Transcript Notes Biochemistry AP

BIOCHEMISTRY
WATER
Chapter 3: Water and Life
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)
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/Adhesion
2. Moderation of temperature
3. Insulation of bodies of water by
floating ice
4. The solvent of life
Trees
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 how 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
slower 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 – the 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 of 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! )
Ice Caps/Glaciers Melting
• Global warming causing ice to melt.
– Loss of habitat
– Less sunlight reflected into space so
increases temp
The structure of ice
Ice, water, and steam
Floating ice and the fitness of the environment
Floating ice and the fitness of the environment: ice fishing
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)
A crystal of table salt dissolving in water
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
Moles
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
• pH = -log [H+]
– For neutral water:
• pH = -log 10-7 = -(-7) = 7
– pH decreases as H+ increases
– A pH of 3 vs. pH of 6 is a 1000 fold
difference (10 fold for each step)
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- +
bicarbonate
H+
proton
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.
The effects of acid precipitation on a forest
Pulp mill
Acid rain damage to statuary, 1908 & 1968
CARBON AND THE
MOLECULAR DIVERSITY
OF LIFE
CHAPTER 4
ISOMERS
• Compounds with the same chemical formula
but different structures
Functional Groups
 See
diagram of functional groups
Three types of isomers
Structural isomers
Functional Groups of Organic Compounds
A comparison of functional groups of female (estradiol) and male (testosterone)
sex hormones
THE STRUCTURE AND
FUNCTION OF LARGE
BIOLOGICAL MOLECULES
CHAPTER 5
Macromolecules
• Polymer – long molecule consisting of
many similar building blocks
connected by covalent bonds
• Monomer – the building blocks
SYNTHESIS AND
BREAKDOWN
• Dehydration synthesis (condensation
reaction) – removal of water to join 2
compounds (anabolic)
• Hydrolysis – addition of water to break
a bond between 2 compounds (catabolic)
The synthesis and breakdown of polymers
CARBOHYDRATES
• Monomers are sugars.
• Used for energy.
• Types of carbohydrates
– Monosaccharides – one sugar
• Examples: glucose, fructose, and galactose
– Disaccharides – two sugars joined
• Examples: lactose, sucrose and maltose
• Joined by glycosidic linkage via dehydration
synthesis
– Polysaccharides – many sugars joined
• Examples: starch, glycogen, chitin
(exoskeletons of insects), and cellulose (fiber)
The structure and classification of some monosaccharides
Linear and ring forms of glucose
Examples of disaccharide synthesis
Storage polysaccharides
Starch and cellulose structures
Starch and cellulose structures
The arrangement of cellulose in plant cell walls
Cellulose digestion: termite and Trichonympha
Cellulose digestion: cow
Chitin, a structural polysaccharide: exoskeleton and surgical thread
LIPIDS
• Little or no affinity for water (hydrophobic)
• No monomers
• Many uses including insulation, store energy,
hormones
• 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
Examples of saturated and unsaturated fats and fatty acids
Saturated and unsaturated fats and fatty acids: butter and oil
The structure of a phospholipid
Two structures formed by self-assembly of phospholipids in aqueous
environments
The synthesis and structure of a fat, or triacylglycerol
Cholesterol, a steroid
An Overview of Protein Functions
PROTEIN
• Monomers are amino acids.
• Many uses including enzymes, antibodies,
receptors, structural, hormones, and
transport
• Protein – consists of one or more
polypeptides with specific 3-D structure
– Polypeptide = polymer of amino acids
– Peptide = bond between 2 amino acids
• There are 20 different amino acids
differing only by the R group
The 20 amino acids of proteins: nonpolar
The 20 amino acids of proteins: polar and electrically charged
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 (weak attractions
from transient partial charges)
– Quaternary – two or more polypeptide chains
aggregated into one functional macromolecule
• Examples: collagen and hemoglobin
The primary structure of a protein
A single amino acid substitution in a protein causes sickle-cell disease
Sickled cells
The secondary structure of a protein
Examples of interactions contributing to the tertiary structure of a protein
The quaternary structure of proteins
Review: the four levels of protein structure
Protein Structure
• The function of a protein depends on
its 4 levels.
• Denaturation – the “unraveling of a
protein” so that the 2nd, 3rd, and 4th
level structures are gone and all that
is left is the primary sequence
– Loss of structure = loss of function
– Causes of denaturation - High temp,
high or low pH
Denaturation and renaturation of a protein
Spider silk: a structural protein
Silk drawn from the spinnerets at the rear of a spider
NUCLEIC ACIDS
• Monomers are nucleotides.
• Nucleotides – are made of sugar,
phosphate, and a N-base
• Examples: DNA, ATP, and RNA
• We will discuss these in great detail
later in the semester! 
The components of nucleic acids
The DNA double helix and its replication
AN INTRODUCTION TO
METABOLISM
CHAPTER 8
Transformations between kinetic and potential energy
Kinetic and potential energy: dam
Kinetic and potential energy: cheetah at rest and running
BIOENERGETICS
• 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
THERMODYNAMICS
• First 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)
– 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
Order as a characteristic of life
• Free energy (G) – energy available to do
work when temperature is uniform
throughout system
G = H – TS
• H = system’s total energy (enthalpy)
• T = temp in Kelvin (° C + 273)
• S = entropy (measure of disorder)
∆ 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
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
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 (or a
reaction that increase entropy is paired
with one that decreases entropy)
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
The structure and hydrolysis of ATP
Energy coupling by phosphate transfer
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
Energy profile of an exergonic reaction
The effect of an enzyme on activation
FThe induced fit between an enzyme and its substrate
The catalytic cycle of an enzyme
• Effects of pH and Temp
– Optimal temperatures and pH ranges exist
for enzymes
– High temperatures can denature an enzyme
– Changes in pH can also denature enzymes
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
Inhibition of enzyme activity
• Allosteric regulation
– Allosteric sites – receptors on
enzymes (not the active site) that
may either inhibit or stimulate
enzyme activity
• 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
• Feedback inhibition
– End product of a pathway acts as a
inhibitor of an enzyme within the
pathway
Allosteric regulation of enzyme activity
Cooperativity
Feedback inhibition