Download Chapter two powerpoint File - District 196 e

2
Chemical Level
of Organization
PowerPoint® Lecture Presentations prepared by
Alexander G. Cheroske
Mesa Community College at Red Mountain
© 2011 Pearson Education, Inc.
Section 1: Atoms and Molecules
• Learning Outcomes
• 2.1 Describe an atom and how atomic structure
affects the mass number and atomic weight
of the various chemical elements.
• 2.2 Explain the relationship between electrons
and energy levels.
• 2.3 Compare the ways in which atoms combine
to form molecules and compounds.
• 2.4 Describe the three states of matter and the
importance of hydrogen bonds in liquid
water.
© 2011 Pearson Education, Inc.
Section 1: Atoms and Molecules
• Studying the human body begins at the
chemical level of organization
• Chemistry studies the structure of matter
• Matter (anything that takes up space and has
mass)
• Mass (the amount of material in matter)
• On Earth, mass is the same as weight
© 2011 Pearson Education, Inc.
Section 1: Atoms and Molecules
• Atoms (smallest stable units of matter)
• Composed of subatomic particles
• Protons (p+, have a positive electrical charge)
• Neutrons (n or n0, are electrically neutral or
uncharged)
• Electrons (e–, much smaller and have a
negative electrical charge)
© 2011 Pearson Education, Inc.
An atom and its important subatomic particles
Protons (p+) have a positive electrical charge.
Neutrons (n or n0) are electrically neutral,
which means they are uncharged.
Electrons (e-) are much smaller and about
1/1800th the mass of either protons or neutrons.
They bear a negative electrical charge.
Figure 2 Section 1
© 2011 Pearson Education, Inc.
1
Section 1: Atoms and Molecules
• Atoms (continued)
• Can be subdivided into:
• Nucleus
• Center of atom
• Contains protons and neutrons
• Electron cloud
• Electrons whirl around the nucleus creating a
cloud
© 2011 Pearson Education, Inc.
The nucleus and electron cloud of
an atom
The nucleus of an atom lies at its center.
The nucleus contains one or more
protons and it may contain neutrons as
well. The mass of the atom is primarily
determined by the numbers of protons
and neutrons in the nucleus.
The electrons in the atom whirl around
the nucleus, creating an electron cloud.
Figure 2 Section 1
© 2011 Pearson Education, Inc.
2
Section 1: Atoms and Molecules
• Molecule (atoms interact to produce larger,
more complex structures)
• Everything is made up of varying combinations
of atoms
• Unique characteristics of an object result from
the particular atoms and how they interact
• Mass of any object represents the sum of
masses of its atoms
© 2011 Pearson Education, Inc.
A molecule, which consists
of two or more atoms
Figure 2 Section 1
© 2011 Pearson Education, Inc.
3
Module 2.1: Atoms
• Atoms
• Normally contain equal numbers of protons and
electrons
• Atomic number (number of protons)
• Mass number (number of protons + number of
neutrons)
• Element (pure substance of only atoms with
the same atomic number)
© 2011 Pearson Education, Inc.
Module 2.1: Atoms
• Electron cloud
• Electrons whirl around the nucleus
• Are attracted to positively charged protons in
nucleus
• Example of an electrical force
• Electron shell
• Two-dimensional representation of the electron
cloud
• Outermost shell represents the surface of the
electron cloud
© 2011 Pearson Education, Inc.
Module 2.1: Atoms
• Isotopes
• Atoms of a single element but with different
numbers of neutrons
• Distinguishable based on mass
© 2011 Pearson Education, Inc.
Electron shell models of three isotopes of hydrogen
Electron shell
Hydrogen-1
mass number: 1
Hydrogen-2,
deuterium
mass number: 2
Hydrogen-3,
tritium
mass number: 3
Figure 2.1
© 2011 Pearson Education, Inc.
2
Module 2.1: Atoms
• Atomic weight (actual mass of an atom)
• Expressed in daltons (also known as the
atomic mass unit or amu)
• One dalton is very close to the weight of one
proton or one neutron
• Is an average mass number reflecting the
proportion of different isotopes
• Example: hydrogen
• Atomic number = 1
• Atomic weight = 1.0079
© 2011 Pearson Education, Inc.
A representation of an element’s atomic weight, which
reflects the proportions of the various isotopes of that
element
Atomic weight of a mixture of
Atomic weight of hydrogen-1 = 1 hydrogen isotopes = 1.0079
Figure 2.1
© 2011 Pearson Education, Inc.
3
Module 2.1: Atoms
• Principal elements
• Thirteen most abundant elements to total body weight
• Trace elements
• Fourteen other elements in the body in very small
amounts
• Each element has a chemical symbol
• Most are connected with English names
• Example: O for oxygen
• A few are from other languages
• Example: Na for sodium, from Latin natrium
© 2011 Pearson Education, Inc.
Figure 2.1
© 2011 Pearson Education, Inc.
4
Figure 2.1
© 2011 Pearson Education, Inc.
4
Figure 2.1
© 2011 Pearson Education, Inc.
4
Figure 2.1
© 2011 Pearson Education, Inc.
4
Module 2.1 Review
a. Define atom.
b. Describe trace elements.
c. How is it possible for two samples of
hydrogen to contain the same number of
atoms yet have different weights?
© 2011 Pearson Education, Inc.
Module 2.2: Electrons
•
Electrons occupy a series of concentric energy
levels or electron shells
•
Reactive elements have unfilled outer energy
levels
•
Will react with other elements to fill outer shell
•
Once outer shell is full, the atom is stable and
unreactive
•
Examples: hydrogen and helium
•
Inert elements have full outermost energy levels
•
Do not participate in chemical processes
•
Examples: helium and neon
© 2011 Pearson Education, Inc.
Module 2.2: Electrons
•
Energy levels (electron shells)
•
First level: holds 2 electrons
•
Second level: holds 8 electrons
•
Third level: holds 8 electrons
© 2011 Pearson Education, Inc.
Reactive elements
Two examples of reactive elements, which have
unfilled outermost energy levels
Inert elements
Two examples of inert elements, which do not
react in chemical processes because their
outermost energy levels are filled
The first energy
level can hold a
maximum of two
electrons.
Hydrogen has
one electron in
the first energy
level.
Helium has two
electrons in the
first energy level.
Hydrogen, H
Atomic number: 1
Mass number: 1
1 electron
Lithium has one
electron in the
second energy
level; it is
extremely
reactive.
The second and
third energy levels
can each contain
up to eight
electrons.
Lithium, Li
Atomic number: 3
Mass number: 6
(3 protons + 3 neutrons)
3 electrons
© 2011 Pearson Education, Inc.
Helium, He
Atomic number: 2
Mass number: 4
(2 protons + 2 neutrons)
2 electrons
Neon, Ne
Atomic number: 10
Mass number: 20
(10 protons + 10 neutrons)
10 electrons
Neon has eight
electrons in the
Second energy
level; it does not
React with other
atoms.
