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© 2013 Pearson Education, Inc.

Molecular Interactions

Chapter 2About This Chapter

• Molecules and bonds

• Noncovalent interactions

• Protein interactions

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Biomolecules

• Organic molecules contain carbon– Biomolecules are associated with living organisms

– Four groups– Carbohydrates, lipids, proteins, nucleotides

– Polymers made of repeating unit

– Conjugated proteins: protein molecules combined with another kind of biomolecule. (e.g., lipoproteins; blood transport molecules)

– Glycosylated molecules: molecules to which a carbohydrate has been attached. (e.g., glycoproteins, glycolipids) in cell membranes

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Figure 2.1-1 REVIEW – Biochemistry of Lipids

Fatty Acids

Fatty acids are long chains of carbon atoms bound t o hydrogens,with a carbon (–COOH) or “acid” group at one end of t he chain.

Saturated fatty acids have no double bonds between carbons, s othey are “saturated” with hydrogens. The more saturat ed a fatty acidis, the more likely it is to be solid at room tempe rature.

Palmitic acid, a saturated fatty acid

Linolenic acid, a polyunsaturated fatty acid

Oleic acid, a monounsaturated fatty acid

Polyunsaturated fatty acids have two or more double bonds between carbons in the chain.

Monounsaturated fatty acids have one double bond between twoof the carbons in the chain. For each double bond, the moleculehas two fewer hydrogen atoms attached to the carbon chain.

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Figure 2.1-2 REVIEW – Biochemistry of Lipids

Formation of Lipids

Glycerol plus onefatty acidproduces a monoglyceride .

Monoglyceride

Glycerol plus twofatty acidsproduces a diglyceride .

Glycerol plus threefatty acids produces a triglyceride(triacylglycerol).More than 90% oflipids are in the formof triglycerides.

Glycerol is a simple 3-carbonmolecule that makes up thebackbone of most lipids.

Glycerol Fatty acid

Diglyceride

Fatty acid

Fatty acid

Fatty acid

Fatty acid

Fatty acid

GLYCEROL

GLYCEROL

Triglyceride GLYCEROL

Fatty acid

Figure 2.1-3 REVIEW – Biochemistry of Lipids

Lipid-Related Molecules

Steroids are lipid-related moleculeswhose structureincludes fourlinked carbonrings.

SteroidsEicosanoids

In addition to true lipids, this category includes three types of lipid-related molecules.

Eicosanoids { eikosi, twenty} aremodified 20-carbon fatty acids with acomplete or partial carbon ring atone end and two long carbon chain“tails.”

Phospholipids have 2 fatty acidsand a phosphate group (–H 2PO4).Cholesterol and phospholipids areimportant compounds of animalcell membranes.

Phospholipids

Prostaglandin E 2 (PGE2)

Eicosanoids, such as thromboxanes,leukotrienes, and prostaglandins, actas regulators of physiologicalfunctions.

Cortisol

Cholesterol is the primary sourceof steroids in the human body. Fatty acid

Fatty acid

GLYCEROL

Phosphate group

P

Figure 2.2-1 REVIEW – Biochemistry of Carbohydrates

Monosaccharides

Five Carbon Sugars (Pentoses)

Forms the sugar-phosphatebackbone of RNA

Forms the sugar-phosphate backbone of RNA

Notice that the only differencebetween glucoseand galactose isthe spatialarrangement ofthe hydroxyl(–OH) groups.

Six Carbon Sugars (Hexoses)

Monosaccharides are simple sugars. The most common monosaccharides are the building blocksof complex carbohydrates and have either five carbo ns, like ribose, or six carbons, like glucose.

Ribose Deoxyribose Glucose (dextrose)Fructose Galactose

Figure 2.2-2 REVIEW – Biochemistry of Carbohydrates

Disaccharides

Disaccharides consist of glucoseplus another monosaccharide. Sucrose (table sugar)

* In shorthand chemical notation,the carbons in the rings andtheir associated hydrogenatoms are not written out.Compare this notation to theglucose structure in the rowabove.

