Download Hydrogen bond

TEST REVIEW – EXAM 2
CHAPTER 2
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Terms to Know
• Molecule
• Compound
• Atom
• Mixture
• Element
• Suspensions
• Neutron
• Solutions
• Cation
• Nonpolar covalent bond
• Hydrogen bond
• Ionic bond
• Polar covalent bond
• Colloids
• Chemical energy
• Radiant energy
• Electrical energy
• Mechanical energy
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Matter
•
Anything that has mass and occupies
space
•
States of matter:
1. Solid—definite shape and volume
2. Liquid—definite volume, changeable shape
3. Gas—changeable shape and volume
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Mass and Weight
• the mass of an object is a fundamental property
of the object
• a numerical measure of its inertia
• measure of the amount of matter in the object.
• definitions of mass often seem circular because it
is such a fundamental quantity that it is hard to
define in terms of something else
• the usual symbol for mass is m and its SI unit is
the kilogram
• the weight of an object is the force of gravity on
the object (w = mg)
6
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A. Weight
B. Matter
C. Mass
D. Energy
1. Can be measured only by its effects on matter.
2. Anything that occupies space and has mass.
3. Although a man who weighs 175 pounds on
Earth would be lighter on the moon and heavier
on Jupiter, his ________ would not be different.
4. Is a function of and varies with gravity.
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Energy Concepts
• What is energy?
•
The capacity to perform work
• What is the difference between potential and kinetic energy?
•
Stored vs. motion
• Energy is neither created nor destroyed but…
•
Converted from one form to another
•
This property is called the conservation of energy
• What is the usual way in which energy is “lost?”
•
Through heat
• What type of energy is heat?
•
Kinetic due to random motion of atoms
•
Heat is generated by friction (in this example between atoms and air)
• Heat is highly __________ energy and highest amount of _________.
•
Disordered, entropy
• Chemical energy is a form of ____________ energy.
•
Potential
• What is the primary form of chemical energy in living organisms?
•
ATP
• What is cellular respiration? What are the byproducts?
•
Conversion of glucose into ATP through reduction of oxygen forming water and carbon
dioxide
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Forms of Energy
• Chemical energy — stored in bonds of
chemical substances
• Electrical energy — results from movement of
charged particles
• Mechanical energy — directly involved in
moving matter
• Radiant or electromagnetic energy — exhibits
wavelike properties (i.e., visible light,
ultraviolet light, and X-rays)
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Composition of Matter
• Elements
• Cannot be broken down by ordinary chemical means
• Each has unique properties:
• Physical properties
• Are detectable with our senses or are
measurable
• Chemical properties
• How atoms interact (bond) with one another
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Composition of Matter
• Atoms
• Unique building blocks for each element
• Atomic symbol: one- or two-letter chemical
shorthand for each element
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Atomic Structure
• Neutrons
• No charge
• Mass = 1 atomic mass unit (amu)
• Protons
• Positive charge
• Mass = 1 amu
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Atomic Structure
• Determined by numbers of subatomic
particles
• Nucleus consists of neutrons and protons
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Atomic Structure
• Electrons
• Orbit nucleus
• Equal in number to protons in atom
• Negative charge
• 1/2000 the mass of a proton (0 amu)
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Identifying Elements
• Atoms of different elements contain different
numbers of subatomic particles
• Compare hydrogen, helium and lithium (next
slide)
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Proton
Neutron
Electron
Hydrogen (H)
(1p+; 0n0; 1e–)
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Helium (He)
(2p+; 2n0; 2e–)
Lithium (Li)
(3p+; 4n0; 3e–)
Figure 2.2
Identifying Elements
• Atomic number = number of protons in
nucleus
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Identifying Elements
• Mass number = mass of the protons and
neutrons
• Mass numbers of atoms of an element are not
all identical
• Isotopes are structural variations of elements
that differ in the number of neutrons they
contain
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Radioisotopes
• Valuable tools for biological research and
medicine
• Cause damage to living tissue:
• Useful against localized cancers
• Radon from uranium decay causes lung
cancer
