Chapter 2

Principles of Human Anatomy and Physiology, 11e 1

Chapter 2

The Chemical Level of Organization

Lecture Outline


INTRODUCTION

• Since chemicals compose your body and all body activities are chemical in nature, it is important to become familiar with the language and fundamental concepts of chemistry.


Chapter 2 The Chemical Level of Organization

• Matter– elements– atoms and molecules

• Chemical bonds• Chemical energy• Chemical reactions• Inorganic compounds• Organic compounds


Basic Principles

• Chemistry is the science of the structure and interactions of matter.

• Matter is anything that occupies space and has mass.– Mass is the amount of matter a substance contains– weight is the force of gravity acting on a mass.

• Describe two ways that you could change your weight. Describe two ways that you could change your weight.


HOW MATTER IS ORGANIZED

• Chemical Elements– All forms of matter are composed of chemical elements

which are substances that cannot be split into simpler substances by ordinary chemical means.

– Elements are given letter abbreviations called chemical symbols.

• Oxygen (O), carbon (C), hydrogen (H), and nitrogen (N) make up 96% of body weight.

• These elements, together with calcium (Ca) and phosphorus (P) make up 98.5% of total body weight.

– Trace elements are present in tiny amounts • copper, tin, selenium & zinc


Review

• Table 2.1 lists the major and trace elements of the human body.


Structure of Atoms

• Units of matter of all chemical elements are called atoms. An element is a quantity of matter composed of atoms of the same type.

• Atoms consist of a nucleus, which contains positively charged protons and neutral (uncharged) neutrons, and negatively charged electrons that move about the nucleus in energy levels (Figure 2.1).


Structure of Atoms

• Atoms are the smallest units of matter that retain the properties of an element

• 3 types of subatomic particles– protons, neutrons and electrons

• Nucleus: protons (p+) & neutrons (neutral charge)

• Electrons (e-) surround the nucleus as a cloud (electron shells are designated regions of the cloud)


Electron Shells

• Most likely region of the electroncloud in which to find electrons

• Each electron shell can hold onlya limited number of electrons

– first shell can hold only 2 electrons

– 2nd shell can hold 8 electrons

– 3rd shell can hold 18 electrons

– higher shells (up to 7) hold many more electrons

• Number of electrons = number of protons

• Each atom is electrically neutral; charge = 0


Structure of Atoms

• Electrons revolve around the nucleus of an atom tending to spend most of the time in specific atomic regions, called shells (Figure 2.1a).– Each shell can hold a certain maximum number of

electrons.– The first shell, the one nearest the nucleus, can hold a

maximum of 2 electrons; the second shell, 8; the third shell;18,the fourth shell, 18; and so on (Figure 2.1b).

– The number of electrons in an atom of a neutral element always equals the number of protons.


Atomic Number and Mass Number

• Atomic Number– The number of protons in the nucleus of an atom– The number of protons in the nucleus makes the atoms

of one element different from those of another as illustrated in Figure 2.2.

– Since all atoms are electrically neutral, the atomic number also equals the number of electrons in each atom.


Atomic Number & Mass Number

• Atomic number is number of protons in the nucleus. .• Mass number is the sum of its protons and neutrons.



• The mass number of an atom is the sum of the numbers of protons and neutrons.– Different atoms of an element that have the same

number of protons but different numbers of neutrons are called isotopes.

• Isotopes– Stable isotopes do not change their nuclear structure

over time.– Certain isotopes called radioactive isotopes are unstable

because their nuclei decay to form a simpler and thus more stable configuration.

– Radioactive isotopes can be used to study both the structure and function of particular tissues.


Atomic Mass

• Mass is measured as dalton (atomic mass unit)– neutron has mass of 1.008 daltons– proton has mass of 1.007 daltons– electron has mass of 0.0005 dalton

• Atomic mass (atomic weight) is close to the mass number of its most abundant isotope.



