Transcript Figure 5.15 The 20 amino acids of proteins

WATER
Chapter 3: Water and the
Fitness of the Environment
Water
• Cells are 70-90% water
• Three-fourths of Earth’s surface is
covered by water
• Water is the biological medium for
life on Earth
• Water must be present for life (as
we know it)
Figure 3.1 Hydrogen bonds between water molecules
Water’s polarity
• Polar – opposite ends of molecule
have opposite charges
– Opposing charges due to oxygen’s
electronegativity
• Oxygen has partial negative
• Hydrogens have partial positive
• Water forms H bonds (opposite
charges attract)
– ~15% water molecules in our body
are bonded to four partners
Four Properties of Water
Allowing for Life
1. Cohesion
2. Moderation of temperature
3. Insulation of bodies of water by
floating ice
4. The solvent of life
Figure 3.2x Trees
Figure 3.3 Walking on water
Cohesion
• Cohesion - H bonding keeps water
molecules close together
• Makes water a “structured” liquid
• Adhesion – the clinging of one substance to
another
– Ex. water sticks to sides of a glass
• Both cohesion and adhesion help water
move up from roots to plants to leaves
• Surface tension – measure of difficult it is
to stretch or break the surface of a liquid
– Water has a high surface tension due to
H bonds – almost like a thin film on the
surface.
Moderation of Temperature
• Heat – measure amount of total kinetic
energy (cal, kcal, or J)
– 1 cal = amount of heat needed to raise
the temp of 1g of water 1°C
– 1000 cal = 1 kcal
– 1 cal = 4.184 J
• Temperature – measures the intensity
of heat due to average kinetic energy of
molecules (°C)
• Specific heat – amount of heat that
must be absorbed or lost for 1 g to
change its temp by 1 °C
– Specific heat of water = 1 cal/g/°C
– Compared to most substances, water’s
specific heat is quite high.
– This means water will changes its temp
less when it absorbs or gives off heat.
– Why? H bonds need heat to break and
heat is released when bonds are formed.
– High specific heat allows large bodies of
water to absorb lots of heat in summer
without raising temp too high. In winter,
gradual cooling of water helps warm the
air.
– High specific heat helps stabilize ocean
temp to better support marine life.
• Heat of vaporization – quantity of heat a
liquid must absorb for 1 g to be converted
to gaseous state
• Water has a high heat of vaporization (580
cal heat needed to evaporate 1 g water at
25°C) because H bonds must be broken
before molecules can change to gas.
– Evaporative cooling – as liquid
evaporates, the surface cools because
the hottest molecules leave as gas and
cooler molecules are left behind
• Why sweat? Evaporative cooling in
progress!
• Why do we sweat more on humid days?
Insulation of Bodies of
Water by Floating Ice
• Water is more dense than ice. At 4°C
water is at its most dense state, then as it
cools to O°C, the molecules freeze. The H
bonds keep the water molecules slightly
apart (like a lattice) so air pockets form
within ice.
• Lakes, oceans etc. would freeze solid if ice
was more dense than water. During the
summer, only upper few inches of ocean
would thaw making life as we know it
impossible (not to mention ice skating and
hockey! )
Figure 3.5 The structure of ice (Layer 2)
Figure 3.5x1 Ice, water, and steam
Figure 3.6 Floating ice and the fitness of the environment
Figure 3.6x1 Floating ice and the fitness of the environment: ice fishing
Figure 3.6x2 Ice floats and frozen benzene sinks
The Solvent of Life
• Solvent - dissolving agent in a solution
• Solute – the substance that is
dissolved
• Aqueous solution – water is the
solvent
• Water can dissolve many substances,
but obviously not all!
– Hydrophilic – likes water
– Hydrophobic – repel water (nonpolar and
nonionic)
Figure 3.7 A crystal of table salt dissolving in water
Figure 3.8 A water-soluble protein
• Concentrations
– Molecular mass – sum of mass of atoms
(daltons)
• Molecular mass of CO2 = 12 + 16(2) = 44
daltons
– 1 mole = 6.02 x 1023 molecules
– 1 mole CO2 = 44 grams
– Molarity (M) = mole/liter
Figure 3.x2 Moles
Unnumbered Figure (page 47) Chemical reaction: hydrogen bond shift
pH
• H+ (protons) occasionally move from
one water molecules to another
(disassociation).
• If water loses a H+ then it becomes
OH- (hydroxide ion).
• If water gains a H+ then it becomes
H3O+ (hydronium ion).
• In pure water, the OH- and H+
concentrations are equal.
• Acid – increases H+ concentrations
• Base – decreases H+ concentrations
• For pure (neutral) water at 25°C:
– [OH-] [H+] = 10-14
– [H+] = 10-7
– [OH-] = 10-7
• If enough acid is added to increase the
[H+] to 10-4, then the [OH-] will decrease by
an equivalent amount or 10-10
• Because concentrations can vary by factors
of 100 trillion, scientists use a log scale.
• pH = -log [H+]
– For neutral water pH = -log 10-7 = -(-7) = 7
– A pH of 3 vs. ph of 6 is a 1000 fold difference
(10 fold for each step)
Figure 3.9 The pH of some aqueous solutions
Buffers
• Buffers – minimize changes in pH by being
able to release or take in H+
• Buffers keeps blood between 7 and 7.8
• The equation below shifts right to
decrease pH and left to increase pH
(bicarbonate buffer)
– H2CO3
Carbonic acid
HCO3- + H+
bicarbonate
Acid Rain
• Acid precipitation – below 5.6
• Caused primarily by increased levels of
sulfur and nitrogen oxides released from
the burning of fossil fuels
• Acid precip can damage lakes, streams, and
soil.
• Acid can make harmful heavy metals more
soluble in water.
Figure 3.10 The effects of acid precipitation on a forest
Figure 3.10x1 Pulp mill
Figure 3.10x2 Acid rain damage to statuary, 1908 & 1968
CARBON AND THE
MOLECULAR DIVERSITY
OF LIFE
CHAPTER 4
Figure 3.10x2 Acid rain damage to statuary, 1908 & 1968
ISOMERS
 Compounds
with the same chemical
formula but different structures
FUNCTIONAL GROUPS
 See
diagram of functional groups
Figure 4.6 Three types of isomers
Figure 4.6ax Structural isomers
Table 4.1 Functional Groups of Organic Compounds
Figure 4.8 A comparison of functional groups of female (estradiol) and male
(testosterone) sex hormones
Figure 4.8x1 Estrone and testosterone
THE STRUCTURE AND
FUNCTION OF
MACROMOLECULES
CHAPTER 5
Macromolecules
– long molecule
consisting of many similar
building blocks connected by
covalent bonds
 Monomer – the building blocks
 Polymer
SYNTHESIS AND BREAKDOWN
 Dehydration
synthesis
(condensation reaction) –
removal of water to join 2
compounds
 Hydrolysis – addition of water to
break a bond between 2
compounds
Figure 5.2 The synthesis and breakdown of polymers
CARBOHYDRATES



