Exam 1

Molecules vs compounds
2. Trace elements of the body
3. Covalent vs Ionic vs Hydrogen bonds
4. Solution vs suspension vs emulsion vs colloid
5. Anabolic vs Catabolic reactions
6. Functional groups of amino acids
7. Most abundant elements of the body
8. Atomic number and atomic mass
9. Atomic structure
10. What atomic particle determines chemical properties; bonding capabilities
11. What happens to salt when it is put in water
12. Anion vs Cation
13. Properties of water
14. What is # of molecules per volume?
15. pH scale
16. How your body maintains constant blood pH
17. Types of reactions
18. Exergonic vs endergonic
19. Factors that increase or decrease the rate of reaction
20. Functional groups
21. What makes a compound organic?
22. Hydrolysis vs Dehydration Synthesis
23. Elements found in each macromolecules
24. Examples of macromolecules
25. Structure of macromolecules
26. Hydrophobic vs hydrophilic
27. Functions of macromolecules
28. Denaturation
29. Levels of structure found in protein
30. DNA and RNA
31. Reactants vs products of a reaction

32. Why is carbon so versatile?
33. Function of minerals in the body
34. Function of a buffer system
35. Examples of buffer systems
36. Bicarbonate buffer system reaction
37. Buffer in blood
38. How the kidneys help regulate pH
39. metabolic alkalosis; metabolic acidosis, urinary alkalosis, urinary acidosis, respiratory acidosis; diabetic

acidosis

Biochemistry
• The study of the molecules that compose living organisms
• Carbohydrates, fats, proteins, and nucleic acids
• Useful for understanding cellular structures, basic
physiology, nutrition, and health

Expected Learning Outcomes:
• Identify the elements of the body from their symbols.
• Distinguish between elements and compounds.
• State the functions of minerals in the body.
• Explain the basis for radioactivity and the types and
hazards of ionizing radiation.
• Distinguish between ions, electrolytes, and free radials.
• Define the types of chemical bonds.

The Chemical Elements
A chemical element is the simplest form of matter to have unique
chemical properties
• Each identified by an atomic number—number of protons in the
nucleus
• Periodic table arranges elements (represented as 1-2 letter
symbols) by atomic number
• 91 naturally occurring elements
• 24 play roles in humans; 6 are most abundant (98.5% body weight)
• Oxygen, carbon, hydrogen, nitrogen, calcium, and phosphorus
• Trace elements present in minute amounts, but play vital roles
• Some are minerals—inorganic elements extracted from soil by
plant, passed up food chain to humans
• 4% of body weight, mostly calcium and phosphorus
• Body structure (bones, teeth), enzyme function, nerve/muscle cell functions

2.1b Atomic Structure
Greek philosopher coined the term atom (“indivisible”) as the
smallest unit of matter
Neils Bohr proposed planetary model of atomic structure in 1913;
useful as a schematic, but not accurate structure
• Nucleus—center of atom, composed of protons and neutrons
• Protons+
(p ) : single (+) charge; mass = 1 atomic mass unit (amu)
• Neutrons0
(n ) : no charge; mass = 1 amu
• Atomic mass is approx equal to total number of protons and neutrons
• Electrons-
(e ) in concentric clouds (electron shells or energy levels)
surrounding nucleus
• Have a single (−) charge, very low mass
• An atom is electrically neutral, as number of electrons equals number of
protons
• Valence electrons in the outermost shell and determine chemical
bonding properties of an atom

2.1c Isotopes and Radioactivity 1
Isotopes are varieties of an element that differ only in the
number of neutrons and therefore in atomic mass
• Extra neutrons increase atomic weight
• Isotopes of an element are chemically similar because
they have the same number of valence electrons
• Atomic weight (relative atomic mass) of an element
accounts for the fact that an element is a mixture of
isotopes

Isotopes and Radioactivity 2
Isotopes of an element have identical chemical behavior, but can
differ in their physical behavior
• Radioisotopes—unstable isotopes that decay and give off
radiation in a process called radioactivity
• Every element has at least one radioisotope
• Intense radiation can be ionizing (ionizing radiation)—ejects
electrons, destroys molecules, creates free radicals—and can
cause genetic mutations and cancer
• Examples: UV radiation, X-rays, alpha particles, beta particles, gamma
rays
• Physical half-life of radioisotopes—time required for 50% to
decay to a stable state
• Biological half-life of radioisotopes—time required for 50% to
disappear from the body