• Atomic Mass– The atomic mass, also called the atomic weight, of an

element is the average mass of all its naturally occurring isotopes and reflects the relative abundance of isotopes with different mass numbers.

– The mass of a single atom is slightly less than the sum of the masses of its neutrons, protons, and electrons because some mass (less than1%) was lost when the atom’s components came together to form an atom.


Ions, Molecules, & Compounds

• Ions are formed by ionization– an atom that gave up or gained an electron– written with its chemical symbol and (+) or (-)

• Molecule– atoms share electrons– written as molecular formula showing the number

of atoms of each element (H2O)


Ions, Molecules, Free Radicals, and Compounds

• If an atom either gives up or gains electrons, it becomes an ion - an atom that has a positive or negative charge due to having unequal numbers of protons and electrons.

• When two or more atoms share electrons, the resulting combination is called a molecule (Figure 2.3a).


Free Radicals• A free radical is an electrically charged atom or group of atoms with an

unpaired electron in its outermost shell (Figure 2.3b).• Unstable and highly reactive; can become stable

– by giving up an electron– taking an electron from another molecule (example: breaking apart

important body molecules)• Antioxidants are substances that inactivate oxygen-derived free radicals


Free Radicals & Your Health

• Produced in your body by absorption of energy in ultraviolet light in sunlight, x-rays, by breakdown of harmful substances, & during normal metabolic reactions

• Linked to many diseases -- cancer, diabetes, Alzheimer, atherosclerosis and arthritis

• Damage may be slowed with antioxidants such as vitamins C and E, selenium & beta-carotene (precursor to vitamin A)


CHEMICAL BONDS

• The atoms of a molecule are held together by forces of attraction called chemical bonds.

• The likelihood that an atom will form a chemical bond with another atom depends on the number of electrons in its outermost shell, also called the valence shell.


CHEMICAL BONDS

• An atom with a valence shell holding eight electrons (2 electrons for hydrogen and neon) is chemically stable, which means it is unlikely to form chemical bonds with other atoms.

• To achieve stability, atoms that do not have eight electrons in their valence shell (or 2 in the case of H and He) tend to empty their valence shell or fill it to the maximum extent.

• octet rule.– Atoms with incompletely filled outer shells tend to

combine with each other in chemical reactions to produce a chemically stable arrangement of eight valence electrons for each atom.


Ionic Bonds

• When an atom loses or gains a valence electron, ions are formed (Figure 2.4a).– Positively and negatively charged ions are attracted to

one another. – Cations are positively charged ions that have given up

one or more electrons (they are electron donors).– Anions are negatively charged ions that have picked up

one or more electrons that another atom has lost (they are electron acceptors).


Ionic Bonds

• When this force of attraction holds ions having opposite charges together, an ionic bond results.– Sodium chloride is formed by ionic bonds (Figure 2.4)

• In general, ionic compounds exist as solids but some may dissociate into positive and negative ions in solution. Such a compound is called an electrolyte.

• Table 2.2 lists the names and symbols of the most common ions and ionic compounds in the body.


The Ionic Bond in Sodium Chloride

• Sodium loses an electron to become Na+ (cation)

• Chlorine gains an electron to become Cl- (anion)

• Na+ and Cl- are attracted to each other to form the compound sodium chloride (NaCl) -- table salt

• Ionic compounds generally exist as solids


Covalent Bonds

• Covalent bonds are formed by the atoms of molecules sharing one, two, or three pairs of their valence electrons.– Covalent bonds are common and are the strongest

chemical bonds in the body.– Single, double, or triple covalent bonds are formed by

sharing one,two, or three pairs of electrons, respectively (Figures 2.5a – d).

• Covalent bonds may be nonpolar or polar.– In a nonpolar covalent bond, atoms share the electrons

equally; one atom does not attract the shared electrons more strongly than the other atom (Figures 2.5a –.d).