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Proton
Neutron
Electron
Hydrogen (1H)
(1p+; 0n0; 1e–)
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Deuterium (2H)
(1p+; 1n0; 1e–)
Tritium (3H)
(1p+; 2n0; 1e–)
Figure 2.3
Chemically Inert Elements
• Stable and unreactive
• Outermost energy level fully occupied or
contains eight electrons
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(a)
Chemically inert elements
Outermost energy level (valence shell) complete
8e
2e
Helium (He)
(2p+; 2n0; 2e–)
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2e
Neon (Ne)
(10p+; 10n0; 10e–)
Figure 2.5a
Chemically Reactive Elements
• Outermost energy level not fully occupied by
electrons
• Tend to gain, lose, or share electrons (form
bonds) with other atoms to achieve stability
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(b)
Chemically reactive elements
Outermost energy level (valence shell) incomplete
1e
Hydrogen (H)
(1p+; 0n0; 1e–)
6e
2e
Oxygen (O)
(8p+; 8n0; 8e–)
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4e
2e
Carbon (C)
(6p+; 6n0; 6e–)
1e
8e
2e
Sodium (Na)
(11p+; 12n0; 11e–)
Figure 2.5b
Molecules and Compounds
• Most atoms combine chemically with other
atoms to form molecules and compounds
• Molecule — two or more atoms bonded
together (e.g., H2 or C6H12O6)
• Compound — two or more different kinds of
elements bonded together (e.g., C6H12O6)
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Mixtures vs. Compounds
• Mixtures
• No chemical bonding between components
• Can be separated physically, such as by
straining or filtering
• Heterogeneous or homogeneous
• Compounds
• Can be separated only by breaking bonds
• All are homogeneous
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Mixtures
• Most matter exists as mixtures
• Two or more components physically
intermixed
• Three types of mixtures
• Solutions
• Colloids
• Suspensions
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Solutions
• Homogeneous mixtures
• Usually transparent, e.g., atmospheric air or
seawater
• Solvent
• Present in greatest amount, usually a liquid
• Solute(s)
• Present in smaller amounts
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Colloids and Suspensions
• Colloids (emulsions)
• Heterogeneous translucent mixtures, e.g.,
cytosol
• Large solute particles that do not settle out
• Undergo sol-gel transformations
• Suspensions:
• Heterogeneous mixtures (blood)
• Large visible solutes tend to settle out
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Solution
Colloid
Suspension
Solute particles are very
tiny, do not settle out or
scatter light.
Solute particles are larger
than in a solution and scatter
light; do not settle out.
Solute particles are very
large, settle out, and may
scatter light.
Solute
particles
Solute
particles
Solute
particles
Example
Example
Example
Mineral water
Gelatin
Blood
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Figure 2.4
Mixtures vs. Compounds
• Mixtures
• No chemical bonding between components
• Can be separated physically, such as by
straining or filtering
• Heterogeneous or homogeneous
• Compounds
• Can be separated only by breaking bonds
• All are homogeneous
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• Heterogeneous, will not settle.
• Heterogeneous, will settle.
• Homogeneous, will not settle.
• Will not scatter light.
A)Suspensions B) Solutions
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C) Colloids
• Nonpolar covalent bond
• Hydrogen bond
• Ionic bond
• Polar covalent bond
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Chemical Bonds
• Electrons occupy up to seven electron shells
(energy levels) around nucleus
• Octet rule: Except for the first shell which is
full with two electrons, atoms interact in a
manner to have eight electrons in their
outermost energy level (valence shell)
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Figure 2.9
Ionic Bonds
• Ions are formed by transfer of valence shell
electrons between atoms
• Anions (– charge) have gained one or more
electrons
• Cations (+ charge) have lost one or more
electrons
• Attraction of opposite charges results in an
ionic bond
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Formation of an Ionic Bond
• Ionic compounds form crystals instead of
individual molecules
• NaCl (sodium chloride)
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Sodium atom (Na)
(11p+; 12n0; 11e–)
Chlorine atom (Cl)
(17p+; 18n0; 17e–)
+
–
Sodium ion (Na+)
Chloride ion (Cl–)
Sodium chloride (NaCl)
(a) Sodium gains stability by losing one electron, and
chlorine becomes stable by gaining one electron.
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(b) After electron transfer, the oppositely
charged ions formed attract each other.
Figure 2.6a-b
CI–
Na+
(c) Large numbers of Na+ and Cl– ions
associate to form salt (NaCl) crystals.
