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PowerPoint® Lecture Slides
prepared by
Betsy C. Brantley
Valencia College
CHAPTER
2
Chemical Level
of Organization
© 2013 Pearson Education, Inc.
Chapter 2 Learning Outcomes
• Section 1: Atoms and Molecules
• 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.
• Section 2: Chemical Reactions
• 2.4
• Use chemical notation to symbolize chemical reactions.
• 2.5
• Distinguish among the major types of chemical reactions that
are important for studying physiology.
© 2013 Pearson Education, Inc.
Chapter 2 Learning Outcomes
• Section 3: The Importance of Water in the Body
• 2.6
• Explain how the chemical properties of water affect the solubility
of inorganic and organic molecules.
• 2.7
• Discuss the importance of pH and the role of buffers in body
fluids.
• Section 4: Metabolites and Nutrients
• 2.8
• Discuss the structures and functions of carbohydrates.
• 2.9
• Discuss the structures and functions of lipids.
• 2.10
• Discuss the structures and diverse functions of various types of
lipids: steroids, phospholipids, and glycolipids.
© 2013 Pearson Education, Inc.
Chapter 2 Learning Outcomes
• 2.11
• Discuss protein structure and the essential functions of proteins
within the body.
• 2.12
• Explain how enzymes function within the body.
• 2.13
• Discuss the structure and function of high-energy compounds.
• 2.14
• Compare and contrast the structures and functions of DNA and
RNA.
© 2013 Pearson Education, Inc.
Chemical Level of Organization (Section 1)
• Chemistry studies structure of matter
• Matter – anything that takes up space and has
mass
• Mass – the amount of material in matter
• On Earth, mass equivalent to weight
© 2013 Pearson Education, Inc.
Atoms and Molecules (Section 1)
• Atom – smallest stable unit of matter
• Composed of subatomic particles
• Protons
• Have a positive electrical charge
• Neutrons
• Are electrically neutral (uncharged)
• Electrons
• Have a negative electrical charge
• Are much smaller than protons or neutrons (about
1/800 the mass)
© 2013 Pearson Education, Inc.
Atoms are composed of subatomic particles
Protons
Neutrons
Electrons
© 2013 Pearson Education, Inc.
Figure 2 Section 1 1
Atomic Structure (Section 1)
• An atom can be subdivided into:
• Nucleus
• At the center of an atom
• Contains one or more protons
• May also contain neutrons
• Mass of atom determined by number of protons and
neutrons
• Electron cloud
• Created by whirl of electrons around the nucleus
© 2013 Pearson Education, Inc.
Nucleus and electron cloud of an atom
Nucleus
Electron cloud
© 2013 Pearson Education, Inc.
Figure 2 Section 1 2
Molecular Structure (Section 1)
• When atoms interact, they produce larger, more
complex structures called molecules
• All matter composed of arrangements of atoms
• Variation in matter characteristics results from types of
atoms and ways they interact
© 2013 Pearson Education, Inc.
Molecule forms when atoms interact
© 2013 Pearson Education, Inc.
Figure 2 Section 1 3
Atomic Number and Mass Number (2.1)
• Atoms normally contain equal numbers of protons,
neutrons, and electrons
• Atomic number
• Total number of protons in an atom
• Mass number
• Total number of protons and neutrons in an atom
• Element
• Substance composed only of atoms with same atomic
number
© 2013 Pearson Education, Inc.
Hydrogen – The Simplest Atom (2.1)
• Chemical symbol H
• Atomic number of 1
• Contains 1 proton and 1 electron
• Proton in the center of the atom (the nucleus)
• Electron whirls around the nucleus in the electron
cloud
• The negatively charged electron is attracted to the
positively charged proton, so it stays in "orbit"
• Electron orbit is often depicted as a circular electron
shell
© 2013 Pearson Education, Inc.
Hydrogen ion representations
Electron cloud representation
Electron shell representation
© 2013 Pearson Education, Inc.
Figure 2.1 1
Isotopes (2.1)
• Atoms of single element can differ in number of neutrons
• Isotopes
• Atoms with same number of protons but different numbers of
neutrons
• Identical chemical properties
• Different mass number
• For example:
• Hydrogen with 1 proton and 0 neutrons = mass number 1
• Hydrogen with 1 proton and 1 neutron = mass number 2
• Hydrogen with 1 proton and 2 neutrons = mass number 3
© 2013 Pearson Education, Inc.
Electron-shell model
Electron-shell model
Electron
shell
Hydrogen-1
mass number: 1
© 2013 Pearson Education, Inc.
Hydrogen-2,
deuterium
mass number: 2
Hydrogen-3,
tritium
mass number: 3
Figure 2.1 2
Atomic Weight (2.1)
• Actual mass of an atom
• Expressed in atomic mass units (amu) or daltons
• One amu close to weight of one proton or neutron
• Equals average mass of an element, including different
isotopes in proportion
• For example:
• Hydrogen atomic number = 1 (one proton)
• Hydrogen atomic weight = 1.0079
• Not all hydrogen atoms have 0 neutrons
• 0.015 percent have 1 neutron (mass number 2)
• Lower percentage have 2 neutrons (mass number 3)
© 2013 Pearson Education, Inc.
Atomic weight of hydrogen and isotopes
Average mass
Average mass
amu
amu
Atomic weight of hydrogen-1 = 1
Atomic weight of a mixture of
hydrogen isotopes = 1.0079
© 2013 Pearson Education, Inc.