Maltose

Glucose* ++++ Fructose Glucose ++++ Glucose

Lactose

Galactose ++++ Glucose

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Figure 2.2-3 REVIEW – Biochemistry of Carbohydrates

Polysaccharides

Polysaccharides are glucosepolymers. All living cells storeglucose for energy in the formof a polysaccharide.

** Chitin and celluloseare structural polysaccharides.

Chitin** Glycogen

Animals

in invertebrateanimals

Glucosemolecules

Digestion of starchor glycogen yields

maltose.

Cellulose**Humans cannotdigest celluloseand obtain itsenergy, even

though it is themost abundantpolysaccharide

on earth.

Plants

Starch

Yeastsand bacteria

Dextran

Figure 2.3-1 REVIEW – Biochemistry of Proteins

Amino Acids

The R groups differ in their size, shape,and ability to form hydrogen bonds orions. Because of the different R groups,each amino acid reacts with othermolecules in a unique way.

The nitrogen (N) in the amino groupmakes proteins our major dietarysource of nitrogen.

All amino acids have a carboxyl group (–COOH), an a mino group(–NH2), and a hydrogen attached to the same carbon. The fourthbond of the carbon attaches to a variable “R” group .

Figure 2.3-2 REVIEW – Biochemistry of Proteins

Amino Acids in Natural Proteins

A few amino acids do not occur in proteins but have importantphysiological functions.

Twenty different amino acids commonly occur in natu ralproteins. The human body can synthesize most of the m, but atdifferent stages of life some amino acids must be o btainedfrom diet and are therefore considered essential am ino acids.

Amino AcidOne-LetterSymbol

Three-LetterAbbreviation

Asparagine

Alanine

Arginine

Asparagine or aspartic acid

Aspartic acid

Cysteine

Glycine

IsoleucineHistidine

Glutamine or glutamic acid

Glutamic acid

Glutamine

Leucine

Proline

Serine

Phenylalanine

LysineMethionine

TyrosineValine

TryptophanThreonine

Gly

IleHis

Leu

Pro

Ser

Phe

LysMet

TyrVal

TrpThr

Asn

Ala

Arg

Asx

Asp

Cys

Glx

Glu

Gln

G

IH

L

PS

F

KM

YV

WT

N

A

R

B

D

C

Z

E

Q

Note:

• Creatine: a molecule that stores energy when it binds to a phosphate group

• Homocysteine: a sulfur-containing amino acid that in excessis associated with heart disease

• γγγγ-amino butyric acid (gamma-amino butyric acid) or GABA: achemical made by nerve cells

Figure 2.3-4 REVIEW – Biochemistry of Proteins

Primary Structure

Structure of Peptides and Proteins

Sequence of amino acids

The 20 protein-forming amino acids assemble into po lymerscalled peptides. The sequence of amino acids in a p eptide chainis called the primary structure . Just as the 26 letters of our alphabet combine to create different words, the 20 amino acidscan create an almost infinite number of combination s.

Peptides range in length from two to two million amino acids :

• Proteins: >100 amino acids• Polypeptide: 10–100 amino acids• Oligopeptide {oligo-, few}: 2–9 amino acids

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Figure 2.4-2 REVIEW – Nucleotides and Nucleic Acids

A nucleotide consists of (1) one or more phosphategroups, (2) a 5-carbon sugar, and (3) a carbon-nitr ogenring structure called anitrogenous base.

Nucleotide

Base

Sugar

Phosphate

Figure 2.4-3 REVIEW – Nucleotides and Nucleic Acids

Nitrogenous Base

Pyrimidines have a single ring.

Purines have a double ring structure.