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Figure 2.6c
Covalent Bonds
• Formed by sharing of two or more valence
shell electrons
• Allows each atom to fill its valence shell at
least part of the time
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Reacting atoms
Resulting molecules
+
Molecule of
Hydrogen
Carbon
methane gas (CH4)
atoms
atom
(a) Formation of four single covalent bonds:
carbon shares four electron pairs with four
hydrogen atoms.
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or
Structural
formula
shows
single
bonds.
Figure 2.7a
Covalent Bonds
• Sharing of electrons may be equal or unequal
• Equal sharing produces electrically balanced
nonpolar molecules
• CO2
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Reacting atoms
Resulting molecules
+
Oxygen
atom
or
Oxygen
atom
Molecule of
oxygen gas (O2)
(b) Formation of a double covalent bond: Two
oxygen atoms share two electron pairs.
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Structural
formula
shows
double
bond.
Figure 2.7b
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Figure 2.8a
Covalent Bonds
• Unequal sharing by atoms with different
electron-attracting abilities produces polar
molecules
• H2O
• Atoms with six or seven valence shell
electrons are electronegative, e.g., oxygen
• Atoms with one or two valence shell
electrons are electropositive, e.g., sodium
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Figure 2.8b
Hydrogen bonds
•
The bonds of a water molecule represent ________ _______ type of bond. Also known as a ________.
•
•
Oxygen has a greater affinity for the electrons and is therefore more _____________. Whereas, hydrogen
has a lesser attraction for electrons is more _____________.
•
•
•
Negative, positive
The attraction between the negative oxygen end of one water compound to the positive hydrogen end of
another water represents a ___________ bond.
•
•
Electronegative, electropositive
The oxygen end of the molecule is therefore slightly more _________ and the hydrogen ends are slightly
more _________.
•
•
Polar covalent, dipole
Hydrogen
Hydrogen bonds are strong bonds. (T/F)
•
False
•
They are easily broken
Hydrogen bonds may inter- or intramolecular. (T/F)
•
True
• The unique properties of water are attributable to hydrogen bonds. Some of the properties
include….
•
Cohesion, high boiling point, why ice floats, high heat of vaporization, high heat capacity
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+
–
Hydrogen bond
(indicated by
dotted line)
+
+
–
–
–
+
+
+
–
(a) The slightly positive ends (+) of the water
molecules become aligned with the slightly
negative ends (–) of other water molecules.
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Figure 2.10a
Chemical Reactions
• Occur when chemical bonds are formed,
rearranged, or broken
• Represented as chemical equations
• Chemical equations contain:
• Molecular formula for each reactant and
product
• Relative amounts of reactants and products,
which should balance
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Patterns of Chemical Reactions
• Synthesis (combination) reactions
• Decomposition reactions
• Exchange reactions
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Synthesis Reactions
• A + B  AB
• Always involve bond formation
• Anabolic
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Dehydration Synthesis and Hydrolysis
• What is dehydration synthesis?
• Removal of a water molecule to form a new covalent
bond
• What is hydrolysis?
• The addition of a water molecule to break a covalent
bond
• What is anabolism?
• Forming new bonds to build something bigger.
Requires energy (endergonic)
• What is catabolism?
• Breaking bonds to make something smaller. Large
molecules down to subunits.
• Releases energy (exergonic).
