8

10–9

10–10

10–11

10–12

10–13

10–14 14

Decreasing

Increasing

•  Neutralization occurs when an acidic or basic solution is returned to neutral –  Acids neutralized by adding base

•  e.g., medications to neutralize stomach acid containing base –  Bases neutralized by adding acid

•  Buffers help prevent pH changes if excess acid or base is added –  Act to accept H+ from excess acid or donate H+ to neutralize base

•  carbonic acid (weak acid) and bicarbonate (weak base) buffer blood pH •  both help maintain pH in a critical range

Biological Macromolecules

Biological Macromolecules: General Characteristics

•  Organic molecules, molecules that contain carbon –  Most are a component of living organisms –  Biological macromolecules (biomolecules) are a subset

•  Inorganic molecules, all other molecules

•  Four classes of biomolecules in living systems: I.  Lipids II.  Carbohydrates III.  Nucleic acids IV.   Proteins

Polymers •  Molecules made up for repeating subunits, termed

monomers –  Monomers identical or similar in chemical structure –  Examples are carbohydrates, nucleic acids, proteins

•  carbohydrates with sugar monomers •  nucleic acids with nucleotide monomers •  proteins with amino acid monomers

•  Dehydration synthesis (condensation) –  Occurs during the synthesis of biomolecules –  One subunit looses an –H –  Other subunit loses an –OH –  New covalent bond formed and water produced

•  Hydrolysis reaction –  Occurs during the breakdown of biomolecules –  An –H added to one subunit –  An –OH added to another subunit

Dehydration

Synthesis

Hydrolysis

Digestion

•  h"ps://www.youtube.com/watch?v=ZMTeqZLXBSo

Biological Macromolecules: Lipids

I.  Lipids –  Diverse group of fatty, water-insoluble compounds –  Not composed of monomers –  Function as stored nutrients, cellular membrane

components, and hormones –  Occur in four primary classes:

1)   Triglycerides 2)   Phospholipids 3)   Steroids 4)   Eicosanoids

1) Triglycerides: energy storage –  Most common form of lipid in living things –  Used for long-term energy storage in adipose tissue –  Also used for structural support, cushioning, and insulation –  Formed from a glycerol molecule and three fatty acids

•  Fatty acids – Are varied in length – Are varied in the number of double bonds

»  saturated if they lack double bonds »  unsaturated if they have one double bond »  polyunsaturated if they have two or more double bonds

Triglyceride

H C H C

H C H C H

C H C H

H C H H

H C H C

H C H H

O O O O

H C H C

H C H C H

C H C H

H C H H

H C H C

H C H H

C O C O C O

Hydroxyl groups Carboxylic acid

Lipogenesis (occurs through dehydration)

Lipolysis (occurs through hydrolysis)

Glycerol Three fatty acids (vary in length and number of double bonds between carbons)

Triglyceride (b)

Triglyceride

2) Phospholipids: Membranes –  Amphipathic molecules forming cell membranes –  Phospholipid structure similar to a triglyceride

•  one end of the glycerol with a polar phosphate group with another organic group

•  group substitutes for a fatty acid •  glycerol, phosphate, and organic group is polar

–  known as the hydrophilic head •  fatty acid group is nonpolar

–  known as the hydrophobic tails

3) Steroids: ringed structures including some hormones –  Composed of hydrocarbons arranged in a multiringed structure –  Differ in the side chains extending from their rings –  Include cholesterol, steroid hormones (e.g., testosterone), and bile salts

•  cholesterol component of membranes •  precursor to other steroid synthesis

4) Eicosanoids: locally acting hormones –  Modified 20-carbon fatty acids –  Synthesized from arachidonic acid, membrane component –  Local signaling molecules –  Have functions in the inflammatory response, in the nervous system,

and all body systems –  Four classes:

•  prostaglandins •  prostacyclins •  thromboxanes •  leukotrienes

Major Classes of

Lipids

Biological Macromolecules: Carbohydrates

II.  Carbohydrates –  An –H and an –OH usually attached to every carbon –  Chemical formula is (CH2O)n

•  n the number of carbon atoms

–  Monosaccharides •  simple monomers

–  Disaccharides •  formed from two monosaccharides

–  Polysaccharides •  formed from many monosaccharides

Glucose and Glycogen •  Glucose

–  Six-carbon carbohydrate –  Most common monosaccharide –  Primary nutrient supplying energy to cells –  Concentration carefully maintained –  Bound into the polysaccharide glycogen during glycogenesis