Adenine (A) Guanine (G) Cytosine (C) Thymine (T) Uracil (U)

Figure 2.4-4 REVIEW – Nucleotides and Nucleic Acids

Deoxyribose{de-, without: oxy-, oxygen}

5-carbon Sugar

Ribose

Figure 2.4-5 REVIEW – Nucleotides and Nucleic Acids

Phosphate

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Figure 2.4-6 REVIEW – Nucleotides and Nucleic Acids

Nucleotide

Single Nucleotide Molecules

Base Sugar Phosphate Groupsconsists of Other Component++++ ++++ ++++ Function

Cell-to-cell communication

Energy capture and transfer

ATP

ADP

NAD

==== Adenine Ribose 3 phosphate groups

Adenine Ribose 2 phosphate groups

Adenine 2 Ribose 2 phosphate groups Nicotinamide

Adenine

++++

++++

++++

++++FAD

cAMP Adenine

====

====

====

==== ++++

Ribose

Ribose

++++

++++

++++

++++

++++

2 phosphate groups Riboflavin ++++

++++

Molecules and Bonds

• Bonds link atoms

• Bonds store and transfer energy

• Molecules versus weaker interactions

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Functional Groups

• Combinations of elements that occur frequently in biological molecules

• Move along molecules as a single unit

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Table 2.1 Common Functional Groups

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Figure 2.5-1 REVIEW – Atoms and Molecules

Helium, He

Protons ++++ neutronsin nucleus ====atomic mass

Protons:determinethe element(atomic number)

Neutrons:determinethe isotope

Electrons:• form covalent bonds• gained or lost createions

• capture and storeenergy

• create free radicals

in orbitals around the nucleus

Atoms

Helium (He) has twoprotons and twoneutrons, so itsatomic number ==== 2,

and its atomic mass ==== 4

2 or more atomsshare electronsto form

Molecules

Water (H2O)

Major EssentialElements

Minor EssentialElements

H, C, O, N, Na,Mg, K, Ca, P,S, Cl

Li, F, Cr, Mn, Fe, Co, Ni,Cu, Zn, Se, Y, I, Zr, Nb,Mo, Tc, Ru, Rh, La

Role of Electrons

• Covelent bond: Electrons shared between atoms

• Ions: Atom or molecule gains or looses an electron and caries a net charge.

• High-energy electrons: Electrons capture energy from environment and transfer it to other atoms

• Free radicals: Unstable molecules with unpaired electron

Isotopes and Ions

An atom that gains orloses neutrons becomes anisotope of the same element.

An atom that gains or loses electrons becomes an ionof the same element.

1H, Hydrogen

2H, Hydrogen isotope

H+, Hydrogen ion

gains aneutron

loses anelectron

Figure 2.5-2 REVIEW – Atoms and Molecules Figure 2.5-3 REVIEW – Atoms and Molecules

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Table 2. 2 Important Ions of the Body

Types of Chemical Bonds – Covalent Bonds

• Covalent bonds– Share a pair of electrons

– Single, double, and triple bonds

– Polar versus nonpolar molecules

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Figure 2.6a-b REVIEW – Molecular Bonds

Covalent Bonds

Nonpolar Molecules

Polar Molecules

Covalent bonds result when atoms share electrons.These bonds require the most energy to make or brea k.

Nonpolar molecules have an evendistribution of electrons. Forexample, molecules composedmostly of carbon and hydrogen tendto be nonpolar.

Polar molecules have regions ofpartial charge ( δδδδ+ or δδδδ -). The mostimportant example of a polarmolecule is water.

δδδδ+ δδδδ+

δδδδ -δδδδ -

Fatty acidHydrogen

Carbon

Negative pole

Positive pole

Water molecule

Types of Chemical Bonds – Ionic Bonds

• Ionic bonds– Atoms gain or lose electrons

– Opposite charges attract

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Figure 2.6c REVIEW – Molecular Bonds

Noncovalent Bonds

Ionic Bonds

Na NaCl Cl

−−−−++++

Sodium atom Chlorine atomSodium ion (Na +)

Chloride ion (Cl −−−−)

Ionic bonds are electrostatic attractions between i ons. A common example is sodium chloride.

Sodium gives up its one weakly heldelectron to chlorine, creating sodium andchloride ions, Na + and Cl −−−−.