Monosaccharides
 Examples: glucose, fructose, and galactose
 One sugar
Disaccharides
 Examples: Lactose, sucrose and maltose
 Two sugars
 Joined by glycosidic linkage via dehydration
synthesis
Polysaccharides
 Examples: starch, glycogen, and cellulose
 Many sugars
Figure 5.3 The structure and classification of some monosaccharides
Figure 5.4 Linear and ring forms of glucose
Figure 5.5 Examples of disaccharide synthesis
Figure 5.6 Storage polysaccharides
Figure 5.7a Starch and cellulose structures
Figure 5.7b,c Starch and cellulose structures
Figure 5.8 The arrangement of cellulose in plant cell walls
Figure 5.x1 Cellulose digestion: termite and Trichonympha
Figure 5.x2 Cellulose digestion: cow
Figure 5.9 Chitin, a structural polysaccharide: exoskeleton and surgical thread
LIPIDS





Little or no affinity for water (hydrophobic)
Examples: Fat, phospholipids, and steroids
Fats – composed of glycerol (an alcohol) and
fatty acids
 Saturated – no double bonds in carbon
chain
 Unsaturated – at least one double bond in
carbon chain
Phospholipids – make up plasma
membrane
Steroids – consist of 4 carbon rings;
examples include cholesterol, vitamin D,
estrogen, and testosterone
Figure 5.11 Examples of saturated and unsaturated fats and fatty acids
Figure 5.11x Saturated and unsaturated fats and fatty acids: butter and oil
Figure 5.12 The structure of a phospholipid
Figure 5.13 Two structures formed by self-assembly of phospholipids in
aqueous environments
Figure 5.10 The synthesis and structure of a fat, or triacylglycerol
Figure 5.14 Cholesterol, a steroid
Figure 5.14x Cholesterol
Table 5.1 An Overview of Protein Functions
PROTEIN
 Polypeptide
– polymer of amino
acids
 Protein – consists of one or more
polypeptides with specific 3-D
structure
 There are 20 different amino
acids differing only by the R
group
Figure 5.15 The 20 amino acids of proteins: nonpolar
Figure 5.15 The 20 amino acids of proteins: polar and electrically charged
Figure 5.16 Making a polypeptide chain