Figure 2.1 3
Elements (2.1)
• Human body contains 27 elements
• 13 of those are considered "common"
• 14 of those are trace elements (present in very small
amounts)
• 92 elements exist in nature
• Another 24 or so created in research laboratories
• Each element has a chemical symbol
• Based on:
• English names (e.g., O for oxygen, C for carbon)
• Names in other languages (e.g., Na for sodium from
the Latin natrium)
© 2013 Pearson Education, Inc.
© 2013 Pearson Education, Inc.
Figure 2.1 4
Module 2.1 Review
a. Define an element.
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?
© 2013 Pearson Education, Inc.
Electrons and Energy Levels (2.2)
• Atoms are electrically neutral
• Every positive proton is balanced by a negative electron
• Electrons occupy an orderly series of energy
levels
• Can be diagrammed as a series of concentric electron
shells
• First shell (closest to nucleus) is the lowest energy level
© 2013 Pearson Education, Inc.
Reactive versus Inert Elements (2.2)
• Outermost energy level is atom's "surface"
• Atoms with unfilled outer shells (reactive)
• Tend to react with other atoms to fill outer shell
• Atoms with full outer shells (inert)
• More stable
• Do not readily react with other atoms
• For example, helium and neon
• Called inert gases
© 2013 Pearson Education, Inc.
Reactive and inert elements
Reactive elements
Inert elements
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
The second and
third energy levels
can each contain up
to eight electrons.
Lithium has
one electron
in the second
energy level; it
is extremely
reactive.
Lithium, Li
Atomic number: 3
Mass number: 6
(3 protons + 3 neutrons)
3 electrons
© 2013 Pearson Education, Inc.
Helium, He
Atomic number: 2
Mass number: 4
(2 protons + 2 neutrons)
2 electrons
Neon has eight
electrons in the
second energy
level; it does
not react with
other atoms.
Neon, Ne
Atomic number: 10
Mass number: 20
(10 protons + 10 neutrons)
10 electrons
Figure 2.2 1 – 2
Cations (2.2)
• Reactive elements gain, lose, or share electrons to
fill outermost shells
• Losing an electron means:
• Fewer electrons (negative) than protons (positive)
• Net positive charge
• Called a positive ion or cation
• One missing electron = charge of +1
• More electrons missing = more positive charge (e.g.,
+2, +3, +4)
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Cation has a positive charge
+
Sodium atom, Na (reactive)
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Sodium ion, Na+ (stable)
Figure 2.2 3
Anions (2.2)
• In a reaction to achieve stability, gaining an
electron means:
• More electrons (negative) than protons (positive)
• Net negative charge
• Called a negative ion or anion
• One extra electron = charge of –1
• More electrons gained = more negative charge
(e.g., –2, –3, –4)
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Anion has a negative charge
Chlorine atom, Cl (reactive)
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Chloride ion, Cl– (stable)
Figure 2.2 4
Chemical Bonds (2.2)
• Atoms interact to stabilize outer energy levels
• Often results in formation of chemical bonds
• Hold atoms together after end of reaction
© 2013 Pearson Education, Inc.
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.
© 2013 Pearson Education, Inc.
Compounds (2.3)
• Chemical bonding creates new chemical entities
• Compound
• Chemical substance made of atoms of two or more
different elements
• Type of bond holding atoms together does not matter
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Ionic Bonds (2.3)
• One of most common types of chemical bonds
• Created by electrical attraction between cations
(positive ions) and anions (negative ions)
• Involve transfer of one or more electrons from one
ion to the other
© 2013 Pearson Education, Inc.
Ionic bond between sodium and chloride
Step 1: Formation of sodium and
chloride ions. The sodium atom
loses an electron to the chlorine
atom. This produces two stable ions
with filled outer energy levels.
Sodium atom
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 ion (Na+)
Sodium chloride (NaCl)
Chlorine atom
© 2013 Pearson Education, Inc.
Chloride ion (Cl–)
Figure 2.3 1
Crystal of sodium chloride
Chloride ions (Cl–)
Sodium ions (Na+)
© 2013 Pearson Education, Inc.
Figure 2.3 2
Covalent Bonds (2.3)
• Formed between atoms
• Involve sharing of electrons between atoms
(instead of gaining or losing)
• Form molecules
• Atoms of one or more elements held together by
covalent bonds
• Typically, electrons are equally shared, producing a
nonpolar molecule
© 2013 Pearson Education, Inc.
Examples of Nonpolar Molecules (2.3)
• Hydrogen molecule
• Pair of hydrogen atoms sharing an electron from each atom
• One electron from each atom shared – single covalent bond
• Oxygen molecule
• Pair of oxygen atoms sharing electrons from each atom
• Two pairs of electrons shared – double covalent bond
• Carbon dioxide
• One carbon atom sharing electrons with two oxygen atoms
• Two pairs of electrons shared with each oxygen – two double
covalent bonds
© 2013 Pearson Education, Inc.
© 2013 Pearson Education, Inc.
Figure 2.3 3
Unequal Sharing of Electrons and Polar
Molecules (2.3)
• When electrons spend more time around one
atom of a molecule, that atom has a slightly
negative charge
• Water molecule
• Oxygen atom carries slightly negative charge (δ–)
• Hydrogen atoms carry slightly positive charge (δ+)
• Forms a polar molecule
• Bonds are called polar covalent bonds
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Hydrogen and oxygen form water by covalent bonding
Hydrogen
atom
Hydrogen
atom
Oxygen atom
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Water molecule
Figure 2.3 4
Water is a polar molecule
Hydrogen atom
Positive pole
Oxygen atom
Negative pole
© 2013 Pearson Education, Inc.