Covalent Bonds

• Atoms share electrons to form covalent bonds

• Electrons spend most of the time between the 2 atomic nuclei– single bond = share 1 pair– double bone = share 2 pair– triple bond = share 3 pair

• Polar covalent bonds share electrons unequally between the atoms involved


Polar Covalent Bonds

• Unequal sharing of electrons between atoms. (Figure 2.6).

• In a water molecule, oxygen attracts the hydrogen electrons more strongly– Oxygen has greater electronegativity as indicated by the

negative Greek delta sign.


Hydrogen Bonds

• Approximately 5% as strong as covalent bonds• Useful in establishing links between molecules or

between distant parts of a very large molecule

• Large 3-D molecules areoften held together by a large number of hydrogen bonds.


Hydrogen Bonds

Hydrogen bonds are weak intermolecular bonds; they serve as links between molecules (usually).

• help determine three-dimensional shape (Figure 2.7) • give water considerable cohesion which creates a

very high surface tension


Chemical Reactions

• New bonds form and/or old bonds are broken.• Metabolism is “the sum of all the chemical reactions in the body.”• Law of conservation of mass

– The total mass of reactants equals the total mass of the products. (Count the number of atoms of each element below.)


CHEMICAL REACTIONS

• A chemical reaction occurs when new bonds are formed or old bonds break between atoms (Figure 2.8).– The starting substances of a chemical reaction are

known as reactants.– The ending substances of a chemical reaction are the

products.• Remember: In a chemical reaction, the total mass of the

reactants equals the total mass of the products (law of conservation of mass).

• Metabolism refers to all the chemical reactions occurring in an organism and involves links between chemical reactions in different parts of the body, or even different parts of a cell.


Forms of Energy and Chemical Reactions

• Energy is the capacity to do work.• Kinetic energy is the energy associated with matter in

motion.– Temperature is an indirect measure of molecular motion.

• Potential energy is energy stored by matter due to its position.– Chemical energy is a form of potential energy stored in

the bonds of compounds or molecules.• The total amount of energy present at the beginning and

end of a chemical reaction is the same; energy can neither be created nor destroyed although it may be converted from one form to another (law of conservation of energy).


Energy and Chemical Reactions

Example: Chemical reactions involve energy changes.

• forms of energy– eg., bonds within sugar molecules

• potential energy = stored energy – eg., molecular vibration measured as temperature

• kinetic energy = energy of motion

• If the chemical bonds within sugar are broken, energy of sugar can be used to heat the body or create movement.


Energy Transfer in Chemical Reactions

• An exergonic reaction is one in which the bond being broken has more energy than the one formed so that extra energy is released, usually as heat (occurs during catabolism of food molecules).

• An endergonic reaction is just the opposite and thus requires that energy be added, usually from a molecule called ATP, to form a bond, as in bonding amino acid molecules together to form proteins.


Energy Transfer in Chemical Reactions

• Reactions in living systems usually involve both kinds of reactions occurring together.– exergonic reactions release energy– endergonic reactions absorb energy

• You will learn of many examples in human metabolism that involve coupled exergonic and endergonic reactions; the energy released from one reaction will drive the other.– Glucose breakdown releases energy, which is

used to build ATP molecules (that store the energy for later use in other reactions.)


Activation Energy• Atoms, ions & molecules

are continuously moving& colliding.

• Activation energy is thecollision energy needed to break bonds & begin areaction.

• Increases in concentration &

temperature, increase the

probability of collision– more particles are in a given space when the concentration

is higher– particles move more rapidly when temperature is raised


Activation Energy

• Example:

Firewood does not spontaneously combust. Why?

Give an example of an oxidation reaction that has a relatively low activation energy.


Factors that influence the chance that a collision will occur and cause a chemical reaction include:

• Concentration• Temperature• Catalysts are chemical compounds that speed up chemical

reactions by lowering the activation energy needed for a reaction to occur (Figure 2.10).– A catalyst does not alter the difference in potential energy

between the reactants and products. It only lowers the amount of energy needed to get the reaction started.