•  Liver and skeletal muscle store excess glucose following a meal –  Broken down from glycogen during glycogenolysis

•  Liver breaks down glucose from glycogen as needed

C C C H H

H (b) Glycogen (a) Glucose

Glycogenolysis

Glycogenesis

Glucose and Glycogen

Other types of carbohydrates •  Other monosaccharides

–  Five carbon monosaccharides (pentose sugars) •  ribose and deoxyribose in nucleic acids

•  Disaccharides, two sugars bonded together –  E.g., sucrose (table sugar), lactose (milk sugar), and maltose (malt

sugar)

Other types of carbohydrates (continued) •  Polysaccharides, three or more sugars bonded

–  Glycogen most common in animals –  Starch and cellulose found in plants

•  plant starch a major nutritional source of glucose •  celullose a source of fiber (nondigestible substances)

Other Simple Carbohydrates

H H H O H

H H H H H

H H HO H

O H O H H

O H H H

O H H O

Monosaccharides

6–carbon sugars (hexose)

5–carbon sugars (pentose)

CH2OH OHCH2 OHCH2

CH2OH CH2OH CH2OH CH2OH CH2OH

Galactose Fructose Deoxyribose Ribose

Maltose Lactose Sucrose

OH OH HO

OH OH OH

OH CH2OH

OH CH2OH OH

Disaccharides

OH OH HO

OH CH2OH

OH OH OH

*You do NOT have to memorize these structures!

Biological Macromolecules: Nucleic Acids

III. Nucleic acids –  Macromolecules that store and transfer genetic

information in cells –  Two classes deoxyribonucleic acid (DNA) and

ribonucleic acid (RNA) •  both polymers of nucleotide monomers

The nucleotide monomer –  Three components: sugar,

phosphate group, and a nitrogenous base

•  sugar a five-carbon pentose

•  phosphate group attached at carbon five

•  nitrogenous base attached to same sugar at carbon one

•  nitrogenous base with single-ring or double-ring structure

O P O–

–O O O

P N N N Phosphate group

(a) Nucleotide monomer

Nitrogenous base

OH in RNA

H in DNA

The nucleotide monomer (continued) •  Five types of nitrogenous bases

– Single-ring bases: pyrimidines •  cytosine, uracil, and thymine

– Double-ring nitrogenous bases: purines •  adenine and guanine

Nucleic Acids Pu

NH2 NH2

C N C O

C O O C

O N C C

C N C HC

C O Uracil (U)

(RNA only)

NH NH HC

Guanine (G) (both DNA and RNA)

Cytosine (C) (both DNA and RNA)

Adenine (A) (both DNA and RNA)

(b) Nitrogenous bases

Thymine (T) (DNA only)

Deoxyribonucleic acid (DNA) –  Double-stranded nucleic acid –  Found in chromosomes in the nucleus and in mitochondria –  Has deoxyribose sugar, phosphate, and one of four nitrogenous bases:

•  adenine, guanine, cystosine, or thymine •  does not contain uracil

–  Double-strands held together by hydrogen bonds •  form between complementary bases •  thymine paired with adenine; guanine paired with cytosine

Ribonucleic acid (RNA) –  Single-stranded nucleic acid –  Found in the nucleus and within cytoplasm of the cell –  Has ribose sugar, phosphate, and one of four nitrogenous bases:

•  adenine, guanine, cystosine, or uracil •  does not contain thymine

Nucleic Acids

P O A O G

(c) RNA (single–stranded)

(d) DNA (double–stranded)

Nucleotide

Sugar–phosphate “backbone”

Nitrogenous base

Deoxyribose sugar

Phosphate group

Unique to RNA

Phosphate group

Nitrogenous base

Ribose sugar

Nucleotide

Phosphodiester bonds

Unique to DNA

3′ 3′

5′ Hydrogen bonds

between nitrogenous bases

Other important nucleotides •  Adenosine triphosphate (ATP)