Isotopes and Radioactivity 3
Standard measure of radiation dosage is the sievert (Sv)
• 5 Sv or more is usually fatal
• Standard acceptable exposure = 50 mSv per year
Background radiation
• Natural sources such as radon gas and cosmic rays
• Average = 2.4 mSv per year
Artificial sources of radiation
• X-rays, color TVs, and so on
• Average = 0.6 mSV per year

Radiation and Madame Curie
First woman to receive Nobel Prize
(1903)
First woman in world to receive a
PhD
• Coined term radioactivity
• Discovered radioactivity of
polonium and radium
• Trained physicians in use of X-
rays and pioneered radiation
therapy as cancer treatment
Died of radiation poisoning at age 67

2.1d Ions, Electrolytes, and Free Radicals 1
An ion is a charged particle (atom or molecule) with unequal
number of protons and electrons
• Ionization—transfer of electrons from one atom to another
• Anion—particle that has a net negative charge due to gain
of electrons
• Cation—particle that has a net positive charge due to loss
of electrons
• Ions with opposite charges are attracted to each other
Ions, Electrolytes, and Free Radicals 2
Ions (continued)
• Salts—electrically neutral compounds of cations and
anions; readily dissociate in water into ions and act as
electrolytes
• Examples: Sodium chloride, calcium chloride
• Electrolytes—substances that ionize in water and form
solutions capable of conducting electric current
• Functions of electrolytes:
• Chemical reactivity, osmotic effects, electrical excitability of nerve
and muscle
• Electrolyte balance is one of the most important considerations in
patient care (imbalances can lead to coma or cardiac arrest)

Ions, Electrolytes, and Free Radicals 3
Ions (continued)
• Free radicals—unstable, highly reactive particles with an
unusual number of electrons
• Produced by normal metabolic reactions, radiation, certain
chemicals
• Trigger reactions that destroy molecules, and can cause cancer,
death of heart tissue, and aging
• Example: superoxide anion,2O-
• Antioxidants—chemicals that neutralize free radicals
• Example: superoxide dismutase (SOD) is an antioxidant enzyme
that converts superoxide anion into oxygen and hydrogen peroxide
• Selenium, vitamin E, vitamin C, and carotenoids are antioxidants
obtained through the diet
2.1e Molecules and Chemical Bonds 1
Atoms can combine to form molecules
• Molecule—particle composed of two or more atoms
united by a chemical bond
• Compound—molecule composed of two or more different elements
• Can be represented by a molecular formula, which identifies
constituent elements and how many atoms of each are present
• Also can be represented by a structural formula, which identifies the
location of each atom
• Isomers—molecules with identical molecular formulae but different
arrangements of their atoms

Molecules and Chemical Bonds 2
• The molecular weight (MW) of a compound is the sum of
the atomic weights of its atoms.
• Example: to calculate the MW of glucose (C6H12O6):
66 C atoms × 12 amu each = 72 amu
12 H atoms × 1 amu each = 12 amu
6 O atoms × 16 amu each = 96 amu
Molecular weight (MW) = 180 amu

Molecules and Chemical Bonds 3
Chemical bonds hold atoms together within a molecule, or attract
one molecule to another
• Ionic bonds—attraction of a cation to an anion
• Example: sodium and chloride ions bond to form sodium chloride
• Relatively easily broken by something more attractive, such as water
• Covalent bonds—atoms share one or more pairs of electrons
• Single covalent bond: nuclei share 1 pair of electrons
• Double covalent bond: nuclei share 2 pairs of electrons
• If electrons are shared equally, it’s a nonpolar covalent bond; example:
carbon atoms bonding together
• If electrons shared unequally, it’s a polar covalent bond; example: hydrogen
bonding with oxygen, electrons spend more time by oxygen

Molecules and Chemical Bonds 4
A hydrogen bond is a weak attraction between a slightly
positive hydrogen atom in one molecule and a slightly
negative oxygen or nitrogen atom in another atom
• Relatively weak bonds, but very important to physiology
• Water molecules are attracted to each other by hydrogen bonds.
• Large molecules (DNA and proteins) are shaped in part by the
formation of hydrogen bonds within them.