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Decomposition Reactions
• AB  A + B
• Reverse synthesis reactions
• Involve breaking of bonds
• Catabolic
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Oxidation-Reduction (Redox) Reactions
• Decomposition reactions: Reactions in which
fuel is broken down for energy
• Also called exchange reactions because
electrons are exchanged or shared differently
• Electron donors lose electrons and are
oxidized
• Electron acceptors receive electrons and
become reduced
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Chemical Reactions
• All chemical reactions are either exergonic or
endergonic
• Exergonic reactions — release energy
• Catabolic reactions
• Endergonic reactions — products contain
more potential energy than did reactants
• Anabolic reactions
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Chemical Reactions
• All chemical reactions are theoretically reversible
• A + B  AB
• AB  A + B
• Chemical equilibrium occurs if neither a forward nor
reverse reaction is dominant
• Many biological reactions are essentially irreversible
due to
• Energy requirements
• Removal of products
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Rate of Chemical Reactions
• Rate of reaction is influenced by:
•  temperature   rate
•  particle size   rate
•  concentration of reactant   rate
• Catalysts:  rate without being chemically
changed
• Enzymes are biological catalysts
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Classes of Compounds
• Inorganic compounds
• Water, salts, and many acids and bases
• Do not contain carbon
• Organic compounds
• Carbohydrates, fats, proteins, and
nucleic acids
• Contain carbon, usually large, and are
covalently bonded
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Water
• 60%–80% of the volume of living cells
• Most important inorganic compound in living
organisms because of its properties
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Properties of Water
• High heat capacity
• Absorbs and releases heat with little
temperature change
• Prevents sudden changes in temperature
• High heat of vaporization
• Evaporation requires large amounts of heat
• Useful cooling mechanism
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Properties of Water
• Polar solvent properties
• Dissolves and dissociates ionic substances
• Forms hydration layers around large charged
molecules, e.g., proteins (colloid formation)
• Body’s major transport medium
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+
–
+
Water molecule
Salt crystal
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Ions in solution
Figure 2.12
Properties of Water
• Reactivity
• A necessary part of hydrolysis and dehydration
synthesis reactions
• Cushioning
• Protects certain organs from physical trauma,
e.g., cerebrospinal fluid
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Salts
• Ionic compounds that dissociate in water
• Contain cations other than H+ and anions
other than OH–
• Ions (electrolytes) conduct electrical currents
in solution
• Ions play specialized roles in body functions
(e.g., sodium, potassium, calcium, and iron)
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Acids and Bases
• Both are electrolytes
• Acids are proton (hydrogen ion) donors
(release H+ in solution)
• HCl  H+ + Cl–
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Acids and Bases
• Bases are proton acceptors (take up H+ from
solution)
• NaOH  Na+ + OH–
• OH– accepts an available proton (H+)
• OH– + H+  H2O
• Bicarbonate ion (HCO3–) and ammonia
(NH3) are important bases in the body
because of buffering properties
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Acid-Base Concentration
• Acid solutions contain [H+]
• As [H+] increases, acidity increases, pH
decreases
• Alkaline solutions contain bases (e.g., OH–)
• As [H+] decreases (or as [OH–] increases),
alkalinity increases, pH increases
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pH: Acid-Base Concentration
• pH = the negative logarithm of [H+] in
moles per liter
• Neutral solutions:
• Pure water is pH neutral (contains equal
numbers of H+ and OH–)
• pH of pure water = pH 7: [H+] = 10 –7 M
• All neutral solutions are pH 7
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pH: Acid-Base Concentration
• Acidic solutions
•  [H+],  pH
• Acidic pH: 0–6.99
• pH scale is logarithmic: a pH 5 solution has
10 times more H+ than a pH 6 solution
• Alkaline solutions
•  [H+],  pH
• Alkaline (basic) pH: 7.01–14
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Concentration
(moles/liter)
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Examples
[OH–]
[H+]
pH
100
10–14
14
1M Sodium
hydroxide (pH=14)
10–1
10–13
13
Oven cleaner, lye
(pH=13.5)
10–2
10–12
12
10–3
10–11
11
10–4
10–10
10
10–5
10–9
9
10–6
10–8
8
10–7
10–7
7 Neutral
10–8
10–6
6
10–9
10–5
5
10–10
10–4
4
10–11
10–3
3
10–12
10–2
2
10–13
10–1
1
10–14
100
0
Household ammonia
(pH=10.5–11.5)
Household bleach
(pH=9.5)
Egg white (pH=8)
Blood (pH=7.4)
Milk (pH=6.3–6.6)
Black coffee (pH=5)
Wine (pH=2.5–3.5)
Lemon juice; gastric
juice (pH=2)
1M Hydrochloric
acid (pH=0)
Figure 2.13
Acid-Base Homeostasis
• pH change interferes with cell function and
may damage living tissue
• Slight change in pH can be fatal
• pH is regulated by kidneys, lungs, and
buffers
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Buffers
• Mixture of compounds that resist pH changes
• Convert strong (completely dissociated) acids
or bases into weak (slightly dissociated) ones
• Carbonic acid-bicarbonate system
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Organic Compounds
• Contain carbon (except CO2 and CO, which
are inorganic)
• Unique to living systems
• Include carbohydrates, lipids, proteins, and
nucleic acids
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Organic Compounds
• Many are polymers — chains of similar units (monomers
or building blocks)
• Synthesized by dehydration synthesis
• Broken down by hydrolysis reactions
• How are polymers formed?
• By dehydration synthesis
• What reactions break down polymers into monomers?
• By hydrolysis
• What molecule is essential to this process?