–  Nucleotide composed of nitrogenous bases adenine, ribose sugar, and three phosphate groups

–  Covalent bonds between last two phosphate groups •  release energy when broken

–  Central molecule in chemical energy transfer within cells

Biological Macromolecules: Proteins

IV.   Proteins –  serve a vast array of functions –  Serve as catalysts (enzymes) in metabolic reactions –  Act in defense –  Aid in transport –  Contribute to structural support –  Cause movement –  Perform regulation –  Provide storage

Protein FuncCons

General protein structure •  Proteins composed of one or more strands of

monomers •  Monomers are amino acids

–  20 total in living organisms –  Have an amine and a carboxylic acid functional group

•  both covalently linked to same carbon atom –  Carbon also covalently bonded to a hydrogen and different side chain

structures •  referred to as the R group •  distinguish different amino acids from one another

General protein structure (continued) •  Amino acids covalently linked by peptide bonds –  Formed during dehydration synthesis reaction –  Occur between amine group of one amino acid and the carboxylic group of

another •  – H lost from the amine group •  – OH lost from the carboxylic acid

–  N-terminal end has free amine group –  C-terminal end has free carboxyl group

Proteins

A m i n e

R OH C C H N O H H R OH C C H N O H H

R O H C C H N O H H

R C N H H O H C O

R C C H N O N H

C O H R C N H H

R C C O N H

C O H R C N H H

R C C O N H

C O H R C N H H

R C C O N H

R C H C O H H

Amino acid Peptide bond

Peptide bond Amine Carboxylic acid

R group (1 of 20 different structures)

N–terminal C–terminal

Polymer protein

Carboxylic acid

Amino Acids N

O H C C O H

C C O H

H C C O H C C O H

C C O H

O C C C O H C C O H

C C O H

C C O H S

C C O H S H

O H C C O H C C C O H

C C O H O C

C C O H

C C C O H H

C C C O H

C C C O H Glycine (Gly)

Valine (Val)

Isoleucine (Ile)

Leucine (Leu)

Phenylalanine (Phe)

Tryptophan (Trp)

Allows bends in protein chain

Forms disulfide bond

Always the first amino acid in a protein sequence (may be removed following synthesis of protein)

Glutamic acid (Glu)

Lysine (Lys)

Histidine (His)

(–) Charge (+) Charge

Aspartic acid (Asp)

Arginine (Arg)

Tyrosine (Tyr)

Glutamine (Gln)

Asparagine (Asn)

Serine (Ser)

Threonine (Thr)

Methionine (Met)

Proline (Pro)

Cysteine (Cys)

NH2 OH NH2

CH2 CH2

OH OH NH2 NH2

C C OH NH2

CH2 NH2+

NH2 OH OH

CH2 CH2

NH2 OH

OH CH3

NH2 OH OH

NH2 OH OH NH2

CH2 CH2

CH3 CH3 CH3

OH OH NH2 NH2 OH NH2

CH2 CH3 CH3

OH NH2

NH2 NH2

Alanine (Ala)

Protein Structure: Amino Acid Sequence and Protein Conformation

•  Primary structure, linear sequence of amino acids

N H H C C O R N

Primary structure

Linear sequence of amino acids joined by peptide bonds

Peptide bond

Amino acid

•  Secondary structures, structural patterns from hydrogen bonds –  Confer unique characteristics –  Two types:

•  alpha helix, spiral coil •  beta sheet, planar pleat

arrangement

Secondary structure

Structural patterns within a protein that result from hydrogen bonds formed between amino acids

Hydrogen bonds

Beta sheet (planar pleats) Alpha helix (spiral coil)

•  Tertiary structure, final three-dimensional shape of polypeptide chain –  Two categories

distinguished: •  globular proteins,

compact shape •  fibrous proteins,

extended linear molecules (c)

Globular protein Fibrous protein

Tertiary structure Final 3–dimensional shape of a protein, which

contains repeating secondary structures

•  Quaternary structure, present in proteins with two or more polypeptide chains –  E.g., hemoglobin

with its four polypeptide chains

Quaternary structure

Molecule composed of two or more separate proteins

Three fibrous proteins Globular

protein

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