The sodium and chloride ions both have stableouter shells that are filled with electrons. Becaus eof their opposite charges, they are attracted toeach other and, in the solid state, the ionic bondsform a sodium chloride (NaCl) crystal.

Ions

• Ions are charged atoms – Cations

– Lost electrons

– Positively charged (+)

– Anions – Gained electrons

– Negatively charged (−)

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Types of Chemical Bonds – Hydrogen and Van der Waals

• Hydrogen bonds– Weak and partial

– Water surface tension

• Van der Waals forces – Weak and nonspecific

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Figure 2.6d REVIEW – Molecular Bonds

Hydrogen Bonds

Hydrogenbonding

Hydrogen bonds form betweena hydrogen atom and a nearbyoxygen, nitrogen, or fluorineatom. So, for example, thepolar regions of adjacentwater molecules allow themto form hydrogen bondswith one another. Hydrogen bonding

between water moleculesis responsible for thesurface tension of water.

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Aqueous Solutions

• Aqueous– Water-based

• Solution– Solute dissolves in solvent

• Solubility– Ease of dissolving

– Hydrophilic

– Hydrophobic

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Figure 2.7-2 REVIEW – Solutions

TERMINOLOGY

Concentration ==== solute amount/volume of solution

A solute is any substance that dissolves in a liquid. The de gree towhich a molecule is able to dissolve in a solvent i s the molecule’s solubility. The more easily a solute dissolves, the higher itssolubility.

A solvent is the liquid into which solutes dissolve. In biolo gicalsolutions, water is the universal solvent.

A solution is the combination of solutes dissolved in a solven t. The concentration of a solution is the amount of solute per unitvolume of solution.

Figure 2.7-3 REVIEW – Solutions

EXPRESSIONS OF SOLUTE AMOUNT

Example

• Mass (weight) of the solute before it dissolves. Usually given in grams (g) or milligrams (mg).

• Molecular mass is calculated from the chemical formula of a molecu le. This is the mass ofone molecule, expressed in atomic mass units (amu) or, more often, in daltons (Da), where1 amu ==== 1 Da.

Molecular mass ==== SUM ××××atomic massof each element

the number of atomsof each element[[[[ ]]]]

What is themolecular massof glucose,C6H12O6? Carbon

Hydrogen

Oxygen 6

612

12.0 amu ×××× 6 ==== 72

1.0 amu ×××× 12 ==== 12

16.0 amu ×××× 6 ==== 96

Molecular mass of glucose ==== 180 amu (or Da)

AnswerElement # of Atoms Atomic Mass of Element

• Moles (mol) are an expression of the number of solute mol ecules, without regard for theirweight. One mole ==== 6.02 ×××× 1023 atoms, ions, or molecules of a substance. One mole of asubstance has the same number of particles as one m ole of any other substance, just as adozen eggs has the same number of items as a dozen roses.

• Gram molecular weight. In the laboratory, we use the molecular mass of a s ubstance tomeasure out moles. For example, one mole of glucose (with 6.02 ×××× 1023 glucose molecules)has a molecular mass of 180 Da and weighs 180 grams . The molecular mass of a substanceexpressed in grams is called the gram molecular wei ght.

• Equivalents (eq) are a unit used for ions, where 1 equivalent ==== molarity of the ion ×××× thenumber of charges the ion carries. The sodium ion, with its charge of ++++1, has one equivalentper mole. The hydrogen phosphate ion (HPO 4

2-) has two equivalents per mole. Concentrationsof ions in the blood are often reported in milliqui valents per liter (meq/L).