Molecules and Chemical Bonds 5

Van der Waals forces are weak, brief attractions between

neutral atoms

• Fluctuation in electron density within an atom creates polarity

for a moment, and attracts adjacent atom for a very short time

• Only 1% as strong as a covalent bond, but play important role

in physiology (for example, protein folding)

2.2 Water and Mixtures
Expected Learning Outcomes:
• Define mixture and distinguish between mixtures and
compounds.
• Describe the biologically important properties of water.
• Show how three kinds of mixtures differ from each other.
• Define acid and base and interpret the pH scale.
• Discuss some ways in which the concentration of a
solution can be expressed, and the kinds of information
we can derive from the different units of measure.

2.2 Introduction
• Body fluids are complex mixtures of chemicals
• Mixtures—consist of substances that are physically
blended but not chemically combined; each substance
retains its own chemical properties

2.2a Water 1
Most mixtures in our bodies consist of chemicals dissolved or
suspended in water
• Water is 50 to 75% of body weight
• Polar covalent bonds and a V-shape give water a set of
properties that account for its ability to support life
• Solvency
• Cohesion
• Adhesion
• Chemical reactivity
• Thermal stability

Water 2
Properties of water:
• Solvency—ability to dissolve other chemicals
• Water is the universal solvent because it dissolves more
substances than any other solvent
• Metabolic reactions depend on solvency of water
• Hydrophilic substances dissolve in water; are polarized or charged
• Hydrophobic substances do not dissolve in water; are nonpolar or
neutral
• To be soluble in water, molecule must be polarized or
charged
• Example: attractions to water overpower ionic bond in NaCl
• Water forms hydration spheres around each ion and the salt
dissolves; water’s negative pole faces+
Na , its positive pole faces Cl^-

Water 3
Properties of water (continued):
• Adhesion—tendency of one substance to cling to another
• Water adheres to membranes reducing friction around organs
• Cohesion—tendency of molecules of the same substance
to cling to each other
• Water is very cohesive due to its hydrogen bonds
• Surface film on surface of water is due to molecules being held
together by surface tension
• Chemical reactivity—ability to participate in chemical
reactions
• Water ionizes into+ -
H and OH ; ionizes other chemicals (acids,
salts); involved in hydrolysis and dehydration synthesis reactions

Water 4
Properties of water (continued):
• Thermal stability due to high heat capacity—amount of
heat needed to raise the temperature of 1g of a substance
by 1°C
• Calorie (cal) is the base unit of heat; 1 cal is the amount of heat to
raise the temperature of 1g of water by 1°C
• Water stabilizes internal temperature; hydrogen bonds resist
temperature increases by inhibiting molecular motion

2.2b Solutions, Colloids, and Suspensions 1
Mixtures of other substance in water classified as solutions,
colloids, and suspensions
• Solution—consists of particles called the solute mixed
with a more abundant substance (usually water) called the
solvent
• Solute can be gas, solid, or liquid
• Solutions are defined by the following properties:
• Solute particles under 1 nm
• Solute particles do not scatter light
• Will pass through most membranes
• Will not separate on standing

Solutions, Colloids, and Suspensions 2
Mixtures (continued)
• Colloids
• Colloids in the body are often mixtures of protein and water
• Many can change from liquid to gel state within and between cells
• Colloids are defined by the following physical properties:
• Particles range from 1–100 nm in size
• Scatter light and are usually cloudy
• Particles too large to pass through semipermeable membrane
• Particles remain permanently mixed with the solvent when mixture
stands

Solutions, Colloids, and Suspensions 3
Mixtures (continued)
• Suspension
• Defined by the following physical properties:
• Particles exceed 100 nm
• Too large to penetrate selectively permeable membranes
• Cloudy or opaque in appearance
• Separates on standing
• Example: blood cells in blood plasma
• Emulsion—suspension of one liquid in another
• Examples: oil-and-vinegar salad dressing; fat in breast milk

2.2c Acids, Bases, and pH 1
Substances can be acids or bases depending on their
tendency to release or bind+
H ions
• An acid is a proton donor; releases+
H ions in water
• A base is a proton acceptor; accepts-+ ions or releaseH s OH
ions in water
Acidity is measured by pH scale
• Derived from the molarity of+
H
• pH of 7.0 is neutral=+ -
(H OH )
• pH of less than 7 is acidic+ -
(H OH )
• pH of greater than 7 is basic- +
(OH H )
• Maintaining normal (slightly basic) pH of blood is crucial
for physiological functions
• Buffers are chemical solutions that resist changes in pH

Acids, Bases, and pH 2
pH (continued)
• pH is the negative logarithm of hydrogen ion molarity
• ph = -log [H+]
+
log H
• Example: if [H+] = 10^-3, then pH = -log [10^-3] =3
• A change of one number on the pH scale represents a
tenfold change in+
H concentration
• pH 4.0 is 10 times as acidic as pH 5.0