– A catalyst helps to properly orient the colliding particles of matter so that a reaction can occur at a lower collision speed.

– The catalyst itself is unchanged at the end of the reaction; it is often re-used many times.


Effectiveness of Catalysts

• Catalysts speed up chemical reactions by lowering the activation energy.


Catalysts or Enzymes

Example:

• Normal body temperatures and concentrations are low enough that many chemical reactions are effectively blocked by the activation energy barrier.– Lactose typically reacts very slowly with water to break down

into two simple sugars called glucose and galactose.– Lactase, an enzyme (catalyst) orients the colliding particles

(lactose and water) properly so that they touch at the spots that make the reaction happen.

– Thousands of lactose/water reactions may be catalyzed by one lactase enzyme.

– Without lactase, the lactose will remain undigested in the intestines and often causes diarrhea and cramping.


Types of Chemical Reactions

• Synthesis reactions occur when two or more atoms, ions, or molecules combine to form new and larger molecules. These are anabolic reactions, meaning that bonds are formed. (Figure 2.8)

• In a decomposition reaction, a molecule is broken down into smaller parts. These are catabolic reactions, meaning that chemical bonds are broken in the process.

• Exchange reactions involve the replacement of one atom or atoms by another atom or atoms.

• In reversible reactions, end products can revert to the original combining molecules.


Synthesis Reactions--Anabolism

• Two or more atoms, ions or molecules combine to form new & larger molecules

• All the synthesis reactions in the body together are called anabolism

• Usually are endergonic because they absorb more energy than they release

• Example– combining amino acids to form a protein molecule


Decomposition Reactions--Catabolism

• Large molecules are split into smaller atoms, ions or molecules

• All decomposition reactions occurring together in the body are known as catabolism

• Usually are exergonic since they release more energy than they absorb


Exchange Reactions

• Substances exchange atoms– consist of both synthesis and decomposition

reactions• Example

– HCl + NaHCO3 gives rise to H2CO3 + NaCl– ions have been exchanged between substances


Reversible Reactions

• Chemical reactions can be reversible.– Reactants can become products or products can revert to

the original reactants• Indicated by the 2 arrows pointing in opposite directions

between the reactants and the products

AB A + B


Oxidation-Reduction Reactions

• Oxidation is the loss of electrons from a molecule (decreases its potential energy)– acceptor of the electron is often oxygen– commonly oxidation reactions involve removing a

hydrogen ion (H+) and a hydride ion (H-) from a molecule

– equivalent to removing 2 hydrogen atoms = 2H• Reduction is the gain of electrons by a molecule

– increases its potential energy• In the body, oxidation-reduction reactions are coupled &

occur simultaneously


INORGANIC COMPOUNDS AND SOLUTIONS

• Inorganic compounds usually lack carbon and are simple molecules; whereas organic compounds always contain carbon and hydrogen, usually contain oxygen, and always have covalent bonds.


Water

– Water is the most important and abundant inorganic compound in all living systems.

– An important property of water is its polarity, the uneven sharing of valence electrons that confers a partial negative charge near the one oxygen atom and two partial positive charges near the two hydrogen atoms in the water molecule (Figure 2.6).

– Water enables reactants to collide to form products


Water in Chemical Reactions

• Water is the ideal medium for most chemical reactions in the body and participates as a reactant or product in certain reactions.

• Hydrolysis breaks large molecules down into simpler ones by adding a molecule of water.

• Dehydration synthesis occurs when two simple molecules join together, eliminating a molecule of water in the process.


Water as a solvent

• In a solution the solvent dissolves the solute.• Substances which contain polar covalent bonds and

dissolve in water are hydrophilic, while substances which contain non polar covalent bonds are hydrophobic.

• The polarity of water and its bent shape allow it to interact with several neighboring ions or molecules. (Figure 2.11)

• Water’s role as a solvent makes it essential for health and survival.