Figure 2.7-7 REVIEW – Solutions

EXPRESSIONS OF VOLUME

Volume is usually expressed as liters (L) or millil iters (mL)(milli-, 1/1000). A volume convention common in medicine i sthe deciliter (dL), which is 1/10 of a liter, or 10 0 mL.

deci- (d)

milli- (m)

micro- ( µµµµ)

nana- (n)

pico- (p)

Prefixes

1/10

1/1000

1/1,000,000

1/1,000,000,000

1/1,000,000,000,000

1 ×××× 10-1

1 ×××× 10-3

1 ×××× 10-6

1 ×××× 10-9

1 ×××× 10-12

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Figure 2.7-8 REVIEW – Solutions

EXPRESSIONS OF CONCENTRATION

Answer

Answer

Example

Example

• Percent solutions. In a laboratory or pharmacy, scientists cannot meas ure out solutes bythe mole. Instead, they use the more conventional m easurement of weight. The solute concen-tration may then be expressed as a percentage of th e total solution, or percent solution. A10% solution means 10 parts of a solute per 100 par ts of total solution. Weight/volumesolutions, used for solutes that are solids, are us ually expressed as g/100 mL solution ormg/dL. An out-of-date way of expressing mg/dL is mg % where % means per 100 parts or 100mL. A concentration of 20 mg/dL could also be expre ssed as 20mg%.

Solutions used forintravenous (IV)infusions are oftenexpressed aspercent solutions.How would youmake 500 mL of a5% dextrose(glucose) solution?

5% solution ==== 5 g glucose dissolved in water to make afinal volume of 100 mL solution.

5 g glucose/100 mL ==== ? g/500 mL

25 g glucose with water added to give a final volum e of500 mL

• Molarity is the number of moles of solute in a liter of solu tion, and is abbreviated as eithermol/L or M. A one molar solution of glucose (1 mol/ L, 1 M) contains 6.02 ×××× 1023 molecules ofglucose per liter of solution. It is made by dissol ving one mole (180 grams) of glucose inenough water to make one liter of solution. Typical biological solutions are so dilute thatsolute concentrations are usually expressed as millimoles per liter (mmol/l or mM).

What is themolarity of a5% dextrosesolution?

5 g glucose/100 mL ==== 50 g glucose/1000 mL (or 1 L)

1 mole glucose ==== 180 g glucose

50 g/L ×××× 1 mole/180 g ==== 0.278 moles/L or 278 mM

Figure 2.8a REVIEW – Molecular Interactions

Hydrophilic Interactions

Molecules that have polarregions or ionic bondsreadily interact with the polarregions of water. This enablesthem to dissolve easily inwater. Molecules thatdissolve readily in water aresaid to be hydrophilic(hydro-, water ++++ philos,loving).

Water molecules interact withions or other polar molecules toform hydration shells aroundthe ions. This disrupts thehydrogen bonding betweenwater molecules, therebylowering the freezing tempera-ture of water (freezing pointdepression).

NaCl in solution Glucose molecule in solution

Hydrationshells

Glucosemolecule

Watermolecules

Cl-

Na+

Figure 2.8b REVIEW – Molecular Interactions

Hydrophobic Interactions

Because they have an evendistribution of electrons andno positive or negative poles,nonpolar molecules have noregions of partial charge, andtherefore tend to repel watermolecules. Molecules likethese do not dissolve readilyin water and are said to behydrophobic (hydro-, water++++ phobos, fear). Moleculessuch as phospholipids haveboth polar and nonpolarregions that play critical rolesin biological systems and inthe formation of biologicalmembranes.

Phospholipid molecules have polar heads and nonpolar tails. Phospholipids arrange themselves sothat the polar heads are in contact withwater and the nonpolar tails aredirected away from water.

This characteristic allows the phospho-lipid molecules to form bilayers, thebasis for biological membranes thatseparate compartments.

Water

Water

Hydrophilic head

Hydrophobic tailsHydrophilic head

Polar head(hydrophilic)

Nonpolarfatty acid

tail(hydrophobic)

Molecular models Stylized model

Molecular Shape and Function

• Molecular bonds determine shape– Shape is closely related to function

• Proteins have the most complex and varied shapes– Primary structure: amino acid sequence

– Secondary structures: alpha-helix and beta-pleated sheets– Fibrous proteins

– Tertiary structure– Globular proteins vs fibrous

– Quaternary structure– Multiple polypeptide subunits

– Disulfide bonds (S-S) – covelent bond© 2013 Pearson Education, Inc.