2.2d Other Measures of Concentration 1
Solutions measured in terms of their concentration of solute;
several different expressions of concentration
• Weight per volume
• Weight of solute in a given volume of solution
• Example: IV saline contains 8.5 g NaCl per liter of solution.
• Common biology units: milligrams per deciliter (mg/dl)
• Example: Serum cholesterol may be 200 mg/dl
• Percentage
• Might be weight of solute (solid) per volume
• Example: 5% dextrose solution has 5 g solute in 100 ml solution.
• Might be volume of solute (liquid) per volume of solution
• Example: 70% ethanol has 70 ml of ethanol in 100 ml solution.

Other Measures of Concentration 2
Expressions of concentration (continued)
• Molarity (M)
• Number of moles of solute per liter of solution
• One mole is the number of grams equal to its molecular weight
• Often the most physiological meaningful measure of concentration
• Body fluids usually quantified in millimolar (mM) concentrations

Other Measures of Concentration 3
Expressions of concentration (continued)
• Milliequivalents per liter (mEq/L)
• Expression of electrolyte concentration
• Accounts for millimolar concentration of solute and electrical charge
of its particles
• Important for nerve firing, the heartbeat, muscle contractions, and
deliver of intravenous fluids

2.3 Energy and Chemical Reactions
Expected Learning Outcomes
• Define energy and work, and describe some types of
energy.
• Understand how chemical reactions are symbolized by
chemical equations.
• List and define the fundamental types of chemical
reactions.
• Identify the factors that govern the speed and direction of
a reaction.
• Define metabolism and its two subdivisions.
• Define oxidation and reduction, and relate these to
changes in the energy content of a molecule.

2.3a Energy and Work
Energy is the capacity to do work
• To do work means to move something, such as a muscle
or a molecule
• Potential energy—energy stored in an object, but not
currently doing work; example: water behind a dam
• Chemical energy—potential energy in molecular bonds
• Free energy—potential energy available in a system to do work
• Kinetic energy—energy of motion, energy doing work
• Examples: muscle movements, flow of ions, vibration of eardrum
• Heat—kinetic energy of molecular motion
• Electromagnetic energy—kinetic energy of photons
• Electrical energy has both potential and kinetic forms

2.3b Classes of Chemical Reactions 1
A chemical reaction is a process in which a covalent or
ionic bond is formed or broken
• A chemical equation symbolizes the course of a chemical
reaction
• Reactants (on left) → products (on right)
• Classes of chemical reactions include:
• Decomposition reactions
• Synthesis reactions
• Exchange reactions

Classes of Chemical Reactions 2
Classes of chemical reactions (continued)
• Decomposition reactions
• Large molecule breaks down into two or more smaller ones
• AB → A + B
• Synthesis reactions
• Two or more small molecules combine to form a larger one
• A + B → AB
• Exchange reactions
• Two molecules exchange atoms or group of atoms
• AB + CD → AC + BD
• Example: stomach acid (HCl) and sodium bicarbonate (NaHCO3)
from the pancreas combine to form NaCl and H2CO3

Classes of Chemical Reactions 3
• Reversible reactions can proceed in either direction
under different circumstances
• Symbolized with double-headed arrow
• Example:+   +3- +
2 2 2 3CO H O H CO HCO H
• An important reaction in respiratory, urinary, and digestive
physiology
• Reversible reactions follow the law of mass action
• Direction of reaction determined by relative abundance of
substances on either side of equation
• Equilibrium is reached when ratio of products to reactants is stable

2.3c Reaction Rates
Reactions occur when molecules collide with enough force
and correct orientation
Reaction rates increase when:
• Concentration of reactants increases
• Temperature rises
• A catalyst is present
• Enzyme catalysts bind to reactants and hold them in orientations that
facilitate the reaction
• Catalysts are not changed by the reaction and can repeat the process
frequently

2.3d Metabolism, Oxidation, and Reduction 1
Metabolism—all chemical reactions of the body; two
divisions:
• Catabolism
• Energy-releasing (exergonic) decomposition reactions
• Breaks covalent bonds
• Produces smaller molecules
• Anabolism
• Energy-storing (endergonic) synthesis reactions
• Requires energy input
• Example: production of protein or fat
• Catabolism and anabolism are inseparably linked
• Anabolism is driven by energy released by catabolism