Figure 2.3 5
Hydrogen Bonds (2.3)
• Bonds form between polar molecules
• For example, between water molecules
• Small positive charge on hydrogen atom of one
molecule
• Attracted to small negative charge on oxygen atom of
another molecule
• Weak attractive force is called a hydrogen bond
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Hydrogen bonds between water molecules
KEY
Hydrogen
Oxygen
Hydrogen
bond
© 2013 Pearson Education, Inc.
Figure 2.3 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.
© 2013 Pearson Education, Inc.
Chemical Reactions (Section 2)
• Chemical reactions are constantly occurring in our
cells
• New chemical bonds are formed; existing bonds
broken
• Reactants are atoms in the reacting substances
• Products are the results of the reactions
• Metabolism includes all the reactions in the body at any
moment
© 2013 Pearson Education, Inc.
Chemical Reactions in Cells (Section 2)
• Cells use chemical reactions
• To provide energy needed for maintaining homeostasis
and essential functions such as:
• Growth
• Maintenance and repair
• Cell division
• Secretion
• Contraction
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Metabolism in a cell
Essential activities
• Maintenance and
repair
• Growth
• Division
• Special functions
Typical
cell
Energy transfer
and use
Substances
absorbed
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Chemical
reactions
Figure 2 Section 2 1
Work and Energy (Section 2)
• Work
• Movement of an object or change in physical structure of matter
• Can be macroscopic (e.g., moving muscles) to microscopic (e.g.,
formation of molecules)
• Energy
• Capacity to perform work
• Kinetic energy
• Energy of motion (e.g., throwing a ball)
• Can be transferred to another object
• Potential energy
• Stored energy (e.g., stretched spring)
• Has the potential to do work
© 2013 Pearson Education, Inc.
Work and Energy Transfer (Section 2)
• Cellular work includes:
• Producing complex molecules
• Moving material into or out of a cell
• Transfer of energy not 100 percent efficient
• Skeletal muscle cells contain potential energy (position
of proteins, covalent bonds)
• With contraction, potential energy is converted to kinetic
energy (movement)
• Some of the energy is released as heat
• Body temperature increases with exercise
© 2013 Pearson Education, Inc.
Chemical Notation – Atoms and Molecules (2.4)
• Chemical notation
• Allows precise and brief description of complex events
• May be used to calculate weights of reactants in a reaction
• Atoms
• Symbol of an element indicates one atom of that element
• Number preceding symbol indicates number of atoms (e.g., 2H)
• Molecules
• Chemical (or molecular) formula provides information about
elements and number of atoms in a molecule
• Subscript following the symbol indicates the number of atoms of
that element (e.g., H2O)
© 2013 Pearson Education, Inc.
Chemical Notation – Reactions (2.4)
• Reactants at the beginning
• Reaction produces one or more products
• Arrow indicates direction of reaction from
reactants to products (e.g.,
)
• Atoms not created or destroyed, just rearranged
• Numbers of atoms of each element on each side
of the arrow must be equal for an equation to be
balanced
© 2013 Pearson Education, Inc.
Chemical Notation – Ions (2.4)
• Superscript plus or minus sign following an
element symbol indicates an ion
• Single plus sign indicates cation with charge of +1
• Single minus sign indicates anion with
charge of –1
• If more than one electron gained or lost, charge
indicated by superscript number before the plus or
minus sign (e.g., Ca2+)
© 2013 Pearson Education, Inc.
© 2013 Pearson Education, Inc.
Figure 2.4 1
Activation Energy (2.4)
• Minimum energy required to activate reactants in a
reaction and allow reaction to proceed
• Outside the body, may be acquired by extremes in
temperature, pressure, or lethal chemical factors
• Inside the body, cells use special proteins called
enzymes
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Activation energy without an enzyme
Energy
In the external environment, extreme
conditions can provide the activation
energy. For example, complex sugars can
be broken down in a laboratory by boiling
them in an acidic solution.
Activation
energy
Progress of reaction
© 2013 Pearson Education, Inc.
Figure 2.4 2
Enzymes and Chemical Reactions (2.4)
• Enzymes
• Promote chemical reactions
• Lower the required activation energy
• Allow reactions to proceed under conditions compatible
with life
• Function as catalysts
• Accelerate chemical reaction without being permanently
changed or consumed
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Enzymes reduce activation energy
Energy
Specific enzymes lower the activation
energy so that important cellular
reactions will occur.
Activation
energy
Progress of reaction
© 2013 Pearson Education, Inc.
Figure 2.4 3
Enzymes and activation energy
Specific enzymes lower the activation
energy so that important cellular
reactions will occur.
Activation
energy
Progress of reaction
© 2013 Pearson Education, Inc.
Energy
Energy
In the external environment, extreme
conditions can provide the activation
energy. For example, complex sugars can
be broken down in a laboratory by boiling
them in an acidic solution.
Activation
energy
Progress of reaction
Figure 2.4 2 – 3
Metabolic Pathway (2.4)
• Series of complex reactions occurring in the body
• Each reaction interlocking with next step
• Each reaction controlled by specific enzyme
• Reaction sequence called metabolic pathway
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Metabolic pathway
Step 1
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Step 2
Step 3
and so on.
Figure 2.4 4
Exergonic and Endergonic (2.4)
• Reactions require activation energy to start
• Reactions then categorized as:
• Exergonic
• Overall net release of energy
• Fairly common in the body
• Tend to generate heat
• Endergonic
• More energy is required to begin than is released
• Include reactions to build molecules
© 2013 Pearson Education, Inc.