Water as a Solvent• Most versatile solvent known

– polar covalent bonds (hydrophilic versus hydrophobic)

– its shape allows each watermolecule to interact with 4 ormore neighboring ions/molecules

• oxygen attracts sodium• hydrogen attracts chloride• sodium & chloride separate as ionic

bonds are broken• hydration spheres surround each ion and

decrease possibility of bonds being reformed

• Water dissolves or suspends many substances


High Heat Capacity of Water

• Water has a high heat capacity.• It can absorb or release a relatively large amount of heat

with only a modest change in its own temperature.• This property is due to the large number of hydrogen ions in

water.

• Heat of vaporization is also high– amount of heat needed to change from liquid to gas– evaporation of water from the skin removes large amount

of heat


Cohesion of Water Molecules

• Hydrogen bonds link neighboring water molecules giving water cohesion

• Creates high surface tension– difficult to break the surface of liquid if molecules

are more attracted to each other than to surrounding air molecules

– respiratory problem causes by water’s cohesive property

• air sacs of lungs are more difficult to inflate


Water as a Lubricant

• Water is a major part of mucus and other lubricating fluids.– mucus in respiratory and digestive systems– synovial fluid in joints– serous fluids in chest and abdominal cavities

• organs slide past one another • It is found wherever friction needs to be reduced or

eliminated


Solutions, Colloids, and Suspensions

• A mixture is a combination of elements or compounds that are physically blended together but are not bound by chemical bonds. Three common liquid mixtures are solutions, colloids, and suspensions.

• Solution: a substance called the solvent dissolves another substance called the solute. Usually there is more solvent than solute in a solution.

• A colloid differs from a solution mainly on the basis of the size of its particles with the particles in the colloid being large enough to scatter light.

• Suspension: the suspended material may mix with the liquid or suspending medium for some time, but it will eventually settle out.


Concentration

• The concentration of a molecule is a way of stating the amount of that molecule dissolved in solution (Table 2.3).

• Percent gives the relative mass of a solute found in a given volume of solution. (example: 100 ml of a 5% glucose solution contains 5 grams of glucose. A liter of the same solution would contain 50 grams of glucose.)

• A mole is the name for the number of atoms in an atomic weight of that element, or the number of molecules in a molecular weight of that type of molecule, with the molecular weight being the sum of all the atomic weights of the atoms that make up the molecule. (example: a mole of glucose has a mass of ___g, so a liter of 2 molar glucose would contain ____x2 grams of glucose.)


Inorganic Acids, Bases & Salts. Acids, bases and salts always dissociate into ions if they are

dissolved in water (Fig 2.12)– acids dissociate into H+and one or more anions

– bases dissociate into OH-and one or more cations

– salts dissociate into anions and cations, none of whichare either H+ or OH-

• Acid & bases react in the body to form salts• A salt, when dissolved in water, dissociates into cations and

anions, neither of which is H+ or OH- (Figure 2.12c). Many salts are present in the body and are formed when acids and bases react with each other.– Electrolytes are important salts in the body that carry

electric current (in nerve or muscle)


Concept of pH

• pH scale runs from 0 to 14 (concentration of H+ in moles/liter)• pH of 7 is neutral

(distilled water -- concentration of OH- and H+ are equal)• pH below 7 is acidic ([H+] > [OH-]).• pH above 7 is alkaline ([H+] < [OH-]).• pH is a logarithmic scale

Example: a change of two or three pH units

pH of 1 contains 10x10=100 more H+ than pH of 3

pH of 8 contains 10x10x10=1000 more H+ than pH of 11


Acid-Base Balance: The Concept of pH

• Biochemical reactions are very sensitive to even small changes in acidity or alkalinity.– pH of blood is 7.35 to 7.45

• A solution’s acidity or alkalinity is based on the pH scale

pH 0 (=100 = 1.0 moles H+/L)

pH 7.0 = 10-7 = 0.0000001 moles H+/L = neutrality or equal numbers of [H+] and [OH-].

pH 14 (= 10-14 = 0.00000000000001 moles H+/L)


Maintaining pH: Buffer Systems

• The pH values of different parts of the body are maintained fairly constant by buffer systems, which usually consist of a weak acid and a weak base.– convert strong acids or bases into weak acids or bases.– Example: carbonic acid-bicarbonate buffer system.