Metabolism, Oxidation, and Reduction 2
Metabolism (continued)
• Oxidation
• A chemical reaction in which a molecule gives up electrons and
releases energy
• Molecule is oxidized when it loses electrons
• The oxidizing agent is the electron acceptor (often oxygen)
• Reduction
• Any chemical reaction in which a molecule gains electrons and
energy
• Molecule is reduced when it accepts electrons
• The reducing agent is the molecule that donates electrons
• Oxidation of one molecule is always accompanied by
reduction of another; called oxidation-reduction (redox)
reactions

2.4 Organic Compounds 1
Expected Learning Outcomes:
• Explain why carbon is especially well suited to serve as
the structural foundation of many biological molecules.
• Identify some common functional groups of organic
molecules from their formulae.
• Discuss the relevance of polymers to biology and explain
how they are formed and broken by dehydration synthesis
and hydrolysis.

2.4 Organic Compounds 2
Expected Learning Outcomes (continued):
• Discuss the types and functions of carbohydrates.
• Discuss the types and functions of lipids.
• Discuss protein structure and function.
• Explain how enzymes function.
• Describe the structure, production, and function of ATP.
• Identify other nucleotide types and their functions
• Identify the principal types of nucleic acids

2.4a Carbon Compounds and Functional Groups 1
Organic chemistry is the study of compounds containing
carbon
• Four categories:
• Carbohydrates
• Lipids
• Proteins
• Nucleic acids

Carbon Compounds and Functional Groups 2
Carbon uniquely suited to form a variety of structures
• Carbon:
• Has four valence electrons; can form four covalent bonds with other
atoms
• Can bind readily with other carbon atoms to form carbon
backbones—long chains, branched molecules, rings
• Readily bonds with hydrogen, oxygen, nitrogen, sulfur, and other
elements
• Carbon backbones carry a variety of functional groups
• Small clusters of atoms attached to carbon backbone
• Determine many of the properties of organic molecules
• Examples: hydroxyl, methyl, carboxyl, amino, phosphate

2.4b Monomers and Polymers 1
Macromolecules are large organic molecules with high
molecular weights
• Most macromolecules are polymers
• Molecules made of a repetitive series of identical or similar subunits
called monomers
• Monomers may be identical or different
• Examples: starch is a polymer of about 3,000 identical glucose
monomers; DNA is a polymer of 4 different nucleotide monomers
• Polymers formed by polymerization—the joining of
monomers

Monomers and Polymers 2
Polymerization (continued)
• Monomers covalently linked together by dehydration
synthesis (condensation) reactions
• A hydroxyl (-OH) group is removed from one monomer, and a
hydrogen (-H) from another; water produced as a by-product
• Hydrolysis is the opposite of dehydration synthesis
• Splitting a polymer in monomers by the addition of water
• Enzyme helps break the covalent bond that links two monomers
together
• A water molecule ionizes into− +
OH and H
•-
OH is added to one monomer
•+
H is added to the other monomer

2.4c Carbohydrates 1
Carbohydrates are hydrophilic organic molecules
• General formula:2(CH O) ,n n = number of carbon atoms
• Glucose, n = 6, so formula is6 12 6C H O
• 2:1 ratio of hydrogen to oxygen
• Names of carbohydrates often built from the root “sacchar-”
and the suffix “-ose” both meaning sugar, sweet
• Examples: sugars and starches

Carbohydrates 2
Monosaccharides are the simplest carbohydrates
• Monomers of larger carbohydrates
• Three important monomers are glucose, galactose, and
fructose
• Produced by digestion of more complex carbohydrates
• Glucose is blood sugar
• All three have

the same molecular formula: C6H12O6
• Isomers of each other
• Ribose and deoxyribose are also monomers
• Part of RNA and DNA, respectively

Carbohydrates 3
Disaccharides are sugars made of two covalently bonded
monosaccharides
• Three important disaccharides:
• Sucrose (table sugar)—glucose + fructose
• Lactose (milk sugar)—glucose + galactose
• Maltose (sugar in grain products)—glucose + glucose
Oligosaccharides are short chains of 3 or more
monosaccharides

The Three Major Disaccharides (Sucrose)

The Three Major Disaccharides (Lactose, Maltose)