Figure 2.2
1
-
2
Module 2.2: Electrons
•
Reactive elements achieve stability by gaining,
losing, or sharing electrons
Cation – Loss of electron, giving atom a positive
charge
•
•
Loss of 1 electron gives the ion a +1 charge
•
Some ions carry charges of +2, +3, or +4
Anion – Gain of electron, giving atom a negative
charge
•
•
•
Gain of 1 electron gives the ion a –1 charge
•
Some ions carry charges of –2, –3, or –4
Stabilizing interactions often form chemical bonds
© 2011 Pearson Education, Inc.
The loss of an electron from a sodium atom, forming a sodium cation
Sodium atom, Na (reactive)
Sodium ion, Na+ (stable)
Figure 2.2
© 2011 Pearson Education, Inc.
3
The gain of an electron by a chlorine atom, forming a chloride anion
Chlorine atom, Cl (reactive)
Chloride ion, Cl– (stable)
Figure 2.2
© 2011 Pearson Education, Inc.
4
Module 2.2 Review
a. Indicate the maximum number of electrons
that can occupy each of the first three
electron shells (energy levels) of an atom.
b. Explain why the atoms of inert elements do
not react with one another or combine with
atoms of other elements.
c. Explain how cations and anions form.
© 2011 Pearson Education, Inc.
Module 2.3: Chemical bonds
•
Chemical bonding creates
•
Molecules (substances with two or more
atoms)
•
Compounds (substances with two or more
different atoms)
© 2011 Pearson Education, Inc.
Module 2.3: Chemical bonds
•
Ionic bonds
•
Created by electrical attraction between
cations and anions
•
Involves the transfer of electrons from one
atom to another to achieve stability
© 2011 Pearson Education, Inc.
Step 1: Formation of sodium and chloride ions. The
sodium atom loses an electron to the chloride atom. This
produces two stable ions with filled outer energy levels.
Step 2: Formation of an ionic bond. Because these ions form
close together, and have opposite charges, they are attracted
to one another. This creates NaCl, an ionic compound.
Sodium atom
Sodium ion (Na+)
Sodium chloride (NaCl)
Chlorine atom
Chloride ion
(CI–)
The formation of the compound sodium
chloride (table salt), involving the formation of
ions and ionic bonding
Chloride ions (CI–)
Sodium ions (Na+)
A crystal of sodium chloride
© 2011 Pearson Education, Inc.
Figure 2.3
1
-
2
Module 2.3: Chemical bonds
•
Covalent bonds
•
Sharing of electrons between atoms
•
Single covalent bond
•
•
One electron contributed by each atom
Double covalent bond
•
Two electrons contributed by each atom
© 2011 Pearson Education, Inc.
Molecules, which consist of one or more elements bound by covalent bonds
resulting from the sharing of electrons
Molecule
Description
Hydrogen (H2)
Hydrogen atoms aren’t found as individuals—they exist as molecules, each
containing a pair of hydrogen atoms. The two atoms share their electrons to fill
their outer energy levels, and the electron pair orbits both nuclei. One electron
is contributed by each atom, so this is called a single covalent bond.
Oxygen (O2)
Carbon
dioxide
(CO2)
An oxygen atom has 6 electrons in its outer energy level. By forming a
double covalent bond with another oxygen atom, an oxygen molecule is
created with a stable outer energy level.
A carbon atom has 4 electrons in its outer energy level, so it needs to gain 4
from other atoms to achieve stability. In a molecule of carbon dioxide, a carbon
atom shares a pair of electrons with each of two oxygen atoms and forms two
double covalent bonds.
Figure 2.3
© 2011 Pearson Education, Inc.
3
Module 2.3: Chemical bonds
•
Nonpolar molecule
•
Electrons shared equally between atoms
•
No electrical charge on the molecule
© 2011 Pearson Education, Inc.
Space-filling models of
two nonpolar molecules
Oxygen (O2)
Carbon dioxide (CO2)
Figure 2.3
© 2011 Pearson Education, Inc.
4
Module 2.3: Chemical bonds
•
Polar molecule
•
Unequal sharing of electrons between atoms
•
Formed by polar covalent bonds
•
Example: water molecule (H2O)
•
Oxygen exerts a stronger attraction for the
electrons than the hydrogens
•
Oxygen carries a slightly negative charge
•
Hydrogens carry a slightly positive charge
© 2011 Pearson Education, Inc.
A water molecule, which is a polar molecule (has
positively and negatively charged portions) that contains
polar covalent bonds
Hydrogen
atom
Positive pole
Oxygen
atom
Negative pole
Figure 2.3
© 2011 Pearson Education, Inc.
6
Module 2.3 Review
a. Name and distinguish between the two most
common types of chemical bonds.
b. Describe the kind of bonds that hold the
atoms in a water molecule together.
c. Relate why we can apply the term molecule to
the smallest particle of water but not to that of
table salt.
© 2011 Pearson Education, Inc.
Module 2.4: Three states of matter
•
Matter exists in one of three states:
1. Solid (particles held tightly together)
•
Maintain their volume and shape at ordinary
temperatures and pressures
2. Liquid (particles held less tightly together)
•
•
Has a constant volume
Container determines shape
3. Gas (particles independent of each other)
•
•
•
Has neither a constant volume nor a fixed shape
Can be compressed or expanded
Will fill a container of any size
© 2011 Pearson Education, Inc.
Figure 2.4
© 2011 Pearson Education, Inc.
1
-
3
Module 2.4: Three states of matter
•
Water
•
Only substance that exists in all three states at
temperatures compatible with life
•
Solid (ice)
•
Liquid (water)
•
•
Exists over a broad range of temperatures due
to interactions among the polar water molecules
Gas (water vapor)
© 2011 Pearson Education, Inc.
Figure 2.4
© 2011 Pearson Education, Inc.
4
Module 2.4: Three states of matter
•
Hydrogen bond
•
Attraction of the small polar charges on
hydrogen atoms to negative charges on other
polar molecules
•
Can change shape of molecules or pull
molecules together
•
Produces surface tension
•
Slows the rate of evaporation at water surface
© 2011 Pearson Education, Inc.
The attraction of hydrogen bonds, which occur
between the small positive and negative charges
on different water molecules
KEY
Hydrogen
Oxygen
Hydrogen
bond
Figure 2.4
© 2011 Pearson Education, Inc.
5
Figure 2.4
© 2011 Pearson Education, Inc.
6
Module 2.4: Three states of matter
•
Water can dissolve a variety of inorganic
compounds due to its polar molecules
•
Seawater contains almost all naturally
occurring elements
•
Our body fluids contain at least 29 elements
© 2011 Pearson Education, Inc.
Module 2.4 Review
a. Describe the different states of matter in
terms of shape and volume.
b. By what means are water molecules attracted
to each other?
c. Explain why small insects can walk on the
surface of a pond, and tears protect the
surface of the eye from dust particles.
© 2011 Pearson Education, Inc.
Section 2: Chemical Reactions
• Learning Outcomes
• 2.5 Use chemical notation to symbolize
chemical reactions.