Module 2.4 Review
a. 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.
b. What is an enzyme?
c. Why are enzymes needed in our cells?
© 2013 Pearson Education, Inc.
Types of Chemical Reactions (2.5)
• Decomposition reactions
• Synthesis reactions
• Exchange reactions
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Decomposition Reactions (2.5)
• Decomposition
• A reaction that breaks a molecule into smaller fragments
• Occurs inside and outside cells
• For example, decomposition reactions in the digestive
tract break down molecules into smaller fragments that
can then be absorbed
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Simple decomposition reaction
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Figure 2.5 A
Hydrolysis (2.5)
• A specific type of decomposition reaction that
involves water
• One of the bonds in a molecule is broken
• Components of water molecule (H and OH) are
added to the fragments
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Hydrolysis
© 2013 Pearson Education, Inc.
Figure 2.5 B
Catabolism (2.5)
• Collective term for decomposition reactions
• Refers to breaking covalent bonds (potential
energy)
• Releases kinetic energy that can perform work
• Body can use energy for growth, movement, and
reproduction
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Catabolism
© 2013 Pearson Education, Inc.
Figure 2.5 C
Synthesis Reactions (2.5)
• Synthesis
• Opposite of decomposition
• Assembles smaller molecules into larger molecules
• May involve combining atoms or molecules
• For example, formation of water
• Always involves formation of new chemical bonds
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Simple synthetic reaction
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Figure 2.5 D
Dehydration Synthesis (2.5)
• Dehydration synthesis (condensation)
• Formation of a complex molecule by removing a water
molecule
• Opposite of hydrolysis
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Dehydration synthesis
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Figure 2.5 E
Anabolism (2.5)
• Collective term for synthesis reactions
• Refers to forming new chemical bonds
• Requires energy
• Energy usually from other catabolic reactions
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Chemical Reactions Are Reversible (2.5)
• Many biological reactions freely reversible
• Can operate in either direction
• Coupled synthesis and decomposition reactions
• At equilibrium, rates of both reactions are in
balance
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Reversible reaction
At equilibrium, the two reaction
rates are in balance.
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Figure 2.5 F
Exchange Reactions (2.5)
• Parts of the reacting molecules are shuffled
around to produce new products
• May involve both decomposition and synthesis
reactions
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Exchange reaction
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Figure 2.5 G
Module 2.5 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?
© 2013 Pearson Education, Inc.
The Importance of Water in the Body
(Section 3)
• Water
• Most important component of your body
• Makes up about 2/3 of total body weight
• Affects all physiological systems
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Properties of Water (Section 3)
• Important properties of water
• Lubrication
• Reactivity
• High heat capacity
• Solubility
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Important properties of water
Lubrication
Water is an effective lubricant,
reducing friction within joints
and in body cavities.
Reactivity
Dehydration
synthesis
In our bodies, chemical
reactions occur in water.
Hydrolysis
High heat capacity
Water has a high heat capacity.
Solubility
Water is a solvent for many subtances.
Solvent
Solutes
© 2013 Pearson Education, Inc.
Solution
Figure 2 Section 3 1
Water Is an Effective Lubricant (Section 3)
• Very little friction between water molecules
• Makes water an effective lubricant
• Even a thin layer of water between surfaces reduces
friction
• In the body, water reduces friction in joints and body
cavities
© 2013 Pearson Education, Inc.
Water and Chemical Reactions (Section 3)
• Chemical reactions occur in water
• Dehydration synthesis and hydrolysis reactions
involve water as a reactant
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Water Has an Unusually High Heat Capacity
(Section 3)
• Heat capacity
• Ability to absorb and retain heat
• Water molecules attracted to one another
• Gives water a high heat capacity
• Temperature has to greatly increase before individual
water molecules break free to become water vapor
• Water carries heat with it when it becomes a gas
• Cooling effect of perspiration/evaporation
• Large amount of water changes temperature very slowly
• Thermal inertia
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Many Compounds Are Soluble in Water
(Section 3)
• Individual particles of many compounds disperse
easily within water
• Solution
• Uniform mixture of two or more substances
• Solvent
• Medium in which other atoms, ions, or molecules are
dispersed
• Solutes
• The dispersed substances
• Aqueous solutions
• Water is the solvent
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Solubility
Solubility
Water is a solvent for many subtances.
Solvent
Solutes
© 2013 Pearson Education, Inc.
Solution
Figure 2 Section 3 1
Ionic Compounds in Water (2.6)
• Ionization (dissociation)
• Process of breaking ionic bonds as ions interact with
poles of water molecule
• A water molecule is polar
• Has a more positive end and a more negative end
• From asymmetric position of hydrogen atoms
• Portion near hydrogen atoms is more positive
• Portion near oxygen atoms is more negative
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Water is polar
Negative
pole
Positive
pole
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Figure 2.6 1
Hydration Spheres around Ions (2.6)
• In solution
• Anions (negative ions) are surrounded by positive poles
of water molecule
• Cations (positive ions) are surrounded by negative
poles of water molecule
• Hydration sphere
• A layer of water molecules around an ion in solution
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Ionic compounds dissociate in water
Sodium chloride crystal
Hydration
spheres
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NaCl in solution
Figure 2.6 2
Hydration Spheres around Molecules (2.6)
• Molecules with polar covalent bonds also attract
polar water molecules
• Water forms hydration sphere around molecule
• If the molecule binds to water strongly, it will
dissolve
• Hydrophilic
• Molecules that interact with water strongly (e.g.,
dissolve)
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Hydration spheres form around large molecules
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Glucose
molecule
Glucose molecule in solution
Figure 2.6 3
Electrolytes (2.6)
• Soluble inorganic molecules whose ions will
conduct an electrical current in solution
• Cations move toward negative side of electrical
field
• Anions move toward positive side of electrical field
• Small electrical currents essential to:
• Muscle contraction
• Nerve function
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An aqueous solution containing anions and cations will conduct an electrical current
© 2013 Pearson Education, Inc.