Carbohydrates 4
Polysaccharides are long chains of monosaccharides (~50
or more, up to thousands)
• Three important polysaccharides:
• Glycogen—energy storage in cells of liver, muscle, brain, uterus,
vagina
• Starch—energy storage in plants that is digestible by humans
• Cellulose—structural molecule in plants that is important for human
dietary fiber (but indigestible to us)

Carbohydrates 5
Functions of carbohydrates:
• Quickly mobilized source of energy
• All digested carbohydrates converted to glucose
• Oxidized to make ATP
• Often conjugated (bound) to lipids and proteins
• Example: lipids, proteins of cell membrane have chains of up to 12 sugars
attached to form glycolipids and glycoproteins, respectively
• Glycoproteins are a major component of mucus
• Proteoglycans—macromolecules that are more carbohydrate
than protein
• Form gels that hold cells and tissues together; fill umbilical cord and eye
• Joint lubrication; responsible for the rubbery texture of cartilage
• Moiety—each component of a conjugated macromolecule

2.4d Lipids 1
Lipids are hydrophobic organic molecules with a high ratio of
hydrogen to oxygen
• More calories per gram than carbohydrates
• Five primary types of lipids in the human body:
• Fatty acids
• Triglycerides
• Phospholipids
• Eicosanoids
• Steroids

Lipids 2
Types of lipids (continued)
• Fatty acids—chains of 4–24 carbon atoms with carboxyl
group on one end and methyl group on the other
• Essential fatty acids must be obtained from food
• Fatty acids are classified as saturated or unsaturated
• Saturated fatty acid—carbon atoms linked by single covalent
bonds
• Molecule contains as much hydrogen as possible (“saturated” with
hydrogen)
• Unsaturated fatty acids—contain some double bonds between
carbons
• Molecule has potential to add hydrogen
• Polyunsaturated fatty acids have multiple double bonds between
carbons

Lipids 3
Types of lipids (continued)
• Triglycerides—three fatty acids linked to glycerol
• Formed by dehydration synthesis; broken down by hydrolysis
• Primary function is energy storage; also help with insulation and
shock absorption (adipose tissue)
• Also called neutral fats because once formed, fatty acid is no
longer acidic
• Dietary oils and fats are triglycerides
• Oils are usually liquid at room or body temperature
• Example: plant-derived polyunsaturated triglycerides (polyunsaturated
fats) such as corn and olive oils
• Saturated fats are solid at room or body temperature
• Example: animal-derived saturated triglycerides (for example, animal
fat)

Trans Fats and Cardiovascular Health
A trans fat is a triglyceride with one or more trans-fatty acids
• Trans-fatty acids—two covalent single C–C bonds angle in
opposite directions (trans means “across from”) on each
side of the C=C double bond
• Carbon chains are straighter than cis-fatty acids; pack more
densely and are solid at room temp
• Abundant in partially hydrogenated oil (PHO), sold as
vegetable shortening; popular for baked goods
• Resists enzymatic breakdown in the human body, remain
in circulation longer, deposits in the arteries; thus, raises
the risk of heart disease

Lipids 4
Types of lipids (continued)
• Phospholipids—similar to triglycerides, but one fatty acid
is replaced by a phosphate group
• Phospholipids are amphipathic
• Fatty acid “tails” are hydrophobic
• Phosphate “head” is hydrophilic
• Phosphate group linked to other functional groups
• Structural foundation of cell membrane

Lipids 5
Types of lipids (continued)
• Eicosanoids—20-carbon compounds derived from a fatty acid called
arachidonic acid
• Hormone-like chemical signals between cells
• Includes prostaglandins
• Function in inflammation, blood clotting, hormone action, labor
contractions, blood vessel diameter
Prostaglandin

Lipids 6
Types of lipids (continued)
• Steroids—lipid with 17 carbon atoms in four rings
• Cholesterol is the “parent” steroid from which other steroids are
synthesized
• Important for nervous system function and structural integrity of all
cell membranes
• 15% of our cholesterol comes from diet
• 85% is internally synthesized (mostly in liver)
• Other steroids include cortisol, progesterone, estrogens,
testosterone, and bile acids

“Good” and “Bad” Cholesterol
There is only one kind of cholesterol
“Good” and “bad” cholesterol refer to droplets of lipoprotein in
the blood that are complexes of cholesterol, fat,
phospholipid, and protein
HDL (high-density lipoprotein) = “good cholesterol”
• Lower ratio of lipid to protein
• May help to prevent cardiovascular disease
LDL (low-density lipoprotein) = “bad cholesterol”
• High ratio of lipid to protein
• Contributes to cardiovascular disease