• Bicarbonate ions (HCO3-) act as weak bases and

carbonic acid (H2CO3) acts as a weak acid.• CO2 + H2O ↔ H2CO3 ↔ H+ + HCO3

-

• Table 2.4 shows pH values for certain body fluids compared to common substances.– gastric juice 1.2 to 3.0; saliva 6.35 to 6.85; bile 7.6 to 8.6

and blood 7.35 to 7.45


ORGANIC COMPOUNDS


Carbon and Its Functional Groups

• The carbon that organic compounds always contain has several properties that make it particularly useful to living organisms.

• It can react with one to several hundred other carbon atoms– forms large molecules of many different shapes.

• Many carbon compounds do not dissolve easily in water– useful materials for building body structures.

• Carbon compounds are mostly or entirely held together by covalent bonds and tend to decompose easily– organic compounds are a good source of energy.


• Many functional groups can attach to carbon skeleton– esters, amino, carboxyl, phosphate groups (Table 2.5)

• Very large molecules are called macromolecules (or “polymers” if all the monomer subunits are similar)

• Isomers have the same molecular formulas but different structures (glucose & fructose are both C6H12O6)

• STRUCTURALFORMULA OFGLUCOSE (Fig 2.14)

Carbon & Its Functional Groups


Carbohydrates

• Carbohydrates provide most of the energy needed for life and include sugars, starches, glycogen, and cellulose.

• Some carbohydrates are converted to other substances which are used to build structures and to generate ATP.

• Other carbohydrates function as food reserves.– About 2-3% of total human body weight

• Carbohydrates are divided into three major groups based on their size: monosaccharides, disaccharides, and polysaccharides (Table 2.6).


Carbohydrates

• formed from C, H, and O– ratio of one carbon atom for each water molecule

• glucose is 6 carbon atoms and 6 water molecules (H20)• source of energy for ATP formation• 2-3 % of total body weight

– glycogen is stored in liver and muscle tissue– sugar building blocks of DNA & RNA

(deoxyribose & ribose sugars)• Only plants produce starches or

cellulose for energy storage


Monosaccharides and Disaccharides: Sugars

• Monosaccharides contain from three to seven carbon atoms and include glucose, a hexose that is the main energy-supplying compound of the body.

• Humans absorb only 3 simple sugars without further digestion in our small intestine– glucose found syrup or honey– fructose found in fruit– galactose found in dairy products

• Disaccharides are formed from two monosaccharides by dehydration synthesis; they can be split back into simple sugars by hydrolysis (Figure 2.15). Glucose and fructose combine, for example, to produce sucrose.


Disaccharides

• Combining 2 monosaccharides by dehydration synthesis releases a water molecule.– sucrose = glucose & fructose– maltose = glucose & glucose– lactose = glucose & galactose (lactose intolerance)


Clinical Application:

• Lactose intolerance is a deficiency of the enzyme lactase. As a result undigested lactose remains in the feces and bacterial fermentation of lactose produces gas.


Polysaccharides

• Polysaccharides are the largest carbohydrates and may contain hundreds of monosaccharides.

• The principal polysaccharide in the human body is glycogen, which is stored in the liver or skeletal muscles. (Figure 2.16)– When blood sugar level drops, the liver hydrolyzes

glycogen to yield glucose which is released from the liver into the blood


Lipids

• Lipids, like carbohydrates, contain carbon, hydrogen, and oxygen; but unlike carbohydrates, they do not have a 2:1 ratio of hydrogen to oxygen.

• They have few polar covalent bonds – hydrophobic– mostly insoluble in polar solvents such as water– combines with proteins (lipoproteins) for transport in

blood

• Table 2.7 summarizes the various types of lipids and highlights their roles in the human body.