• H2O
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(a)
Dehydration synthesis
Monomers are joined by removal of OH from one monomer
and removal of H from the other at the site of bond formation.
Monomer 1
+
Monomer 2
Monomers linked by covalent bond
(b)
Hydrolysis
Monomers are released by the addition of a water molecule, adding OH to one monomer and H to the other.
+
Monomer 1
Monomer 2
Monomers linked by covalent bond
(c)
Example reactions
Dehydration synthesis of sucrose and its breakdown by hydrolysis
Water is
released
+
Water is
consumed
Glucose
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Fructose
Sucrose
Figure 2.14
Carbohydrates
• Sugars and starches
• Contain C, H, and O [(CH20)n]
• Three classes
• Monosaccharides
• Disaccharides
• Polysaccharides
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Carbohydrates
• Functions
• Major source of cellular fuel (e.g., glucose)
• Structural molecules (e.g., ribose sugar in
RNA)
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Monosaccharides
• Simple sugars containing three to seven C
atoms
• (CH20)n n = 3 – 7
• C3H6O3
• C6H12O6
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(a) Monosaccharides
Monomers of carbohydrates
Example
Example
Hexose sugars (the hexoses shown
Pentose sugars
here are isomers)
Glucose
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Fructose
Galactose
Deoxyribose
Ribose
Figure 2.15a
• Glycogen is animals main storage form of glucose. Found in high concentrations in the liver and muscles.
• Starch is plants main storage form of glucose.
• Cellulose is a key structural molecule in plants. Not digestible by humans.
(c) Polysaccharides
Long branching chains (polymers) of linked monosaccharides
Example
This polysaccharide is a simplified representation of
glycogen, a polysaccharide formed from glucose units.
Glycogen
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Figure 2.15c
Lipids
• Contain C, H, O (less than in carbohydrates),
and sometimes P
• Insoluble in water
• Main types:
• Neutral fats or triglycerides
• Phospholipids
• Steroids
• Eicosanoids
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Triglycerides
• Neutral fats — solid fats and liquid oils
• Composed of three fatty acids bonded to a
glycerol molecule
• Main functions
• Energy storage
• Insulation
• Protection
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(a) Triglyceride formation
Three fatty acid chains are bound to glycerol by
dehydration synthesis
+
Glycerol
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3 fatty acid chains
Triglyceride,
or neutral fat
3 water
molecules
Figure 2.16a
Saturation of Fatty Acids
• Saturated fatty acids
• Single bonds between C atoms; maximum
number of H
• Solid animal fats, e.g., butter
• Unsaturated fatty acids
• One or more double bonds between C atoms
• Reduced number of H atoms
• Plant oils, e.g., olive oil
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Phospholipids
• Modified triglycerides:
• Glycerol + two fatty acids and a phosphorus
(P)-containing group
• “Head” and “tail” regions have different
properties (amphipathic)
• Hydrophilic head
• Hydrophobic tail
• Important in cell membrane structure
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(b) “Typical” structure of a phospholipid molecule
Two fatty acid chains and a phosphorus-containing group are
attached to the glycerol backbone.
Example
Phosphatidylcholine
Polar
“head”
Nonpolar
“tail”
(schematic
phospholipid)
Phosphoruscontaining
group (polar
“head”)
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Glycerol
backbone
2 fatty acid chains
(nonpolar “tail”)
Figure 2.16b
Steroids
• Steroids — interlocking four-ring structure
• Cholesterol, vitamin D, steroid hormones, and
bile salts
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(c)
Simplified structure of a steroid
Four interlocking hydrocarbon rings form a steroid.
Example
Cholesterol (cholesterol is the
basis for all steroids formed in the body)
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Figure 2.16c
Other Lipids in the Body
• Other fat-soluble vitamins
• Vitamins A, D, E, and K
• Lipoproteins
• Transport fats in the blood
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Proteins
• Polymers of amino acids (20 types)
• Joined by peptide bonds
• Contain C, H, O, N, and sometimes S and P
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Amine
group
Acid
group
(a) Generalized
structure of all
amino acids.
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(b) Glycine
is the simplest
amino acid.
(c) Aspartic acid
(an acidic amino acid)
has an acid group
(—COOH) in the
R group.
(d) Lysine
(a basic amino acid)
has an amine group
(–NH2) in the R group.