2.4e Proteins 1
A protein is a polymer of amino acids
• Amino acids have a central carbon with three attachments
• Amino group (-NH2)
• Carboxyl group (–COOH)
• R (radical) group
• 20 amino acids used to make the proteins are identical except for the
radical (R) group
• Properties of each amino acid determined by the R group

Proteins 2
A peptide is composed of two or more amino acids joined by
peptide bonds
• Peptide bond
• Joins amino group of one amino acid to carboxyl group of the next
• Formed by dehydration synthesis
• Peptides are named for the number of amino acids they
contain
• Dipeptides (2 amino acids)
• Tripeptides (3 amino acids)
• Oligopeptides (fewer than 10 to15 amino acids)
• Polypeptides (larger than 15 amino acids)

Protein Structure 1
Proteins have a complex three-dimensional shape referred to
as their conformation
• Unique; crucial to function
• Proteins can reversibly change conformation to affect function
• Important examples seen in muscle contraction, enzyme catalysis,
membrane channel opening, and so on
• Denaturation—extreme conformational change that
destroys function
• Extreme heat or pH may cause permanent (irreversible)
denaturation
• Example: cooked egg white becomes opaque and stiff

Protein Structure 2
Proteins have three to four levels of complexity:
• Primary structure
• Sequence of amino acids within protein molecule
• Primary structure is encoded by genes
• Secondary structure
• Coiled or folded shape held together by hydrogen bonds
• Hydrogen bonds between slightly negative C=O and slightly
positive –NH groups
• Most common secondary structures:
• Alpha helix has a springlike shape
• Beta sheet (beta-pleated sheet) has a folded, ribbonlike shape

Protein Structure 3
Protein levels of complexity (continued)
• Tertiary structure
• Further bending and folding of proteins into globular and fibrous
shapes due to hydrophobic–hydrophilic interactions and van der
Waals forces
• Disulfide bridges between cysteine amino acids stabilize tertiary
structure
• Globular proteins
• Compact tertiary structure for proteins within cell membrane and
proteins that move freely in body fluids
• Fibrous proteins
• Slender filaments suited for roles in muscle contraction and
strengthening of skin and hair

Protein Structure 4
Protein levels of complexity (continued)
• Quaternary structure
• Associations of two or more polypeptide chains due to ionic bonds
and hydrophobic–hydrophilic interactions
• Occurs only in some proteins
• Example: hemoglobin has four peptide subunits

Protein Structure 5
• Conjugated proteins contain a non-amino acid moiety
called a prosthetic group covalently bound to them
• Example: hemoglobin contains four complex iron-
containing rings called a heme moiety (see previous slide)

Protein Functions 1
Proteins have more diverse functions than other
macromolecules
• Structure
• Keratin—tough structural protein of hair, nails, skin surface
• Collagen—contained in deeper layers of skin, bones, cartilage, and
teeth
• Communication
• Neurotransmitters, some hormones, and other signaling molecules
are proteins
• Signaling molecules that exert their effects by reversibly binding to a
receptor molecule are called ligands
• The receptors to which the signaling molecules bind are also
proteins

Protein Functions 2
Protein functions (continued)
• Membrane transport
• Channels allow hydrophilic substances to diffuse across cell
membranes
• Carriers help solutes cross cell membranes via active or passive
transport
• Catalysis
• The enzymes that catalyze physiological reactions are usually
globular proteins
• Recognition and protection
• Glycoproteins are important for immune recognition
• Antibodies are proteins

Protein Functions 3
Protein functions (continued)
• Movement
• Molecular motors (motor proteins) are molecules with the ability to
change shape repeatedly
• Cell adhesion
• Proteins bind cells together

2.4f Enzymes and Metabolism
Enzymes are proteins that function as biological catalysts
• Some are ribozymes, composed of RNA and found in
ribosomes
• Enzymes act on one or more substrates
• Speed up chemical reaction by lowering the activation
energy—the energy needed to get a reaction started
• Permit reactions to occur rapidly at body temperature
• Enzyme naming convention:
• Named for substrate with -ase as the suffix
• Examples: amylase catalyzes the hydrolysis of amylose (starch);
lactase catalyzes the hydrolysis of lactose (milk sugar)

Enzyme Structure and Action 1
Enzyme action:
1. Substrate binds to pocket on enzyme called the active
site
2. Formation of enzyme–substrate complex
• Enzyme–substrate specificity is like a lock and key
3. Enzyme releases reaction products
• Enzyme unchanged and can repeat process
• Example: the substrate sucrose is hydrolyzed by sucrase
into the reaction products glucose and fructose