• 2.6 Distinguish among the major types of
chemical reactions that are important for
studying physiology.
• 2.7 Describe the crucial role of enzymes in
metabolism.
© 2011 Pearson Education, Inc.
Section 2: Chemical Reactions
• Cells remain alive by controlling chemical
reactions
• New bonds may form between atoms
• Existing bonds may be broken
• Reactants (reacting substances) form different
products
• Metabolism
• All the cellular reactions of the body at a given
moment
© 2011 Pearson Education, Inc.
Section 2: Chemical Reactions
• Each cell is a “chemical factory” where
chemical reactions control:
• Growth
• Maintenance
• Repair
• Secretion
• Contraction
• Energy transfer and use
© 2011 Pearson Education, Inc.
Figure 2 Section 2
© 2011 Pearson Education, Inc.
1
Section 2: Chemical Reactions
• Work
• Movement of an object or change in physical
structure of matter
• Examples:
• Walking or running
• Synthesis of molecules
• Conversion of liquid water to vapor (evaporation)
© 2011 Pearson Education, Inc.
Section 2: Chemical Reactions
• Energy (capacity to perform work)
• Kinetic energy
• Energy of motion
• Example: skeletal muscle contracting
• Potential energy
• Stored energy (potential to do work)
• Example: skeletal muscle at rest
• Conversion of energy is never 100% efficient
• Some energy is released as heat
• Example: body temperature rises as muscles contract
© 2011 Pearson Education, Inc.
Module 2.5: Chemical notation
• Chemical notation
• A simple “chemical shorthand” to describe
chemical structures and events
© 2011 Pearson Education, Inc.
Module 2.5: Chemical notation
• Symbol of element indicates one atom of that
element
• H = one atom of hydrogen
• O = one atom of oxygen
• Number preceding symbol indicates more than
one atom of that element
• 2 H = two atoms of hydrogen
• 2 O = two atoms of oxygen
© 2011 Pearson Education, Inc.
Module 2.5: Chemical notation
• Subscript following symbol indicates a
molecule with that number of atoms of that
element
• H2 = hydrogen molecule, composed of two
hydrogen atoms
• O2 = oxygen molecule, composed of two
oxygen atoms
• H2O = water molecules, composed of two
hydrogen atoms and one oxygen atom
© 2011 Pearson Education, Inc.
Module 2.5: Chemical notation
• Chemical reaction
• Reactants
• Participants at reaction start
• Usually on the left
• Products
• Generated at end of reaction
• Usually on the right
• 2 H + O  H 2O
© 2011 Pearson Education, Inc.
Module 2.5: Chemical notation
• A superscript plus or minus following symbol
indicates an ion
• Single plus sign indicates cation with +1
• Atom has lost one electron
• Example: Na+
• Single minus sign indicates anion with -1
• Atom has gained one electron
• Example: Cl-
• Number before sign indicates more than one
electron lost or gained
• Example Ca2+
© 2011 Pearson Education, Inc.
Module 2.5: Chemical notation
• Chemical reactions rearrange atoms into new
combinations
• Numbers of atoms must be the same one both
sides of equation (= balanced)
• Unbalanced: H2 + O2  H2O
• Balanced: 2 H2 + O2  2 H2O
© 2011 Pearson Education, Inc.
Module 2.5: Chemical notation
• Mole (mol)
• Quantity (in grams) equal to an element’s atomic
weight
• One mole of a given element always contains the
same number of atoms of a mole of another element
• Weight differs
• Example:
• 1 mole of oxygen
• 16 grams
• 1 mole of hydrogen
• 1 gram
© 2011 Pearson Education, Inc.
A mole, a quantity of an element with a weight
(in grams) equal to that element’s atomic weight
1 mole of oxygen
1 mole of hydrogen
Figure 2.5
© 2011 Pearson Education, Inc.
2
Module 2.5: Chemical notation
• Molecular weight
• Sum of the atomic weights of all atoms
comprising a molecule
• Can be used to calculate quantities of reactants
or products for a chemical reaction
© 2011 Pearson Education, Inc.
Atomic weight of H = 1
Molecular weight of H2 = 2
Two moles of H2 weigh 4 g
Atomic weight of O = 16
Molecular weight of O2 = 32
One mole of O2 weighs 32 g
Molecular weight of 2 H2O:
2 x (2 + 16)
2 x 18 = 36
Two moles of H2O weigh 36 g
The use of molecular weights to determine the amounts of
reactants needed and the amount of product generated
Figure 2.5
© 2011 Pearson Education, Inc.
4
Module 2.5 Review
a. The chemical shorthand used to describe chemical
compounds and reactions effectively is known as
__________________.
b. Using the rules for chemical notation,write the
molecular formula for glucose, a compound
composed of 6 carbon atoms, 12 hydrogen atoms,
and 6 oxygen atoms.
c. Calculate the weight of one mole of glucose. (The
atomic weight of carbon = 12 grams.)
© 2011 Pearson Education, Inc.
Module 2.6: Types of chemical reactions
•
Types of chemical reactions
1. Decomposition
•
Breaks molecule into smaller components
•
•
•
AB  A + B
Example: digestion of food for absorption
Hydrolysis
•
Decomposition using the insertion of water
• AB + H2O  A-H + OH-B
•
Catabolism (katabole, a throwing down)
•
•
Collective decomposition reactions of the body
Releases kinetic energy
• CD  C + D + ENERGY
© 2011 Pearson Education, Inc.
Module 2.6: Types of chemical reactions
•
Types of chemical reactions (continued)
2. Synthesis
•
Combining of atoms or smaller molecules into larger
molecules
•
•
•
Always involves new bond formation
Dehydration synthesis (condensation)
•
•
•
A + B  AB
Removal of water to form a larger molecule
A-H + OH-B  A-B + H2O
Anabolism (anabole, a throwing upward)
•
•
Collective synthesis of new molecules in the body
Requires energy
© 2011 Pearson Education, Inc.
Module 2.6: Types of chemical reactions
•
Many important biological reactions are
reversible
•
A + B ↔ AB
•
Equilibrium
•
Rates at which the two reactions occur are in
balance
© 2011 Pearson Education, Inc.
Module 2.6: Types of chemical reactions
•
Types of chemical reactions (continued)
3. Exchange reactions
•
Reactants are “shuffled” to produce new
products
•
•
AB + CD  AD + CB
Reactants must break apart (decomposition) in
order to form products (synthesis)
© 2011 Pearson Education, Inc.
Module 2.6 Review
a. Identify and describe three types of chemical
reactions important in human physiology.
b. Distinguish the roles of water in hydrolysis and
dehydration synthesis reactions.
c. In cells, glucose, a six-carbon molecule, is
converted into two three-carbon molecules by a
reaction that releases energy. What is the source
of the energy?
© 2011 Pearson Education, Inc.
Module 2.7: Enzymes
•
Most chemical reactions require energy to
activate reactants
•
Activation energy
•
•
Amount of energy to start a reaction
Enzymes
•
Special proteins that lower the activation
energy for a reaction
© 2011 Pearson Education, Inc.