Figure 2.6 4
Common Electrolytes (2.6)
• Electrolytes in proper proportion are critical to life
• Changes in electrolyte levels disturb almost every
vital function
• For example, potassium
• Low levels – paralyze muscles
• High levels – weak, irregular heartbeat
• Electrolyte levels are regulated by the kidneys,
digestive tract, and skeletal system
© 2013 Pearson Education, Inc.
© 2013 Pearson Education, Inc.
Figure 2.6 5
Colloid and Suspension (2.6)
• Colloid
• Solution containing dispersed proteins or other large
molecules
• Remain in solution indefinitely
• For example, liquid Jell-O
• Suspension
• Solution containing larger particles
• Particles will settle out if undisturbed
• For example, whole blood
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Hydrophobic (2.6)
• Molecules with few or no polar covalent bonds
• No positive/negative pole
• Nonpolar
• No hydration spheres form (no poles for the water molecules to be
attracted to)
• Molecules do not dissolve
• Hydrophobic
• Molecules that do not readily interact with water
© 2013 Pearson Education, Inc.
Colloid solution contains dispersed proteins or other large molecules
Fats and oils
Hydrophobic
molecules
Protein
Proteins held
in solution
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Figure 2.6 6
Module 2.6 Review
a. Define electrolytes.
b. Distinguish between hydrophilic and hydrophobic
molecules.
c. Explain how the ionic compound sodium chloride
dissolves in water.
© 2013 Pearson Education, Inc.
Hydrogen and Hydroxide Ions (2.7)
• Water (H2O) can dissociate into hydrogen ions
(H+) and hydroxide ions (OH–)
• Hydrogen ions
• Extremely reactive in solution
• Can break chemical bonds
• Can disrupt cell and tissue function
• Body has to regulate concentration to maintain life
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Water molecules dissociate into H+ and OH-
© 2013 Pearson Education, Inc.
Figure 2.7 1
pH Scale (2.7)
• pH
• Measure of hydrogen ion concentration in body fluids
• Ranges from 0 to 14
• Change in one unit is tenfold change in H+ ion
concentration
• 7 is neutral
• Below 7 is acidic
• Contains more hydrogen ions than hydroxide ions
• Above 7 is basic or alkaline
• Contains more hydroxide ions than hydrogen ions
© 2013 Pearson Education, Inc.
pH scale
Blood pH normally ranges from 7.35 to
7.45.
Urine
Beer,
vinegar,
wine, Tomatoes,
pickles grapes
Stomach
hydrochloric
acid
Saliva,
milk
Extremely Increasing concentration of H+
Decreasing concentration of OH–
acidic
pH 0
2
1
3
4
pH below 7.0 is acidic.
© 2013 Pearson Education, Inc.
5
6
Pure Seawater Household
bleach
water Eggs
Neutral
7
Oven
cleaner
Household
ammonia
Increasing concentration of OH+ Extremely
Decreasing concentration of H–
basic
8
pH of 7.0 is neutral.
9
10
11
12
13
pH above 7.0 is alkaline.
Figure 2.7 2
14
pH in the Blood (2.7)
• Normal pH of blood is 7.35 to 7.45
• Outside this range damages cells and tissues by:
• Breaking chemical bonds
• Changing shapes of proteins
• Altering cellular functions
• Acidosis
• Blood pH below 7.35
• pH below 7.0 causes coma
• Alkalosis
• Blood pH above 7.45
• pH above 7.8 causes uncontrollable, sustained skeletal muscle
contraction
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Acids (2.7)
• Acid
• Any solute that dissociates in solution and releases
hydrogen ions
• Lowering pH
• Also called proton donors
• Strong acid
• Dissociates completely in solution
• Reaction occurs in one direction
• For example, hydrochloric acid (HCl)
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An acid is any solute that dissociates in solution and releases hydrogen ions
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Figure 2.7 3
Bases (2.7)
• Base
• A solute that removes hydrogen ions from a solution
• Acts as a proton acceptor
• May release a hydroxide ion
• Strong base
• Dissociates completely in solution
• For example, sodium hydroxide (NaOH)
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A base is a solute that removes hydrogen ions from a solution
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Figure 2.7 4
Weak Acids and Weak Bases (2.7)
• Do not dissociate completely
• Have less of an impact on pH than strong acids
and bases
• For example, carbonic acid (H2CO3) dissociates
into a hydrogen ion and a bicarbonate ion
(HCO3–)
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Weak acids and weak bases don't dissociate completely
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Figure 2.7 5
Salts (2.7)
• Ionic compound composed of any cation except
hydrogen and any anion except hydroxide
• Held together by ionic bonds
• Many dissociate completely in water, releasing
cations and anions
• For example, sodium chloride (NaCl)
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A salt is an ionic compound consisting of any cation except a hydrogen ion and any anion except a hydroxide ion
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Figure 2.7 6
Buffers and Buffer Systems (2.7)
• Buffers
• Compounds that stabilize the pH of a solution by
removing or replacing hydrogen ions
• Buffer systems
• Help maintain pH within normal limits
• Involve a weak acid and its related salt
• For example, the carbonic acid–bicarbonate buffer
system
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Module 2.7 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?