Enzyme Structure and Action 2
Temperature, pH and other factors can change enzyme
shape and function
• Can alter ability of enzyme to bind to substrate
• Enzymes vary in optimum pH
• Salivary amylase works best at pH 7.0
• Pepsin in stomach works best at pH 2.0
• Temperature optimum for human enzymes is usually near
body temperature (37°C)

Cofactors
Many human enzymes require a nonprotein partner called a
cofactor
• Cofactors may be inorganic or organic
• Inorganic cofactors include the ions iron, copper, zinc, magnesium,
and calcium
• Some of these work by binding to the enzyme, triggering a
conformational change that activates the active site
• Organic cofactors are called coenzymes
• Often derived from vitamins
• Example:NAD+ derived from niacin; acts as an electron shuttle between
the metabolic pathways glycolysis and aerobic respiration

Metabolic Pathways 1
A metabolic pathway is a chain of reactions, each catalyzed
by a different enzyme
• A simple metabolic pathway may be symbolized as
follows:A B C D
α β γ
→ → →
• A is the initial reactant, B and C are intermediates, and D is the end
product
• The Greek letters (α, β, and γ) above the reaction arrows represent
the enzymes that catalyze each step
• A is the substrate for enzyme α
• B is the substrate for enzyme β
• C is the substrate for enzyme γ

Metabolic Pathways 2
• Metabolic pathways are turned on or off by altering
enzyme activity
• This may be achieved many ways including:
• Binding or dissociation of cofactors
• The end product inhibits an enzyme at an earlier step (for example,
product D binds to enzyme a and shuts down production of
intermediate product B)
• Pathway is turned on when end product is in demand (low
concentration) and turned off once concentration increases

2.4g ATP, Other Nucleotides, and Nucleic Acids
Nucleotides—organic compounds with three components:
• Nitrogenous base (single or double carbon–nitrogen ring)
• Sugar (monosaccharide)
• One or more phosphate groups
Example:
• ATP (adenosine triphosphate)—has adenine nitrogenous
base, a ribose sugar, and three phosphate group

Adenosine Triphosphate 1
Adenosine triphosphate (ATP) is the body’s most important
energy-transfer molecule
• Stores energy gained from exergonic reactions
• Releases it within seconds for physiological work
• Holds energy in covalent bonds between phosphates
• Second and third phosphate groups have high energy bonds(~)
• Most energy transfers to and from ATP involve adding or removing
the third phosphate group

Adenosine Triphosphate 2
ATP (continued):
• Hydrolysis of ATP is catalyzed by adenosine
triphosphatases (ATPases)
• Breaks the third high-energy phosphate bond of ATP to produce
adenosine diphosphate (ADP), inorganic phosphatei(P ), and
energy
•iATP ADP P energy→ + +
• Releases 7.3 kcal of energy per mole (505g) of ATP
• Phosphorylation—addition of free phosphate group
(released from ATP hydrolysis) to another molecule to
activate it
• Carried out by enzymes called kinases

Adenosine Triphosphate 3
Energy for ATP synthesis comes primarily from glucose
oxidation
• First stage is glycolysis—splitting glucose into two
pyruvate molecules
• A little ATP made in this process; most chemical energy remains in
the pyruvate
• Fate of pyruvate depends on oxygen supply
• If ATP demand outpaces oxygen supply, pyruvate undergoes
anaerobic fermentation to lactate
• If enough oxygen is available, aerobic respiration occurs in
mitochondria

Other Nucleotides
Guanosine triphosphate (GTP)
• Another nucleotide involved in energy transfer
• In some reactions, donates a phosphate group
Cyclic adenosine monophosphate (cAMP)
• Formed by removal of second and third phosphate groups
from ATP
• In many cases, it’s formation is triggered by a chemical
signal (example: hormone) binding to cell surface
• cAMP becomes “second messenger” within cell

Nucleic Acids
Nucleic acids are polymers of nucleotides
• DNA (deoxyribonucleic acid)
• Contains millions of nucleotides
• Constitutes genes, the instructions for synthesizing proteins
• RNA (ribonucleic acid)
• 70 to 10,000 nucleotides long
• Carries out genetic instruction (encoded in DNA) for synthesizing
proteins
• Assembles amino acids in right order to produce proteins