Module 2.7: Enzymes
•
Catalysts (katalysis, dissolution)
•
Compounds that accelerate reactions without
changing themselves
•
Reactions continue until equilibrium is reached
•
Enzymes perform as catalysts
© 2011 Pearson Education, Inc.
The action of protein catalysts called enzymes in promoting a chemical reaction by lowering its
activation energy
Activation
energy
Progress of reaction
Specific enzymes lower the
activation energy requirements
so that important reactions will
occur.
Energy
Energy
In the external environment, extreme conditions
can provide the activation energy. For example,
complex sugars can the broken down in a
laboratory by boiling them in an acidic solution.
Activation
energy
Progress of reaction
Figure 2.7
© 2011 Pearson Education, Inc.
1
Module 2.7: Enzymes
•
Metabolic pathway
•
Series of reactions (each involving an enzyme)
that support life
•
May absorb or release energy on completion
•
•
Exergonic (exo-, outside)
•
Release energy
•
Common in the body and help to maintain body
temperature
Endergonic (endo-, inside)
•
Absorb energy
© 2011 Pearson Education, Inc.
A metabolic pathway, a series of interlocking steps, each controlled by a specific enzyme
Figure 2.7
© 2011 Pearson Education, Inc.
2
Module 2.7: Enzymes
•
Metabolites (metabole, change)
•
All molecules that can be synthesized or
decomposed in our bodies
Processed by enzymatic reactions
Nutrients
•
•
•
•
Essential metabolites normally obtained from our diet
Can be classified as:
•
•
Organic
•
Always contain carbon and hydrogen
•
Examples: sugars, fats, proteins
Inorganic
•
Generally do not contain carbon and hydrogen
•
Examples: carbon dioxide, water, salts
© 2011 Pearson Education, Inc.
Module 2.7 Review
a. What is an enzyme?
b. Explain the differences between metabolites
and nutrients.
c. Why are enzymes needed in our cells?
© 2011 Pearson Education, Inc.
Section 3: The Importance of Water in the Body
• Learning Outcomes
• 2.8 Explain how the chemical properties of
water affect the solubility of inorganic and
organic molecules.
• 2.9 Discuss the importance of pH and the role
of buffers in body fluids.
© 2011 Pearson Education, Inc.
Section 3: The Importance of Water in the Body
• Water
• Most important constituent of body
• Accounts for up to 2/3 of total body weight
• Changes in body water content can affect all
systems
© 2011 Pearson Education, Inc.
Important Properties of Water
Lubrication
Water is an effective lubricant because
there is little friction between water
molecules. Thus even a thin layer of
water between two opposing surfaces
will greatly reduce friction between
them; water reduces friction within
joints and in body cavities.
Reactivity
In our bodies, chemical reactions
occur in water, and water molecules
are also participants in some
reactions, including hydrolysis and
dehydration synthesis.
Dehydration
synthesis
Hydrolysis
High heat capacity
Heat capacity is the ability to absorb
and retain heat. Water has an
unusually high heat capacity, because
water molecules in the liquid state are
attracted to one
another through
hydrogen
bonding.
• The temperature of water must be
high before individual molecules
have enough energy to break free to
become water vapor, a gas.
• Water carries a great deal of heat
away with it when it finally does
change from a liquid to a gas. This
feature accounts for the cooling
effect of perspiration on the skin.
• A large mass of water changes
temperature very slowly. This
property is called thermal inertia.
Solubility
A remarkable number of inorganic and
organic molecules will dissolve in water.
The individual particles become
dispersed within the water, and the result
is a solution—a uniform mixture of two
or more substances. The medium in
which other atoms, ions, or molecules are
dispersed is called the solvent; the
dispersed substances are the solutes. In
aqueous solutions, water is the solvent.
Figure 2 Section 3
© 2011 Pearson Education, Inc.
Module 2.8: Physiological systems depend
on water
• Many inorganic compounds have ionic bonds
• Ionization or dissociation
• Ionic bonds broken in water due to the polarity
of the water molecule
• Anions surrounded by positive poles of water
molecules
• Cations surrounded by negative poles of water
molecules
• Hydration sphere
• Sheath of water molecules around an ion
© 2011 Pearson Education, Inc.
Module 2.8: Physiological systems depend
on water
• Organic molecules may dissolve in water if
they contain polar covalent bonds
• = Hydrophilic (hydro-, water + philos, loving)
© 2011 Pearson Education, Inc.
The roles of hydration spheres in the ionization or dissociation of an ionic compound (center) and in solutions of
an organic molecule containing polar covalent bonds (right)
Sodium chloride crystal
Negative
pole
Hydration
spheres
H
Positive
pole
Glucose
molecule
Cl–
Na+
NaCl in solution
Glucose molecule in solution
Figure 2.1
© 2011 Pearson Education, Inc.
1
-
3
Module 2.8: Physiological systems depend
on water
• Electrolytes
• Ions in solution that conduct electric current
• Example: NaCl  Na+ + Cl–
• Movement of ions across the plasma
membrane of cells is important in many body
cells
• Muscle
• Nervous
• Ion concentrations in bodily fluids is carefully
regulated
© 2011 Pearson Education, Inc.
The movement of sodium and chloride ions,
dissociated from the electrolyte sodium
chloride, when subjected to an electric field
Cl–
Na+
Figure 2.8
© 2011 Pearson Education, Inc.
4
Figure 2.8
© 2011 Pearson Education, Inc.
5
Fats and oils
Protein
Many organic molecules either lack
polar covalent bonds or have very
few. Such molecules do not have
positive and negative poles and are
said to be nonpolar. When nonpolar
molecules are exposed to water,
hydration spheres do not form and
the molecules do not dissolve.
Molecules that do not readily interact
with water are called hydrophobic
(hi-drō-FŌ-bik; hydro-, water +
phobos, fear). Among the most
familiar hydrophobic molecules are
fats and oils of all kinds.
Body fluids typically contain large
and complex organic molecules,
such as proteins, that are held in
solution by their association with
water molecules.
A colloid, a solution containing dispersed proteins or other large molecules
Figure 2.8
© 2011 Pearson Education, Inc.
6
Module 2.8: Physiological systems depend
on water
• Many organic molecules lack or have very few
polar covalent bonds (nonpolar)
• = Hydrophobic (hydro-, water + phobos, fear)
• Do not readily interact with water
• Examples: fats and oils
© 2011 Pearson Education, Inc.
Module 2.8: Physiological systems depend
on water
• Large organic molecules are usually held in
solution by water molecules
• Example: proteins
• Colloid
• Large molecules that stay in solution indefinitely
• Example: liquid Jell-O
• Suspension
• Larger molecules that may settle out of solution
due to gravity
• Example: blood
© 2011 Pearson Education, Inc.
Module 2.8 Review
a. Define electrolytes.
b. Distinguish between hydrophilic and
hydrophobic molecules.
c. Explain how the ionic compound sodium
chloride dissolves in water.
© 2011 Pearson Education, Inc.