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Metabolites and Nutrients (Section 4)
• Metabolites
• All the molecules that can be synthesized or broken down by chemical
reactions inside our bodies
• Nutrients
• Essential metabolites normally obtained from the diet
• Both categorized as:
• Organic compounds
• Always contain elements carbon and hydrogen and often oxygen
• Carbon chains held together by covalent bonds
• Inorganic compounds
• Usually do not have carbon and hydrogen
• Molecules held together by ionic bonds
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Organic and inorganic compounds
Nutrients and metabolites
include
Organic compounds
Organic compounds always contain carbon and
hydrogen as their main structural ingredients.
Examples: sugars, fats, proteins, and nucleic
acids (RNA, DNA), which are produced by living
organisms.
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Inorganic compounds
Inorganic compounds generally do not
contain carbon and hydrogen as their
main structural ingredients.
Examples: carbon dioxide, oxygen,
water, acids, bases, and salts.
Figure 2 Section 4
Carbohydrates (2.8)
• Organic molecules containing carbon, hydrogen,
and oxygen
• Ratio of these elements is 1:2:1
• Usually account for less than 1.5 percent total
body weight
• Most important function is as source of energy
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Figure 2.8 2
Monosaccharides (2.8)
• Simple sugars
• Most basic unit of a carbohydrate
• Contain from 3 to 7 carbon atoms
• Triose (3 carbon)
• Tetrose (4 carbon)
• Pentose (5 carbon)
• Hexose (6 carbon)
• Heptose (7 carbon)
• Glucose (a hexose)
• Most important metabolic "fuel" in the body
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Glucose molecular structure
Atoms in glucose
may form in a
straight chain or
ring.
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Figure 2.8 1
Three-Dimensional Structure (2.8)
• Determines molecule's fate or function
• Molecules may have same molecular formula but
different structures
• For example, fructose and glucose
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Three-dimensional structure determines fate or function of an organic molecule
Glucose
Fructose
Fructose has the same chemical formula as glucose, but
a different arrangement of atoms.
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Figure 2.8 2
Disaccharides (2.8)
• Formed when two monosaccharides join together
• One example is sucrose
• Very soluble in water
• Formed by dehydration synthesis
• Joining of molecules by the removal of a water molecule
• May be broken into monosaccharides by
hydrolysis
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Dehydration synthesis joins two monosaccharides forming a disaccharide
DEHYDRATION
SYNTHESIS
Glucose
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Fructose
Sucrose
Figure 2.8 3
Hydrolysis breaks down a disaccharide into monosaccharides
HYDROLYSIS
Glucose
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Fructose
Sucrose
Figure 2.8 4
Polysaccharides (2.8)
• Complex carbohydrates formed from multiple
monosaccharides or disaccharides
• Examples include:
• Starch
• Large polysaccharides formed by plants from glucose
• Found in potatoes and grains
• Major dietary source of energy
• Glycogen
• Large polysaccharide formed by animals from glucose
• Muscle cells make and store glycogen
• Broken down when high demand for glucose
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Glycogen is a polysaccharide
Glucose
molecules
Glycogen has many side branches formed by
chains of glucose.
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Figure 2.8 5
Module 2.8 Review
a. List the three structural classes of carbohydrates,
and give an example of each.
b. Cite the C:H:O ratio in carbohydrates and
describe their major functions in the body.
c. Predict the reactants and the type of chemical
reaction involved when muscle cells make and
store glycogen.
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Lipids (2.9)
• Contain carbon, hydrogen, and oxygen
• Carbon-to-hydrogen ratio 1:2
• Contain much less oxygen than carbohydrates
• May contain small quantities of phosphorus,
nitrogen, sulfur
• Include:
• Fats
• Oils
• Waxes
• Most insoluble in water
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Figure 2.9
Fatty Acids (2.9)
• Long carbon chains with hydrogen atoms attached
• One end (the head) has a carboxyl group:
–COOH
• Carbon chain known as the hydrocarbon tail
• Hydrophobic area
• The longer the tail, the lower the solubility
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Fatty acids are long carbon chains with hydrogen atoms attached
Carboxyl group
Hydrocarbon tail
of the fatty acid
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Lauric acid (C12H24O2)
Figure 2.9 1
Saturated Fatty Acid (2.9)
• Each carbon atom in tail
• Has four single covalent bonds
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Saturated fatty acid
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Figure 2.9 2
Unsaturated Fatty Acid (2.9)
• One or more of the covalent bonds between carbon atoms
is a double covalent bond
• Those carbon atoms will bind only 1 hydrogen atom
instead of 2
• Changes the shape of the tail
• Changes the way the body metabolizes the fatty acid
• Monounsaturated fatty acid
• Contains only one double bond
• Polyunsaturated fatty acid
• Contains multiple double bonds
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Unsaturated fatty acid
Double
covalent
bond
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Figure 2.9 3
Fats (2.9)
• Fatty acids can be attached to glycerol
• Triglycerides
• Three fatty acid chains
• One glycerol
• Most common fats in the body
• Monoglyceride
• Glycerol + one fatty acid
• Diglyceride
• Glycerol + two fatty acids
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Fatty acids can be attached to glycerol
Glycerol
Fatty acids
Fatty Acid 1
Saturated
Fatty Acid 2
Saturated
Fatty Acid 3
Unsaturated
HYDROLYSIS
Dehydration synthesis can produce
monoglycerides, diglycerides, or
triglylcerides. Hydrolysis breaks
glycerides into fatty acids and glycerol.
