Download The Chemistry of Life

The Chemistry of Life
The Nature of Matter
• Life depends on chemistry
• When you eat food or inhale oxygen, your body
uses these materials in chemical reactions
that keep you alive
– Just as buildings are made from bricks, steel, glass,
and wood, living things are made from chemical
compounds
– If the first task of an architect is to understand building
materials, then the first job of a biologist is to
understand the chemistry of life
• YOU ARE WHAT YOU EAT!!!!!!!!!
Atoms
• The study of chemistry begins with the basic unit of matter, the
atom
– The Greek word atomos, which means “unable to be cut,” was
first used to refer to matter by the Greek philosopher Democritus
nearly 2500 years ago
– Democritus asked a simple question: If you take an object like a
stick of chalk and break it in half, are both halves still chalk?
• The answer, of course, is yes
• But what happens if you go on? Suppose you break it in half again
and again and again
– Can you continue to divide without limit, or does there come a
point at which you cannot divide the fragment of chalk without
changing it into something else?
– Democritus thought that there had to be a limit
– He called the smallest fragment the atom, a name scientists
still use today
Atoms
• Atoms are incredibly small
• Placed side by side, 100 million atoms
would make a row only about 1 centimeter
long—about the width of your little finger!
• Despite its extremely small size, an
atom contains subatomic particles that
are even smaller
Atoms
• The figure TO THE RIGHT
shows the subatomic particles
in a helium atom
• The subatomic particles that
make up atoms are protons,
neutrons, and electrons
– Protons and neutrons have
about the same mass
– However, protons are
positively charged particles (+)
and neutrons carry no charge
• Their name is a reminder that
they are neutral particles
• Strong forces bind protons and
neutrons together to form the
nucleus, which is at the center
of the atom
Atoms
Atoms
• Helium atoms contain
protons, neutrons, and
electrons:
– The positively charged
protons and uncharged
neutrons are bound
together in the dense
nucleus, while the
negatively charged
electrons move in the
space around the
nucleus
Atoms
• The electron is a negatively charged particle
(−) with 1/1840 the mass of a proton
• Electrons are in constant motion in the space
surrounding the nucleus
• They are attracted to the positively charged
nucleus but remain outside the nucleus
because of the energy of their motion
• Because atoms have equal numbers of
electrons and protons, and because these
subatomic particles have equal but opposite
charges, atoms are neutral before they
react!!!!!!
Elements
• A chemical element is a pure substance that
consists entirely of one type of atom
• More than 100 elements are known, but only about two
dozen are commonly found in living organisms
• Elements are represented by a one- or two-letter symbol
– Example:
• C stands for carbon
• H for hydrogen
• Na for sodium
• The number of protons in an atom of an element is
the element's atomic number
– Carbon's atomic number is 6, meaning that each atom of carbon
has six protons and, consequently, six electrons
• ATOM
–IS THE FUNDAMENTAL UNIT OF
MATTER
–COMPOSED OF SUBATOMIC
PARTICLES
• ATOMIC NUCLEUS
–PROTON
–NEUTRON
–ELECTRONS:ORBIT (OUTSIDE) THE
NUCLEUS
• ATOMIC NUMBER — THE NUMBER
OF PROTONS IN THE NUCLEUS OF
THE ATOM.
–BEFORE AN ATOM REACTS THE
NUMBER OF PROTONS AND
ELECTRONS ARE EQUAL
• MASS NUMBER (ATOMIC MASS) — IS THE SUM
OF THE PROTONS AND NEUTRONS IN THE
NUCLEUS OF THE ATOM.
– amu = .000,000,000,000,000,000,000,001,67 g
– Electron mass is so small (negligible) that when
calculating MASS NUMBER of an atom the mass of
the electron is considered zero ( 0).
•
LOCATION MASS CHARGE
•
( in amu)
• _____________________________________
PROTON
nucleus
1
+1
• _____________________________________
• NEUTRON nucleus
1
0
• _____________________________________
• ELECTRON outside
1/2000
-1
•
nucleus
• _____________________________________
Isotopes
• Atoms of an element can have different numbers of neutrons
– Example:
• Some atoms of carbon have six neutrons, some have seven, and a few have
eight
• Atoms of the same element that differ in the number of
neutrons they contain are known as isotopes
• The sum of the protons and neutrons in the nucleus of an atom is
called its mass number
– Isotopes are identified by their mass numbers.