Lipids = fats

• Formed from C, H and O– fats– phospholipids– steroids– eicosanoids– lipoproteins – some vitamins

• 18-25% of body weight


Triglycerides

• Triglycerides are the most plentiful lipids in the body and provide protection, insulation, and energy (both immediate and stored).– At room temperature, triglycerides may be either solid

(fats) or liquid (oils).– Triglycerides provide more than twice as much energy

per gram as either carbohydrates or proteins.– Triglyceride storage is virtually unlimited.– Excess dietary carbohydrates, proteins, fats, and oils will

be deposited in adipose tissue as triglycerides.


Triglycerides

• Neutral fats composed of a single 3-carbon glycerol molecule and 3 fatty acid molecules (Figure 2.17).

• Very concentrated form of energy– 9 calories/gram compared to 4 for proteins &

carbohydrates– our bodies store triglycerides in fat cells if we eat extra

food


Triglycerides

• 3 fatty acids & one glycerol molecule• Fatty acids attached by dehydration systhesis


Saturation of Triglycerides

• Determined by the number of single or double covalent bonds

• Saturated fats contain single covalent bonds and are covered with hydrogen atoms----lard

• Monounsaturated are not completely covered with hydrogen----safflower oil, corn oil

• Polyunsaturated fats contain even less hydrogen atoms----olive and peanut oil


Clinical Application

• Essential fatty acids (EFA’s) are essential to human health and cannot be made by the human body. They must be obtained from foods or supplements.


Chemical Nature of Phospholipids

head tails


Phospholipids

• Phospholipids are important membrane components.• They are amphipathic, with both polar and nonpolar regions

(Figure 2.18).– a polar head

• a phosphate group (PO4-3) & glycerol molecule• forms hydrogen bonds with water

– 2 nonpolar fatty acid tails• interact only with lipids


Steroids

• Steroids have four rings of carbon atoms (Figure 2.19a).• Steroids include

– sex hormone– bile salts– some vitamins– cholesterol, with cholesterol serving as an important

component of cell membranes and as starting material for synthesizing other steroids.


Four Ring Structure of Steroids


Other Lipids

• Eicosanoids include prostaglandins and leukotrienes.– Lipid type derived from a fatty acid called arachidonic acid– prostaglandins = wide variety of functions

• modify responses to hormones• contribute to inflammatory response• prevent stomach ulcers• dilate airways• regulate body temperature• influence formation of blood clots

– leukotrienes = allergy & inflammatory responses

• Body lipids also include fatty acids; fat-soluble vitamins such as beta-carotenes, vitamins D, E, and K; and lipoproteins.


Proteins

• Proteins give structure to the body, regulate processes, provide protection, help muscles to contract, transport substances, and serve as enzymes (Table 2.8).

• Contain carbon, hydrogen, oxygen, and nitrogen• 12-18% of body weight


Proteins

• Constructed from combinations of 20 amino acids.– dipeptides formed from 2 amino acids

joined by a covalent bond called a peptide bond

– polypeptides chains formed from 10 to 2000 amino acids.


Amino Acid Structure

• Central carbon atom

• Amino group (NH2)

• Carboxyl group (COOH)• Side chains (R groups) vary

between amino acids


Levels of Structural Organization

• Levels of structural organization include – primary– secondary– tertiary– quaternary (Figure 2.22)

• The resulting shape of the protein greatly influences its ability to recognize and bind to other molecules.

• Denaturation of a protein by a hostile environment causes loss of its characteristic shape and function.


Formation of a Dipeptide Bond

• Dipeptides formed from 2 amino acids joined by a covalent bond called a peptide bond– dehydration synthesis

• Polypeptides chains contain 10 to 2000 amino acids.