Module 2.9: pH
• Hydrogen ion (H+)
• Hydrogen atom that has lost electron
• Extremely reactive in solution
• Concentration in body regulated precisely
• Hydroxide ion (OH–)
• Produced when water decomposes (along with H+)
© 2011 Pearson Education, Inc.
The dissociation of a water molecule into
a hydrogen ion, H+, and a hydroxide ion,
OH–
Figure 2.9
© 2011 Pearson Education, Inc.
1
Module 2.9: pH
• pH
• Negative logarithm of the hydrogen ion
concentration (mol/L)
• Ranges from 0–14
• Acidic: below 7
• Neutral: equal to 7
• Alkaline: above 7
© 2011 Pearson Education, Inc.
The pH—the negative logarithm of the
hydrogen ion concentration—of a variety
of common substances
Blood
The pH of blood normally ranges form 7.35 to 7.45. Abnormal fluctuations
in pH can damage cells and tissues by breaking chemical bonds, changing
the shapes of proteins, and altering cellular functions. Acidosis is an
abnormal physiological state caused by low blood pH (below 7.35); a pH
below 7 can produce coma. Alkalosis results from an abnormally high pH
(above 7.45); a blood pH above 7.8 generally causes uncontrollable and
sustained skeletal muscle contractions.
1 mol/L
hydrochloric
acid
Urine
Beer,
vinegar,
wine, Tomatoes,
pickles grapes
Stomach
hydrochloric
acid
Extremely
acidic
pH 0
[H+] 100
(mol/L)*
1 mol/L
sodium
hydroxide
1
10-1
Saliva,
milk
Pure
water
Increasing concentration of H+ Neutral
2
10-2
3
10-3
4
10-4
A solution with a pH below 7 is
acidic (a-SI-dik), meaning that
it contains more hydrogen
ions than hydroxide ions.
5
10-5
6
10-6
7
10-7
Seawater
Eggs
Household
bleach
Household
ammonia
Increasing concentration of OH–
8
10-8
A solution with a pH of 7 is
said to be neutral, because
it contains equal numbers of
hydrogen and hydroxide ions.
9
10-9
10
10-10
11
10-11
Extremely
basic
12
10-12
13
10-13
14
10-14
A pH above 7 is basic, or
alkaline (AL-kuh-lin), meaning
that it has more hydroxide ions
than hydrogen ions.
* One liter of pure water contains about 0.0000001 mol of hydrogen ions and an equal number of
hydroxide ions. In other words, the concentration of hydrogen ions in a solution of pure water
is 0.0000001 mol per liter. This can be written as [H+] = 10–7 mol/L. The brackets around the H+
signify “the concentration of,” another example of chemical notation.
© 2011 Pearson Education, Inc.
Oven
cleaner
Figure 2.9
2
Module 2.9: pH
• Blood pH
• Normally ranges from 7.35–7.45
• Values above or below can alter cell functions
• Acidosis: below 7.35
• Alkalosis: above 7.45
© 2011 Pearson Education, Inc.
Module 2.9: pH
• Acid
• Solute that dissociates and releases hydrogen
ions
• Often referred to as proton donors
• After losing an electron, a hydrogen ion consists
solely as a proton
• Strong acids dissociate completely
• Example: hydrochloric acid (HCl)
• HCl  H+ + Cl–
© 2011 Pearson Education, Inc.
Hydrochloric acid (HCl), a strong acid because it releases hydrogen ions
Figure 2 .9 3
© 2011 Pearson Education, Inc.
Module 2.9: pH
• Base
• Solute that removes hydrogen ions from
solution
• Proton acceptor
• Strong bases dissociate completely
• Example: sodium hydroxide (NaOH)
• NaOH  Na+ + OH–
© 2011 Pearson Education, Inc.
Sodium hydroxide (NaOH), a strong base because it releases hydroxide
ions that remove hydrogen ions from solution
Figure 2 .9 4
© 2011 Pearson Education, Inc.
Module 2.9: pH
• Weak acids and bases fail to dissociate
completely
• Example: carbonic acid (H2CO3)
• H2CO3 ↔ H+ + HCO3–
• In body fluids, carbonic acid reversibly
dissociates to hydrogen ion and bicarbonate
ion
© 2011 Pearson Education, Inc.
Carbonic acid, a weak acid in body fluids that reversibly dissociates into
hydrogen ion and a bicarbonate ion
Figure 2.9
© 2011 Pearson Education, Inc.
5
Module 2.9: pH
• Salt
• Ionic compound with any cation and anion
except:
• H+
• OH–
• Example: NaCl
• NaCl  Na+ + Cl–
© 2011 Pearson Education, Inc.
NaCl, a salt—an ionic compound consisting of any cation except a hydrogen
ion and any anion except a hydroxide ion.
Figure 2.9
© 2011 Pearson Education, Inc.
6
Module 2.9: pH
• Buffers
• Compounds that stabilize pH by removing or
replacing hydrogen ions
• Help to maintain normal pH of body fluids
• Buffer systems
• Usually involve a weak acid and its related salt
• Examples:
• Carbonic acid (H2CO3)
• Sodium bicarbonate (NaHCO3)
© 2011 Pearson Education, Inc.
Module 2.9 Review
a. Define pH.
b. Explain the differences among an acid, a
base, and a salt.
c. What is the significance of pH in physiological
systems?
© 2011 Pearson Education, Inc.
Section 4: Organic Compounds
• Learning Outcomes
• 2.10 Discuss the structures and functions of
carbohydrates.
• 2.11 Discuss the structures and functions of
lipids.
• 2.12 Discuss the structures and diverse
functions of eicosanoids, steroids,
phospholipids, and glycolipids.
• 2.13 Discuss protein structure and the
essential functions of proteins within the
body.
© 2011 Pearson Education, Inc.
Section 4: Organic Compounds
• Learning Outcomes
• 2.14 Explain how enzymes function within the
body.
• 2.15 Discuss the structure and function of highenergy compounds.
• 2.16 Compare and contrast the structures and
functions of DNA and RNA.
© 2011 Pearson Education, Inc.
Section 4: Organic Compounds
• Organic compounds
• Always contain carbon and hydrogen and generally
oxygen
• Many are long chains of carbon linked with covalent
bonds
• Many are soluble in water
• Functional groups
• Attached groupings of atoms that occur commonly in
many organic molecules
• Influence the properties of the overall molecule
• Many allow cells to transfer and capture energy as
high-energy compounds
© 2011 Pearson Education, Inc.
Important Functional Groups of Organic Compounds
Functional Group
Structural Formula*
Importance
Examples
Acts as an acid, releasing H+
to become R–COO–
• Fatty acids
• Amino acids
Amino group,
– NH2
Can accept or release H+,
depending on pH; can form
bonds with other molecules
• Amino acids
Hydroxyl group,
– OH
May link molecules through dehydration synthesis; hydrogen bonding between hydroxyl groups and
water molecules affect solubility
• Carbohydrates
• Fatty acids
• Amino acids
Phosphate group,
– PO4
May link other molecules to form
larger structures; may store
energy
• Phospholipids
• Nucleic acids
• High-energy
compounds
Carboxylic
acid group,
– COOH
* The term R group is used to denote the rest of the molecule, whatever that might be. The R group is also known as a side chain.