(e) Cysteine
(a basic amino acid)
has a sulfhydryl (–SH)
group in the R group,
which suggests that
this amino acid is likely
to participate in
intramolecular bonding.
Figure 2.17
Dehydration synthesis:
The acid group of one
amino acid is bonded to
the amine group of the
next, with loss of a water
molecule.
Peptide
bond
+
Amino acid
Amino acid
Dipeptide
Hydrolysis: Peptide
bonds linking amino
acids together are
broken when water is
added to the bond.
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Figure 2.18
Amino acid
Amino acid
Amino acid
Amino acid
Amino acid
(a) Primary structure:
The sequence of amino acids forms the polypeptide chain.
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Figure 2.19a
a-Helix: The primary chain is coiled
to form a spiral structure, which is
stabilized by hydrogen bonds.
b-Sheet: The primary chain “zig-zags” back
and forth forming a “pleated” sheet. Adjacent
strands are held together by hydrogen bonds.
(b) Secondary structure:
The primary chain forms spirals (a-helices) and sheets (b-sheets).
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Figure 2.19b
Tertiary structure of prealbumin
(transthyretin), a protein that
transports the thyroid hormone
thyroxine in serum and cerebrospinal fluid.
(c) Tertiary structure:
Superimposed on secondary structure. a-Helices and/or b-sheets are
folded up to form a compact globular molecule held together by
intramolecular bonds.
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Figure 2.19c
Quaternary structure of
a functional prealbumin
molecule. Two identical
prealbumin subunits
join head to tail to form
the dimer.
(d) Quaternary structure:
Two or more polypeptide chains, each with its own tertiary structure,
combine to form a functional protein.
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Figure 2.19d
Fibrous and Globular Proteins
• Fibrous (structural) proteins
• Strandlike, water insoluble, and stable
• Examples: keratin, elastin, collagen, and
certain contractile fibers
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Fibrous and Globular Proteins
• Globular (functional) proteins
• Compact, spherical, water-soluble and
sensitive to environmental changes
• Specific functional regions (active sites)
• Examples: antibodies, hormones, molecular
chaperones, and enzymes
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Protein Denaturation
• Shape change and disruption of active sites
due to environmental changes (e.g.,
decreased pH or increased temperature)
• Reversible in most cases, if normal conditions
are restored
• Irreversible if extreme changes damage the
structure beyond repair (e.g., cooking an egg)
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Enzymes
WITHOUT ENZYME
WITH ENZYME
Activation
energy
required
Less activation
energy required
Reactants
Reactants
Product
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Product
Figure 2.20
Substrates (S)
e.g., amino acids
+
Product (P)
e.g., dipeptide
Energy is
absorbed;
bond is
formed.
Water is
released.
H2O
Peptide
bond
Active site
Enzyme (E)
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Enzyme-substrate
complex (E-S)
Enzyme (E)
1 Substrates bind
2
Internal
Product is
at active site.
rearrangements 3
released. Enzyme
Enzyme changes
leading to
returns to original
shape to hold
catalysis occur.
shape and is
substrates in
available to catalyze
proper position.
another reaction.
Figure 2.21, step 3
Enzymes
• What is an enzyme?
• Protein
• Biologic catalyst
• What is a catalyst
• Substance that speeds up a reaction
• What is Ea?
• Energy of activation
• Enzymes do what to a reaction?
• Lower energy of activation (heat, mechanical, chemical, etc)
• Speeds up rxn
• On what does an enzyme act?
• Its substrate
• Enzymes are __________ for their substrates?
• Specific
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Nucleic Acids
• DNA and RNA
• Largest molecules in the body
• Contain C, O, H, N, and P
• Building block = nucleotide, composed of Ncontaining base, a pentose sugar, and a
phosphate group
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Deoxyribonucleic Acid (DNA)
• Four bases:
• adenine (A), guanine (G), cytosine (C), and
thymine (T)
• Double-stranded helical molecule in the cell
nucleus
• Provides instructions for protein synthesis
• Replicates before cell division, ensuring
genetic continuity
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Phosphate
Sugar:
Deoxyribose
Base:
Adenine (A)
Thymine (T)
Adenine nucleotide
Sugar
Phosphate
Thymine nucleotide
Hydrogen
bond
(a)
Sugar-phosphate
backbone
Deoxyribose
sugar
Phosphate
Adenine (A)
Thymine (T)
Cytosine (C)
Guanine (G)
(b)
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(c) Computer-generated image of a DNA molecule
Figure 2.22
Ribonucleic Acid (RNA)
• Four bases:
• adenine (A), guanine (G), cytosine (C), and
uracil (U)
• Single-stranded molecule mostly active
outside the nucleus
• Three varieties of RNA carry out the DNA
orders for protein synthesis
• messenger RNA, transfer RNA, and ribosomal
RNA
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True / False
• Chemical properties are determined primarily by neutrons.