DEHYDRATION
SYNTHESIS
Triglyceride
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Figure 2.9 4
Lipids in the Body (2.9)
• Lipids
• Form essential structural components of all cells
• Important as energy reserves
• Provide roughly twice as much energy as carbohydrates
• Account for 12–18 percent total body weight in men
• Account for 18–24 percent total body weight in women
• Several essential fatty acids for lipid synthesis must be
obtained through diet
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Module 2.9 Review
a. Describe lipids.
b. Compare the structures of saturated and
unsaturated fatty acids.
c. In the hydrolysis of a triglyceride, what are the
reactants and the products?
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Structural Lipids (2.10)
• Include:
• Cholesterol
• Phospholipids
• Glycolipids
• Help form and maintain outer cell membrane and
intracellular membranes
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Structural lipids help form and maintain cell membranes
Glycolipid
Phospholipid
Plasma
membrane
Cholesterol
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Figure 2.10 1
Steroids (2.10)
• Large lipid molecules
• Share distinctive carbon-ring framework
• Differ in functional groups attached to the ring
• Cholesterol
• Found in animal plasma membranes
• Needed to maintain membrane
• Also needed for cell growth and division
• Cortisol
• Made in adrenal cortex
• Important in regulating tissue metabolism
• Estrogen and testosterone
• Involved in regulation of sexual function
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Steroids share a carbon-ring framework
Cholesterol
Cortisol
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Estrogen
Testosterone
Figure 2.10 2
Phospholipids and Glycolipids (2.10)
• Both are synthesized by our cells
• Both are diglycerides (two fatty acid tails attached
to one glycerol)
• Phospholipids
• Phosphate group links diglyceride to nonlipid group
• Glycolipids
• Carbohydrate attached to a diglyceride
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Phospholipids and glycolipids are structurally related
Nonlipid group
Carbohydrate
Phosphate group
Glycerol
Glycerol
Fatty
acids
In a phospholipid, a phosphate group
links a diglyceride to a nonlipid group.
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Fatty
acids
In a glycolipid, a carbohydrate
is attached to a diglyceride.
Figure 2.10 3
Micelles (2.10)
• Hydrophobic tails move away from water
• Hydrophilic heads move toward water
• Micelles
• Droplet formed by phospholipids and glycolipids
• Formed as digestive tract breaks down food
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Hydrophobic tails and hydrophilic heads
Hydrophilic
heads
Hydrophobic
tails
Phospholipid
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Glycolipid
Figure 2.10 4
Micelle structure
Water
Phospholipid
Glycolipid
Micelles
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Figure 2.10 4
Module 2.10 Review
a. Why is cholesterol necessary in the body?
b. Describe the basic functions of steroids,
phospholipids, and glycolipids.
c. Describe the orientations of phospholipids and
glycolipids when they form a micelle.
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Proteins Formed from Amino Acids (2.11)
• Proteins
• Most abundant organic component in human body
• Account for 20 percent of total body weight
• All contain carbon, hydrogen, oxygen, and nitrogen
• May contain small quantities sulfur and phosphorus
• Composed of long chains of amino acids
• 20 different amino acids found in the body
• Typical protein contains 1000 amino acids
• Three-dimensional shape determines functional properties
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Basic structure of an amino acid
Amino group
Central carbon
Carboxyl group
R group (side chain of
variable structure)
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Figure 2.11 1
Protein Formation (2.11)
• Amino acids
• Linked with covalent bonds through dehydration synthesis
• Bonds are called peptide bonds
• Peptides
• Molecules composed of amino acids held together by peptide
bonds
• Dipeptide
• Two amino acids
• Dipeptides can be split into amino acids by hydrolysis
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Dehydration synthesis links two amino acids together
Glycine (gly)
Alanine (ala)
1
Adjacent amino
acids can be
linked together by
a covalent bond
that connects the
carboxyl group of
one amino acid to
the amino group of
another.
3
Peptide bonds can
be broken through
hydrolysis.
2
Peptide bonds form between amino acids.
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Figure 2.11 2
Polypeptides and Proteins (2.11)
• Polypeptide
• More than two amino acids linked by peptide bonds
• Proteins
• Polypeptide chains of more than 100 amino acids
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Four Levels of Protein Structure – Primary and
Secondary (2.11)
• Primary structure
• Sequence of amino acids along length of single
polypeptide
• Secondary structure
• Results from bonds formed between atoms at different
parts of the polypeptide chain
• May create simple spiral (alpha-helix) or flat pleated
sheet
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Four Levels of Protein Structure – Tertiary
(2.11)
• Tertiary structure
• Coiling and folding gives three-dimensional shape
• Results from interactions between polypeptide chain
and surrounding water molecules
• Some contribution from interactions between R groups
of amino acids in protein
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Four Levels of Protein Structure – Tertiary
(2.11)
• Quaternary structure
• Results from interaction between individual polypeptide
chains
• Hemoglobin, a protein
• Contains four polypeptide subunits that form a globular protein
• Keratin and collagen, proteins
• Three linear subunits intertwine, forming fibrous protein
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Levels of protein structure
Primary Structure of a Protein
A1
A2
A4
A3
A5
A6
A7
A8
Secondary Structure
Hydrogen
bond
A2
A1
A5
A1
A2
A3
A4
A5
Hydrogen
bond
A6
A3
A9
A7 A9
or
A9
A8
A7
A6
A11
A12
A13
A14
A10
Pleated sheet
Alpha-helix
Tertiary Structure
Quaternary Structure
or
Hemoglobin
(globular protein)
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Keratin
(fibrous protein)
Figure 2.11 3
Denaturation (2.11)
• Increasing temperature changes protein shape
• Denaturation
• Change in protein tertiary or quaternary structure
• Protein no longer functions
• High body temperature (above 43°C or 110°F)
• Fatal due to denaturation of structural proteins
• Irreparable damage to organs and organ systems
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Module 2.11 Review
a. Describe proteins.
b. What kind of bond forms during the dehydration
synthesis of two amino acids?
c. How does boiling a protein affect its structural
and functional properties?