Figure 2 Section 4
© 2011 Pearson Education, Inc.
Module 2.10: Carbohydrates
• Carbohydrates
• Contain carbon, hydrogen, and carbon in ratio
near 1:2:1
• Examples: sugars and starches
• Roughly 1.5% of total body weight
• Most important as energy sources
© 2011 Pearson Education, Inc.
Module 2.10: Carbohydrates
• Types of carbohydrates
1. Monosaccharide (mono-, single + sakcharon,
sugar)
•
Simple sugar
•
Contains 3–7 carbon atoms
•
Example: glucose (a six-carbon or hexose
sugar)
© 2011 Pearson Education, Inc.
The hexose glucose, a monosaccharide or
simple sugar, the simplest form of
carbohydrate
The atoms in a
glucose molecule
may form either a
straight chain or a
ring. In the body,
the ring form is
more common.
Figure 2 .10
© 2011 Pearson Education, Inc.
1
Module 2.10: Carbohydrates
• Isomers
• Molecules with the same molecular formula but
different structures
• Can be important in molecular function
• Example: glucose and fructose
• Both have molecular formula of C6H12O6
• Structures are different
© 2011 Pearson Education, Inc.
The structures of the monosaccharide
isomers glucose and fructose, both with
chemical formula C6H12O6
Glucose
Fructose
The monosaccharides glucose and fructose are isomers.
Fructose is a hexose found in many fruits. Although its
chemical formula, C6H12O6, is the same as that of glucose,
the arrangement of its atoms differs from that of glucose.
Figure 2 .10
© 2011 Pearson Education, Inc.
2
Module 2.10: Carbohydrates
•
Types of carbohydrates (continued)
2. Disaccharide (di-, two)
•
Two monosaccharides joined
•
Example: sucrose
•
Dehydration synthesis creates disaccharides
•
Hydrolysis breaks them down to
monosaccharides
© 2011 Pearson Education, Inc.
The formation of
the disaccharide
sucrose by
dehydration
synthesis
DEHYDRATION
SYNTHESIS
Glucose
Fructose
Sucrose
Figure 2 .10
© 2011 Pearson Education, Inc.
3
Module 2.10: Carbohydrates
•
Types of carbohydrates (continued)
3. Polysaccharide (poly-, many)
•
Many disaccharides and/or monosaccharides
connected together
Example: starches
•
•
•
•
Major dietary source in potatoes and grains
Broken down into monosaccharides by digestive
system
Glycogen
• Animal starch
• Can be broken down to glucose to supply energy
• Can form from glucose to store energy
© 2011 Pearson Education, Inc.
The production of
two monosaccharides
from a disaccharide
by hydrolysis, the
reverse of
dehydration
synthesis
HYDROLYSIS
Glucose
Fructose
Sucrose
Figure 2 .10
© 2011 Pearson Education, Inc.
4
1
The polysaccharide glycogen, produced when multiple
dehydration synthesis reactions add additional
monosaccharides or disaccharides
Glucose
molecules
The polysaccharide glycogen (GLĪ-kō-jen), or
animal starch, has many side branches
consisting of chains of glucose molecules.
Muscle cells make and store glycogen. When
these cells have a high demand for glucose,
glycogen molecules are broken down; when the
demand is low, they absorb glucose from the
bloodstream and rebuild glycogen reserves.
Figure 2.10
© 2011 Pearson Education, Inc.
5
Figure 2 .10
© 2011 Pearson Education, Inc.
2
Module 2.10 Review
a. List the three structural classes of carbohydrates, and
give an example of each.
b. A food contains organic molecules with the elements
C, H, and O in a ratio of 1:2:1. What class of
compounds do these molecules belong to, and what
are their major functions in the body?
c. Predict the reactants and the type of chemical
reaction involved when muscle cells make and store
glycogen.
© 2011 Pearson Education, Inc.
Module 2.11: Lipids
•
Lipids (lipos, fat)
•
Carbon to hydrogen ratio is near 1:2
•
Much less oxygen compared to carbohydrates with
similar number of carbon atoms
•
Examples: fats, oils, waxes
•
Most are insoluble in water
•
Special transport mechanisms for them in the blood
•
Essential components of all cells
•
Provide twice as much energy as carbohydrates
•
Normally account for 12–24% of total body weight
•
Must obtain some from diet
© 2011 Pearson Education, Inc.
A triglyceride, a lipid produced by dehydration synthesis
between a glycerol molecule and three fatty acids
Glycerol
Fatty acids
Fatty Acid 1
Saturated
Fatty Acid 2
Saturated
Fatty Acid 3
Unsaturated
HYDROLYSIS
DEHYDRATION
SYNTHESIS
Dehydration synthesis can produce a
monoglyceride (mon-ō-GLI-ser-īd),
consisting of glycerol + one fatty acid.
Subsequent reactions can yield a
diglyceride (glycerol + two fatty acids)
and then a triglyceride (glycerol +
three fatty acids). Hydrolysis breaks the
glycerides into fatty acids and glycerol.
Triglyceride
Figure 2 .11
© 2011 Pearson Education, Inc.
4
Module 2.11: Lipids
•
Fatty acids
•
Long carbon chains with attached hydrogens
•
Two ends
1. Head
•
Has carboxylic group (–COOH)
•
Hydrophilic
2. Tail
•
Hydrophobic
© 2011 Pearson Education, Inc.
A fatty acid, the building block of lipids, which includes a tail and
a head composed of a carboxylic acid group: —COOH
The carbon chain attached
to the carboxylic acid group
is known as the hydrocarbon
tail of the fatty acid. The
hydrocarbon tail is
hydrophobic, so fatty acids
have very limited solubility
in water. In general, the
longer the hydrocarbon tail,
the lower the solubility of the
molecule.
Lauric acid (C12H24O2)
© 2011 Pearson Education, Inc.
Figure 2 .11
1
Module 2.11: Lipids
•
Fatty acids
•
Two types
1. Saturated fatty acid
•
Each carbon has four attached hydrogens
2. Unsaturated fatty acid
•
Contains double bonds
•
One double bond = monounsaturated
•
>1 double bond = polyunsaturated
•
Has fewer attached hydrogens
•
Changes metabolism of the molecule
© 2011 Pearson Education, Inc.
Figure 2 .11
© 2011 Pearson Education, Inc.
2
An unsaturated fatty acid, in which one or more of the
single covalent bonds between carbon atoms has
been replaced by a double covalent bond
Double
covalent
bond
Figure 2 .11
© 2011 Pearson Education, Inc.
3
Module 2.11: Lipids
•
Glycerides
•
Fatty acid chains attached to a glycerol
molecule
•
Three types through dehydration synthesis
1. Monoglyceride (glycerol + one fatty acid)
2. Diglyceride (glycerol + two fatty acids)
3. Triglyceride (glycerol + three fatty acids)
© 2011 Pearson Education, Inc.