• No chemical bonding occurs between the components of a
mixture.
• Buffers resist abrupt and large changes in the pH of the body by
releasing or binding ions.
• All organic compounds contain carbon.
• A dipeptide can be broken into two amino acids by dehydration
synthesis.
• The lower the pH, the higher the hydrogen ion concentration.
• Covalent bonds are generally less stable than ionic bonds.
• Hydrogen bonds are comparatively strong bonds.
• The fact that no chemical bonding occurs between the
components of a mixture is the chief difference between mixtures
and compounds.
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True / False
• The pH of body fluids must remain fairly constant for the body to
maintain homeostasis
• A charged particle is generally called an ion.
• Isotopes differ from each other only in the number of electrons
contained.
• About 60% to 80% of the volume of most living cells consists of
organic compounds.
• Lipids are a poor source of stored energy.
• Current information theorizes that omega-3 fatty acids decrease
the risk of heart disease.
• Glucose is an example of a monosaccharide.
• A molecule consisting of one carbon atom and two oxygen atoms
is correctly written as CO2.
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Multiple Choice
• Choose the statement that is false or incorrect.
A) In chemical reactions, breaking old bonds requires energy and forming new bonds releases
energy.
B) Exergonic reactions release more energy than they absorb.
C) A key feature of the body’s metabolism is the almost exclusive use of exergonic reactions by
the body.
D) Endergonic reactions absorb more energy than they release.
• A chemical reaction in which bonds are broken is usually associated with ________.
A) a synthesis B) the consumption of energy C) the release of energy D) forming a larger molecule
• What happens in redox reactions?
A)
the reaction is always easily reversible
B) the electron acceptor is oxidized
B)
both decomposition and electron exchange occur D) the electron donor is reduced
• Choose the answer that best describes fibrous proteins
A) are usually called enzymes B) are very stable and insoluble in water
C) rarely exhibit secondary structure D) are cellular catalysts
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Multiple Choice
•
In liquid XYZ, you notice that light is scattered as it passes through. There is no precipitant in the
bottom of the beaker, though it has been sitting for several days. What type of liquid is this?
A) suspension B) solution C) mixture D) colloid
•
Atom X has 17 protons. How many electrons are in its valence shell?
•
A) 10 B) 5 C) 3 D) 7
•
Which protein types are vitally important to cell function in all types of stressful circumstances?
•
A) catalytic proteins B) molecular chaperones
•
C) regulatory proteins D) structural proteins
•
If atom X has an atomic number of 74 it would have which of the following?
A) 37 protons and 37 neutrons B) 37 protons and 37 electrons
C) 74 protons D) 37 electrons
•
What does the formula C6H12O6 mean?
A) There are 6 calcium, 12 hydrogen, and 6 oxygen atoms.
•
B) There are 12 hydrogen, 6 carbon, and 6 oxygen atoms.
•
C) The substance is a colloid.
•
D) The molecular weight is 24.
•
Two good examples of a colloid would be Jell-O® and ________.
•
A) cytosol B) blood C) urine D) toenails
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Multiple Choice
•
An atom with a valence of 3 may have a total of ________ electrons.
A) 3 B) 8 C) 17 D) 13
•
The chemical symbol O=O means ________.
•
A) zero equals zero
•
B) the atoms are double bonded
•
C) both atoms are bonded and have zero electrons in the outer orbit
•
D) this is an ionic bond with two shared electrons
•
What is a dipole?
A) a type of reaction B) a type of bond
C) an organic molecule D) a polar molecule
•
Amino acids joining together to make a peptide is a good example of a(n) ________ reaction.
A) decomposition B) reversible
C) exchange D) synthesis
•
Which of the following is not considered a factor in influencing a reaction?
A) time B) concentration C) particle size D) temperature
•
Which of the following is not an electrolyte?
A) NaOH B) HCl C) H2O D) Ca2CO3
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