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Enzyme Function – Substrates and Active Sites
(2.12)
• Substrates
• Reactants in an enzymatic reaction
• Active site
• Specific region of the enzyme where substrates must
bind
• Substrates fit into this site like a key in a lock
• Site shape determined by tertiary or quaternary
structure of enzyme
• Binds only to substrates with particular shapes and
charges
• Demonstrating specificity
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Enzyme-Substrate Complex (2.12)
• Formed when substrate binds to active site on enzyme
• Multiple enzymes in each cell
• Each enzyme is active under its own set of conditions
• Enzyme activation or inactivation method of short-term
control over reaction rates and pathways
• Enzymes change shape when substrates bind
• Change is temporary and reversible
• Promotes formation of a product
• Completed product detaches from active site
• Enzyme able to repeat process
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Enzyme action on a substrate
S1
Substrates
S2
Active site
S1
S2
Enzymesubstrate
complex
S1
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S2
Figure 2.12 1 – 4
Module 2.12 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.
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High-Energy Compounds (2.13)
• Some enzymatic reactions require an energy
donor
• High-energy compound
• Contains high-energy bonds
• Covalent bonds that release energy when broken
• Adenosine triphosphate (ATP)
• Most common high-energy compound
• Provides energy to power vital functions
• Contraction of muscles
• Synthesis of proteins, carbohydrates, and lipids
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Formation of ATP (2.13)
• Begins with adenosine
• Composed of adenine and ribose
• Adenosine monophosphate (AMP)
• Adenosine bound to a single phosphate
• Adenosine diphosphate (ADP)
• Adenosine monophosphate with a second high-energy bond to a
phosphate (total 2 phosphates)
• Adenosine triphosphate (ATP)
• Adenosine diphosphate with yet another high-energy bond to a
phosphate (total 3 phosphates)
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The formation of adenosine triphosphate (ATP)
Adenosine triphosphate (ATP)
Adenosine diphosphate (ADP)
Adenosine monophosphate (AMP)
Adenosine
Adenine
Phosphate
groups
Ribose
High-energy
bonds
Adenosine + phosphate =
adenosine monophosphate (AMP)
Adenosine monophosphate
+ phosphate = adenosine
diphosphate (ADP)
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Adenosine diphosphate +
phosphate = adenosine
triphosphate (ATP)
Figure 2.13 1
ATP Formation Is Reversible (2.13)
• ATP is synthesized in one location
• Energy stored in ATP is released when broken
down to ADP in another location
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The formation of ATP from ADP is a reversible reaction
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Figure 2.13 2
Module 2.13 Review
a. Where do cells obtain the energy needed for their
vital functions?
b. Describe ATP.
c. Compare AMP with ADP.
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Nucleic Acids (2.14)
• Organic molecules
• Composed of carbon, hydrogen, oxygen, nitrogen,
and phosphorus
• Divided into two classes
• Deoxyribonucleic acid (DNA)
• Ribonucleic acid (RNA)
• Store and transfer information essential to protein
synthesis
• Composed of long chains of nucleotides
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Nucleotides (2.14)
• Composed of:
• Phosphate group
• Sugar
• Nitrogenous base
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Typical nucleotide structure
Sugar
Phosphate
group
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Nitrogenous
base
Figure 2.14 1
Nitrogenous bases
Nitrogenous bases
Adenine and guanine are found in both DNA and
RNA.
Adenine
Guanine
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)
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Figure 2.14 1
Nucleotide Backbone (2.14)
• Phosphate and sugar of adjacent nucleotides
bound together by dehydration synthesis
• Produces a chain of sugar-to-phosphate-to-sugar
sequence with nitrogenous bases projecting to
one side
• Sequence of nitrogenous bases carries
information
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Nucleotide phosphate and sugar are joined by dehydration synthesis
DEHYDRATION
SYNTHESIS
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Figure 2.14 2
DNA Structure (2.14)
• DNA molecule
• Consists of pair of nucleotide chains
• Called complementary strands
• Strands twist around each other to form a double helix
(like a spiral staircase)
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DNA Structure (2.14)
• DNA molecule
• Opposing nitrogenous bases held together by hydrogen
bonds
• Complementary base pairs due to shapes of the bases
• Adenine can bind only with thymine (A-T)
• Cytosine can bind only with guanine (C-G)
PLAY
Protein Synthesis: DNA Moleclue
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DNA molecular structure
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
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Guanine
Figure 2.14 3
RNA Structure (2.14)
• RNA molecule
• Single chain of nucleotides
• Shape and function depend on order of nucleotides and
interactions between them
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RNA molecular structure
Single-strand
sugar-phosphate
backbone
Nitrogenous
bases
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Figure 2.14 4
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Figure 2.14
Module 2.14 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?
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