Figure 2 .11
© 2011 Pearson Education, Inc.
4
Module 2.11 Review
a. Describe lipids.
b. Describe the structures of saturated and
unsaturated fatty acids.
c. In the hydrolysis of a triglyceride, what are the
reactants and the products?
© 2011 Pearson Education, Inc.
Module 2.12: Diverse functions of lipids
•
Functions of lipids
•
Chemical messengers
•
Components of cellular structures (structural
lipids)
•
Maintain cell membrane
•
Allow separation of different aqueous solutions
© 2011 Pearson Education, Inc.
Module 2.12: Diverse functions of lipids
•
Eicosanoids
•
Derived from arachidonic acid
•
Examples:
•
Leukotrienes
•
•
Produced by cells in response to injury
Prostaglandins
•
Released by cells to coordinate local cellular
activities
© 2011 Pearson Education, Inc.
Figure 2 .12
© 2011 Pearson Education, Inc.
1
Module 2.12: Diverse functions of lipids
•
Steroids
•
Large molecules with four carbon rings
•
Differ in attached functional groups
•
Examples:
•
Cholesterol (chole-, bile + steros, solid)
•
•
Functions to maintain plasma membranes,
growth, and division
Hormones
•
Regulation of sexual and other metabolic
functions
© 2011 Pearson Education, Inc.
Three steroid molecules, which share a distinctive
carbon-ring framework
Cholesterol
Estrogen
© 2011 Pearson Education, Inc.
Testosterone
Figure 2.12
2
Module 2.12: Diverse functions of lipids
•
Phospholipids and Glycolipids
•
Phospholipid
•
•
Phosphate linking a diglyceride to a nonlipid
group
Glycolipid
•
Carbohydrate attached to a diglyceride
© 2011 Pearson Education, Inc.
The structures of
phospholipids and
glycolipids
Nonlipid group
Carbohydrate
Phosphate group
Glycerol
Glycerol
Fatty
acids
In a phospholipid, a phosphate group
links a diglyceride to a nonlipid group.
Fatty
acids
In a glycolipid, a carbohydrate is
attached to a diglyceride.
Figure 2 .12
© 2011 Pearson Education, Inc.
3
Module 2.12: Diverse functions of lipids
•
Phospholipids and Glycolipids
•
Structurally related
•
Can be synthesized by cells primarily from
fatty acids
•
Like fatty acids
•
•
Tails are hydrophobic
•
Other end is hydrophilic
In water, can form large droplets (micelles)
© 2011 Pearson Education, Inc.
Hydrophilic
heads
Water
Phospholipid
Hydrophobic
tails
Phospholipid
Glycolipid
Glycolipid
The arrangement of phospholipid
and glycoplid molecules in
droplets called micelles
Micelles
Figure 2 .12
© 2011 Pearson Education, Inc.
4
Module 2.12 Review
a. Why is cholesterol necessary in the body?
b. Describe the basic functions of eicosanoids,
steroids, phospholipids, and glycolipids.
c. Describe the orientations of phospholipids
and glycolipids when they form a micelle.
© 2011 Pearson Education, Inc.
Module 2.13: Proteins
•
Proteins
•
Most abundant organic molecule in the body
•
In many ways, are most important
•
Normally account for 20% of total body weight
•
Contain carbon, hydrogen, oxygen, and
nitrogen
•
•
Possibly sulfur and phosphorus as well
Consist of long chains of amino acids
•
20 amino acids in the body
•
Typical protein contains 1000 amino acids
© 2011 Pearson Education, Inc.
Module 2.13: Proteins
•
Amino acids
•
All have same structural components
•
Central carbon
•
Amino group
•
Carboxyl group
•
R group (variable side chain)
© 2011 Pearson Education, Inc.
Amino group
Central carbon
Carboxylic acid group
The structural elements of amino
acids, the basic units of proteins
R group (variable side chain
of one or more atoms)
Figure 2 .13
© 2011 Pearson Education, Inc.
1
Module 2.13: Proteins
•
Peptides
•
Amino acids linked through dehydration synthesis
•
Covalent bond connects the carboxylic acid group
of one amino acid to the amino group of another
•
= Peptide bond
•
Dipeptide
•
•
Polypeptides
•
•
Two amino acids linked together
Three or more amino acids linked together
Peptides of over 100 amino acids are called proteins
© 2011 Pearson Education, Inc.
The linkage of two amino acids
by dehydration synthesis, and
the breakage of peptide bonds
by hyrdolysis
Glycine (gly)
Alanine (ala)
Adjacent amino acids
can be linked
together by a
covalent bond that
connects the
carboxylic acid group
of one amino acid to
the amino group of
another.
Peptide bonds can
also be broken
through hydrolysis;
the hydrolysis of a
dipeptide yields a
pair of amino acids.
The bond between amino acids is known as a
peptide bond. Molecules consisting of amino
acids held together by peptide bonds are called
peptides. This molecule is called a dipeptide
because it contains two amino acids.
Figure 2 .13
© 2011 Pearson Education, Inc.
2
Module 2.13: Proteins
•
Protein structure
1. Primary structure
•
Sequence of amino acids
2. Secondary structure
•
Molecule shape changes due to bonds at different
parts of the polypeptide chain
•
Example: hydrogen bonds
3. Tertiary structure
•
•
Coiling and folding giving the protein a final 3-D shape
Usually from interactions of the protein and
surrounding water molecules
© 2011 Pearson Education, Inc.
Module 2.13: Proteins
•
Protein structure (continued)
4. Quaternary structure
•
Interaction between multiple polypeptide chains
to form a protein complex
•
Examples:
•
Hemoglobin
•
•
Binds oxygen in red blood cells
Keratin and collagen
•
© 2011 Pearson Education, Inc.
Fibrous proteins that give strength to tissues
The four levels of structural complexity of proteins
Primary Structure
A polypeptide, a linear chain
of amino acids
Secondary Structure
Hydrogen
bond
Hydrogen bond
or
Alpha-helix
Pleated sheet
Secondary structure
results from bonds
between atoms at
different parts of the
polypeptide chain.
Hydrogen bonding, for
example, may create
either a simple spiral,
known as an
alpha-helix, or a flat
pleated sheet.
Tertiary Structure
Heme
unit
Quaternary Structure
Hemoglobin
(globular protein)
Tertiary structure results from the
complex coiling and folding that
gives a protein its final
three-dimensional shape. Tertiary
structure results primarily from
interactions between the polypeptide
chain and the surrounding water
molecules, and to a lesser extent
from interactions between the R
groups of amino acids in different
parts of the molecule.
Quaternary structure results
from the interaction between
individual polypeptide chains to
form a protein complex. The
protein hemoglobin contains
four polypeptide subunits.