– The figure at right shows the subatomic composition of carbon-12,
carbon-13, and carbon-14 atoms
• The weighted average of the masses of an element's isotopes
is called its atomic mass
– “Weighted” means that the abundance of each isotope in nature is
considered when the average is calculated
• Because they have the same number of electrons, all isotopes
of an element have the same chemical properties
Isotopes of Carbon
• Because they have the same number of
electrons, these isotopes of carbon have the
same chemical properties
– The difference among the isotopes is the number of
neutrons in their nuclei
Isotopes of Carbon
Radioactive Isotopes
• Some isotopes are radioactive, meaning that
their nuclei are unstable and break down at a
constant rate over time:
– The radiation these isotopes give off can be
dangerous, but radioactive isotopes have a number of
important scientific and practical uses:
• Geologists can determine the ages of rocks and fossils by
analyzing the isotopes found in them
• Radiation from certain isotopes can be used to treat cancer
and to kill bacteria that cause food to spoil
• Radioactive isotopes can also be used as labels or “tracers”
to follow the movements of substances within organisms
Chemical Compounds
• In nature, most elements are found combined with
other elements in compounds
– A chemical compound is a substance formed by the
chemical combination of two or more elements in definite
proportions
• Scientists show the composition of compounds by a kind
of shorthand known as a chemical formula:
– Water, which contains two atoms of hydrogen for each atom of
oxygen, has the chemical formula H2O
– The formula for table salt, NaCl, indicates that the elements
from which table salt forms—sodium and chlorine—combine in a
1 : 1 ratio
Chemical Compounds
•
•
The physical and chemical properties of a compound are usually very
different from those of the elements from which it is formed:
– Example:
• Hydrogen and oxygen, which are gases at room temperature, can
combine explosively and form liquid water
• Sodium is a silver-colored metal that is soft enough to cut with a
knife
– It reacts explosively with cold water
• Chlorine is very reactive, too
– It is a poisonous, greenish gas that was used to kill many
soldiers in World War I.
Sodium and chlorine combine to form sodium chloride (NaCl), or table
salt
– Sodium chloride is a white solid that dissolves easily in water. As you
know, sodium chloride is not poisonous
– In fact, it is essential for the survival of most living things
Chemical Bonds
• The atoms in compounds are held together
by chemical bonds
• Much of chemistry is devoted to understanding
how and when chemical bonds form
• Bond formation involves the electrons that
surround each atomic nucleus:
– The electrons that are available to form bonds are
called valence electrons
• The main types of chemical bonds are ionic
bonds and covalent bonds
Ionic Bonds
• An ionic bond is formed when one or more electrons
are transferred from one atom to another
• Recall that atoms are electrically neutral because they
have equal numbers of protons and electrons
– Electrically stable BUT chemically unstable
• An atom that loses electrons has a positive charge
• An atom that gains electrons has a negative charge
• These positively and negatively charged atoms are
known as ions
• ATOMIC NUMBER — THE NUMBER
OF PROTONS IN THE NUCLEUS OF
THE ATOM
–BEFORE AN ATOM REACTS THE
NUMBER OF PROTONS AND
ELECTRONS ARE EQUAL
• ELECTRONS
–Do not move about an atom in definite
orbits
–Only the probability of finding an
electron at a particular place in an atom
can be determined
–Each electron seems to be locked into a
certain area in the electron cloud
• Modern Atomic Theory
–Electrons are arranged in energy
levels
–An energy level represents the most
likely location in the electron cloud
in which an electron can be found
• MODERN ATOMIC THEORY
–Electrons with the lowest energy are
found in the energy level closest to the
nucleus
–Electrons with the highest energy are
found in the energy levels farther from
the nucleus
–Each energy level can hold only a
maximum number of electrons
• MODERN ATOMIC THEORY
–First energy level—maximum of 2
electrons
–Second energy level— maximum
of 8 electrons
–Third energy level—maximum of
18 electrons (three sublevels of
2/8/8)
• Chemical activity—depends on the
arrangement of electrons in the
outermost energy level
• FORCES WITHIN THE ATOM
– Electromagnetic—negative charged electrons are
attracted to the positive charged protons
– Strong Force—prevents the positively charged
protons from repelling each other
• keeps protons together
– Weak Force—responsible for radioactive decay
• neutron changes into a proton and an electron
– Gravity: force of attraction that depends on the
mass of two objects and the distance between them
• ATOMS AND BONDING
–Before an atom reacts it is
electrically neutral (same number
of protons and electrons) —
electrically stable
• ATOMS AND BONDING:
– However the atom might not be chemically
stable:
– Chemical stability depends on the valence
electrons (outermost energy level)
• 1/2/3 electrons — will lose electrons
• 5/6/7 electrons — will gain electrons
• after losing/gaining electrons the atom will
be chemically stable but now will become
electrically unstable
• ATOMS AND BONDING:
–lose 1 electron +1
–lose 2 electrons +2
–lose 3 electrons +3
–gain 3 electrons -3
–gain 2 electrons -2
–gain 1 electron -1
• When the outermost energy level
(valence electrons) is filled, the
atom is chemically stable
• IONIC BONDS—involves the transfer
of electrons
–One atom gains electrons and the
other atom loses electrons resulting
in filled outer energy levels
–Ions are formed (charged atom or
group of atoms — polyatomic)
Ionic Bonds
• The figure above shows how ionic bonds form between
sodium and chlorine in table salt
• A sodium atom easily loses its one valence electron
and becomes a sodium ion (Na+)
• A chlorine atom easily gains an electron and
becomes a chloride ion (Cl−)
Ionic Bonds
• IONIC BONDS:
–The process of removing electrons and
forming positive ions (more protons than
electrons) is called ionization (energy is
absorbed)
–Energy is needed for ionization—
ionization energy:
• Low for atoms with few valence
electrons (metals)
• High for atoms with many valence
electrons (nonmetals)
• IONIC BONDS:
–The process of gaining electrons and
forming negative ions (more electrons
than protons) is called electron affinity
(energy is released)
• Low for atoms with few valence
electrons (metals)
• High for atoms with many valence
electrons (nonmetals)
• IONIC BONDS:
–It is much easier to gain 1 or 2
electrons than to lose 6 or 7
electrons!!!!