Levels of Structural

Organization

• Primary is unique sequence of amino acids• Secondary is alpha helix or pleated sheet folding• Tertiary is 3-dimensional shape of polypeptide chain


Levels of Structural Organization

• Primary…• Secondary …• Tertiary…• Quaternary is relationship of

multiple polypeptide chains


Bonds of Tertiary & Quaternary Structure

• Disulfide bridges stabilize the tertiary structure of protein molecules

• Covalent bonds between sulfhydryl groups of 2 cysteine amino acids


Protein Denaturation

• The function of a protein depends on its ability to bind to another molecule

• Hostile environments such as heat, acid or salts will change a protein’s 3-D shape and destroy its ability to function– raw egg white when cooked is vastly different


Enzymes

• Catalysts in living cells are called enzymes.• Enzymes are highly specific in terms of the “substrate” with

which they react.

• Enzymes are subject to variety of cellular controls.• Enzymes speed up chemical reactions by increasing

frequency of collisions, lowering the activation energy and properly orienting the colliding molecules (Figure 2.23).


Enzymes

• Enzymes are protein molecules that act as catalysts• Enzyme = apoenzyme + cofactor

– Apoenzymes are the protein portion– Cofactors are nonprotein portion

• may be metal ion (iron, zinc, magnesium or calcium)

• may be organic molecule derived from a vitamin• Enzymes usually end in suffix -ase and are named for

the types of chemical reactions they catalyze– oxidase, kinase, and lipase are examples


Enzyme Functionality

• Highly specific• Very efficient

– speed up reaction up to 10 billion times faster

• Under nuclear control– rate of synthesis of enzyme– inhibitory substances– inactive forms of enzyme


Clinical Application

• inherited disorder: galactosemia– Infant lacks enzyme.– Galactose accumulates in the blood causing anorexia.– Treatment is elimination of milk from the diet.


Nucleic Acids: Deoxyribonucleic Acid (DNA) and Ribonucleic Acid (RNA)

• Nucleic acids are huge organic molecules that contain carbon, hydrogen, oxygen, nitrogen, and phosphorus.

• Deoxyribonucleic acid (DNA) forms the genetic code inside each cell and thereby regulates most of the activities that take place in our cells throughout a lifetime.

• Ribonucleic acid (RNA) relays instructions from the genes in the cell’s nucleus to guide each cell’s assembly of amino acids into proteins by the ribosomes.

• The basic units of nucleic acids are nucleotides, composed of a nitrogenous base, a pentose, sugar, and a phosphate group (Figures 2.24a, b).


DNA Structure

• Contains C, H, O, N and phosphorus.

• Each gene of our genetic material is a piece of DNA that controls the synthesis of a specific protein.

• A molecule of DNA is a chain of nucleotides.

• A nucleotide includes:– nitrogenous base (A-G-T-C)– pentose sugar– phosphate group


DNA Fingerprinting

• Used to identify criminal, victim or a child’s parents– need only strand of hair, drop of semen or spot of blood

• Certain DNA segments are repeated several times– unique from person to person


RNA Structure

• Differs from DNA– single stranded– ribose sugar not deoxyribose sugar– uracil nitrogenous base replaces thymine

• Types of RNA within the cell, each with a specific function– messenger RNA– ribosomal RNA– transfer RNA


Adenosine Triphosphate (ATP)

• Temporary molecular storage of energy as it is being transferred from exergonic catabolic reactions to cellular activities– muscle contraction, transport of substances across cell

membranes, movement of structures within cells and movement of organelles

• Consists of 3 phosphategroups attached toadenine & 5-carbonsugar (ribose)

• (Figure 2.25).


Formation & Usage of ATP

• Hydrolysis of ATP (removal of terminal phosphate group by enzyme -- ATPase)– releases energy– leaves ADP (adenosine diphosphate)

• Synthesis of ATP – enzyme ATP synthase catalyzes the addition of

the terminal phosphate group to ADP– energy from 1 glucose molecule is used during

both anaerobic and aerobic respiration to create 36 to 38 molecules of ATP


end

Chapter 2

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