Hemoglobin is found within red
blood cells, where it binds and
or
transports oxygen. It is an
example of a globular protein. In
keratin and collagen, three
linear subunits intertwine, forming
Heme
a fibrous protein. The
unit
three-dimensional shape of a
protein plays an essential role in
Keratin
determining its functional
(fibrous protein) properties.
Figure 2 .13
© 2011 Pearson Education, Inc.
3
Module 2.13: Proteins
•
Protein structure
•
Denaturation
•
Protein shape and function deteriorate
•
Occurs under extreme conditions
•
•
Temperature
•
pH
Irreparable damage can occur to tissues and
organs
© 2011 Pearson Education, Inc.
Module 2.13 Review
a. Describe proteins.
b. What kind of bond forms during the
dehydration synthesis of two amino acids?
c. Why does boiling a protein affect its structural
and functional properties?
© 2011 Pearson Education, Inc.
Module 2.14: Enzymes are regulatory
proteins
•
Enzymes facilitate most everything that
occurs inside the body
•
Substrates
•
Reactants in enzymatic reactions
•
Must bind to a specific region of an enzyme
•
= Active site
•
Binding reinforced by hydrogen bonds
•
Enzyme-substrate complex
•
Once binding to enzyme occurs
Animation: Chemical Reactions: Enzymes
© 2011 Pearson Education, Inc.
Module 2.14: Enzymes are regulatory
proteins
•
•
Structure of interacting molecules is very
important in enzymatic reactions
•
Substrate and enzyme fit in a “lock and key”
fashion
•
Substrate binding results in a temporary,
reversible change in shape of enzyme
Enzyme is free to catalyze other reactions
once product is complete
•
Saturation limit
•
Substrate concentration required to have
maximum rate of reaction
© 2011 Pearson Education, Inc.
The events in the actions of enzymes
Substrates
Active site
Enzymesubstrate
complex
Figure 2 .14
© 2011 Pearson Education, Inc.
1
-
4
Module 2.14 Review
a. Define active site.
b. What are the reactants in an enzymatic
reaction called?
c. Relate an enzyme’s structure to its reaction
specificity.
© 2011 Pearson Education, Inc.
Module 2.15: High-energy compounds
•
High-energy compounds
•
Donate energy to chemical reactions to form
products
•
Contain high-energy bonds (covalent)
•
Release energy when broken
© 2011 Pearson Education, Inc.
Module 2.15: High-energy compounds
•
Adenosine triphosphate (ATP)
•
Most common high-energy compound
•
Provides energy for many vital body functions
•
Examples:
•
Contraction of muscles
•
Synthesis of proteins, carbohydrates, and lipids
© 2011 Pearson Education, Inc.
Module 2.15: High-energy compounds
•
Adenosine triphosphate (ATP) (continued)
•
Formation
•
•
Adenosine monophosphate (AMP)
•
Adenosine
•
Ribose sugar
•
One phosphate
Adenosine diphosphate (ADP)
•
•
Adenosine triphosphate (ATP)
•
•
+ One more phosphate
+ One more phosphate
Formation of ATP from ADP is reversible
© 2011 Pearson Education, Inc.
The reversible reaction by which ATP is
formed from ADP
Figure 2 .15
© 2011 Pearson Education, Inc.
2
Module 2.15 Review
a. Where do cells obtain the energy needed for
their vital functions?
b. Describe ATP.
c. Compare AMP with ADP.
© 2011 Pearson Education, Inc.
Module 2.16: DNA and RNA
•
Nucleic acids
•
•
Large organic molecules
Composed of carbon, hydrogen, oxygen, nitrogen,
and phosphorus
Primary function is storage and transfer of information
•
•
•
•
Particularly synthesis of proteins
Consist of two long chains formed from dehydration
synthesis of subunits (nucleotides)
Two classes
1. Deoxyribonucleic acid (DNA)
2. Ribonucleic acid (RNA)
© 2011 Pearson Education, Inc.
Module 2.16: DNA and RNA
•
Nucleotide components
1. Nitrogenous base
a. Purines
•
Adenine
•
Guanine
b. Pyrimidines
•
Cytosine
•
Thymine (only in DNA)
•
Uracil (only in RNA)
2. Pentose sugar
3. Phosphate group
© 2011 Pearson Education, Inc.
The structure of nucleotides, the subunits
of a nucleic acid
Sugar
Phosphate
group
Nitrogenous
base
Figure 2 .16
© 2011 Pearson Education, Inc.
1
The structures of the two types of nitrogenous
bases of nucleic acids
Nitrogenous bases
The purines adenine and guanine are found in
both DNA and RNA.
Adenine
Guanine
There are three important pyrimidines. DNA and
RNA both contain cytosine, thymine is found
only in DNA, and uracil is found only in RNA.
Cytosine
Thymine
(DNA only)
Uracil
(RNA only)
Figure 2.16
© 2011 Pearson Education, Inc.
2
Module 2.16: DNA and RNA
•
Phosphate and sugars of adjacent
nucleotides are linked together to form
nucleic acids
•
•
Forms the “backbone” of the nucleic acid
Sequence of attached nitrogenous bases
carries the information
© 2011 Pearson Education, Inc.
The production of a linear
sugar-to-phosphate-to-sugar “backbone” of a nucleic
acid by dehydration synthesis
DEHYDRATION
SYNTHESIS
Figure 2.16
© 2011 Pearson Education, Inc.
2
Module 2.16: DNA and RNA
•
DNA
•
Pair of nucleotide chains
•
•
•
= Complementary strands
Twist around each other to form a double helix
Hydrogen bonds between opposing nitrogenous
bases hold two strands together
•
= Complementary base pairs
•
•
Adenine – Thymine (A-T)
Cytosine – Guanine (C-G)
Animation: Protein Synthesis: DNA Molecule
Animation: Protein Synthesis: How DNA
Information is Used to Make Protein
© 2011 Pearson Education, Inc.
The structure of a DNA molecule, in which two
complementary strands (nucleotide chains) are linked by
hydrogen bonds between complementary base pairs
Deoxyribose
Adenine
Thymine
Phosphate group
Hydrogen bond
The two strands of DNA twist
around one another in a
double helix that resembles a
spiral staircase.
DNA strand 1
DNA strand 2
Hydrogen bond
Cytosine
Guanine
Figure 2.16
© 2011 Pearson Education, Inc.
3
Module 2.16: DNA and RNA
•
RNA
•
Single chain of nucleotides
•
Three types
1. Messenger RNA (mRNA)
2. Transfer RNA (tRNA)
3. Ribosomal RNA (rRNA)
© 2011 Pearson Education, Inc.
A molecule of RNA, which consists of
a single chain of nucleotides
single-strand
sugar-phosphate
backbone
Nitrogenous
bases
Figure 2.16
© 2011 Pearson Education, Inc.
4
Figure 2 .16
© 2011 Pearson Education, Inc.
3
Module 2.16 Review
a. Describe nucleic acids.
b. Explain how the complementary strands of
DNA are held together.
c. A large organic molecule composed of ribose,
nitrogenous bases, and phosphate groups is
which kind of nucleic acid?
© 2011 Pearson Education, Inc.