FOUR LEVELS OF PROTEIN
STRUCTURE




Primary – sequences of amino acids
Secondary – hydrogen bonds cause coils and
folds
 Beta pleated sheet and alpha helix
Tertiary – irregular contortions due to various
weak bonds:
 Hydrophobic interactions
 Disulfide bridges
 Ionic bonds
 Van der Waals interactions
Quaternary – two or more polypeptide chains
aggregated into one functional macromolecule
 Examples: collagen and hemoglobin
Figure 5.18 The primary structure of a protein
Figure 5.19 A single amino acid substitution in a protein causes sickle-cell
disease
Figure 5.19x Sickled cells
Figure 5.20 The secondary structure of a protein
Figure 5.22 Examples of interactions contributing to the tertiary structure of a
protein
Figure 5.23 The quaternary structure of proteins
Figure 5.24 Review: the four levels of protein structure
Figure 5.25 Denaturation and renaturation of a protein
Figure 5.21 Spider silk: a structural protein
Figure 5.21x Silk drawn from the spinnerets at the rear of a spider
NUCLEIC ACIDS
 Examples:
DNA and RNA
 Nucleotides – building blocks of
nucleic acids (made of sugar,
phosphate, and a N-base)
 We will discuss these in great
detail later in the semester! 
Figure 5.29 The components of nucleic acids
Figure 5.30 The DNA double helix and its replication
AN INTRODUCTION TO
METABOLISM
CHAPTER 8
Figure 6.2 Transformations between kinetic and potential energy
Figure 6.2x1 Kinetic and potential energy: dam
Figure 6.2x2 Kinetic and potential energy: cheetah at rest and running
BIOENERGETICS
– the study of how
energy flows through living organisms
 Catabolic – reactions or pathways
where a larger molecule is broken
down into smaller molecules
 Anabolic - reactions or pathways
where smaller molecules are joined to
build a larger molecule
 Bioenergetics
THERMODYNAMICS
Law of Thermodynamics –
energy can be transferred and
transformed, but not created nor
destroyed
 Second Law of Thermodynamics –
every energy transfer or
transformation makes the universe
more disordered (have more
entropy)
 First
Entropy – measure of disorder or
randomness
 Most energy transformations involve at
least some energy be changed to heat
 Heat is the lowest form of energy
 Biological order has increased over time
 Second law requires only that
processes increase the entropy of the
universe
 Organisms may decrease entropy but
entire universe must increase entropy

Figure 6.4 Order as a characteristic of life
energy (G) – energy available to
do work when temperature is uniform
throughout system
 Free
G = H – TS
= system’s total energy (enthalpy)
 T = temp in Kelvin (° C + 273)
 S = entropy
H
∆ G = G final state – G starting state
∆ G = ∆ H - T∆ S
 For a spontaneous reaction:
 ∆ G = negative
So must: give up energy (decrease
H) and/or give up order (increase S)
∆ G = 0 at equilibrium
 Free energy increases if move away from
equilibrium and decreases if move toward
equilibrium
Figure 6.5 The relationship of free energy to stability, work capacity, and
spontaneous change
 Exergonic
∆
G = negative
 Spontaneous
 Net release of energy
 Endergonic
 ∆ G = positive
 NOT spontaneous
 Stores free energy in molecules
Figure 6.6 Energy changes in exergonic and endergonic reactions
 Cells
at equilibrium are dead!
 Cells can keep disequilibrium by
having products of one reaction not
accumulate but instead become
reactants of another reaction
 Energy coupling – an exergoinc
reaction drives an endergoinc reaction
Figure 6.7 Disequilibrium and work in closed and open systems
ATP = ADENOSINE TRIPHOSPHATE
ATP + H2O
ADP + Pi
∆ G = -7.3kcal/mol
 ADP = adenosine diphosphate
 Normally the phosphate is bonded to
an intermediate compound which is
then considered phosphorylated
 The reverse reaction is endergonic and
requires +7.3 kcal/mol to make ATP
from ADP
Figure 6.8 The structure and hydrolysis of ATP
Figure 6.9 Energy coupling by phosphate transfer
Figure 6.10 The ATP cycle
ENZYMES



Catalytic proteins or enzymes – change
the rate of reaction without being
consumed by the reaction
Activation energy (EA) – energy needed
to start a reaction
 Energy needed to contort the reactants
so the bonds can change
Enzymes lower activation energy by
enabling reactants to absorb enough
energy to reach transition state at
moderate temps
 Enzymes
are substrate specific
 Substrate – the reactant on
which an enzyme works
 Active site – area on enzyme
where substrate fits
 Induced fit – model of enzyme
activity
Figure 6.12 Energy profile of an exergonic reaction
The effect of an enzyme on activation
Figure 6.14 The induced fit between an enzyme and its substrate
Figure 6.15 The catalytic cycle of an enzyme
 Effects
of pH and Temp
 Optimal temperatures and pH
ranges exist for enzymes
Figure 6.16 Environmental factors affecting enzyme activity


Cofactors – non-protein helpers that bind to
active site or substrate (zinc, iron)
 Coenzymes – cofactors that are organic
(vitamins)
Enzyme Inhibitors – reduce enzyme activity
 Competitive inhibitors – block substrate
from entering active site
 Reversible
 Overcome by adding more substrate
 Noncompetitive inhibitors – bind to another
part of enzyme thereby changing the
enzyme’s shape making it inactive
 Irreversible
 Examples: DDT, sarin gas, and penicillin
Figure 6.17 Inhibition of enzyme activity



Allosteric regulation
 Allosteric sites – receptors on
enzymes (not the active site) that may
either inhibit or stimulate enzyme
activity
Feedback inhibition
 End product of a pathway acts as a
inhibitor of an enzyme within the
pathway
Cooperativity – an enzyme with multiple
subunits where binding to one active site
causes shape changes to rest of subunits
which in turn activates those subunits
Figure 6.18 Allosteric regulation of enzyme activity
Figure 6.19 Feedback inhibition
Figure 6.20 Cooperativity