–Positive (+) ions attract negative
(-) ions resulting in ionic bonds
Ionic Bonds
• The chemical bond in which electrons are
transferred from one atom to another is
called an ionic bond
• The compound sodium chloride (NaCl)
forms when sodium loses its valence
electron to chlorine
Ionic Bonds
• In a salt crystal, there are trillions of
sodium and chloride ions
• These oppositely charged ions have a
strong attraction
• The attraction between oppositely charged
ions is an ionic bond.
• IONIC BONDS:
–The placement of ions in an ionic
compound results in a regular, repeating
arrangement called a crystal lattice
–Gives great stability/high melting points
–Chemical formula shows the ratio of
ions not the actual number present
–Each ionic compound has a
characteristic crystal lattice
arrangement
• COVALENT BONDS — sharing of
electrons
–Results in filled outer energy levels of
both sharing atoms
–The positively charged nucleus of
each atom simultaneously attracts the
negatively charged electrons that are
being shared
Covalent Bonds
• Sometimes electrons are shared by atoms instead of
being transferred
• What does it mean to “share” electrons?
– It means that the moving electrons actually travel in the
orbitals of both atoms
• A covalent bond forms when electrons are shared
between atoms
– When the atoms share two electrons, the bond is called a single
covalent bond
– Sometimes the atoms share four electrons and form a double
bond
– In a few cases, atoms can share six electrons and form a triple
bond
Covalent Bonds
• The structure that
results when atoms are
joined together by
covalent bonds is
called a molecule
• The molecule is the
smallest unit of most
compounds
• The diagram, to the right,
of a water molecule
shows that each
hydrogen atom forms a
single covalent bond with
the oxygen atom
Covalent Bonds
• COVALENT BONDS:
–Combination of atoms formed by a
covalent bond are called molecules
• Molecule is the smallest particle of
a covalently bonded substance
that has all the properties of that
substance
• Chemical formula for a molecule
shows the exact number of atoms
of each element involved in the
bond
–Tend to have low melting points
• COVALENT BONDS:
–Represented by electron dot
diagrams
–Chemical symbol represents the
nucleus and all inner energy
levels
–Dots surrounding the symbol
represent the valence
(outermost) electrons
Covalent Bonds
• Sharing is NOT always equal
• O2 and H2 sharing is equal
• H2O sharing is NOT equal
– Results in the molecule having a slight
electrical charge
– POLAR
Van der Waals Forces
• Because of their structures, atoms of different
elements do not all have the same ability to
attract electrons
• Some atoms have a stronger attraction for
electrons than do other atoms
• Therefore, when the atoms in a covalent
bond share electrons, the sharing is not
always equal
• Even when the sharing is equal, the rapid
movement of electrons can create regions on a
molecule that have a tiny positive or negative
charge
Van der Waals Forces
• When molecules are close together, a slight
attraction can develop between the
oppositely charged regions of nearby
molecules
– Chemists call such intermolecular forces of attraction
van der Waals forces, after the scientist who
discovered them
• Although van der Waals forces are not as
strong as ionic bonds or covalent bonds,
they can hold molecules together, especially
when the molecules are large
Van der Waals Forces
•
•
•
People who keep geckos as pets have already seen van der Waals
forces in action
These remarkable little lizards can climb up vertical surfaces, even smooth
glass walls, and then hang on by a single toe despite the pull of gravity
How do they do it?
– No, they do not have some sort of glue on their feet and they don't have
suction cups
•
A gecko foot is covered by as many as half a million tiny hairlike projections:
– Each projection is further divided into hundreds of tiny, flat-surfaced fibers.
– This design allows the gecko's foot to come in contact with an extremely large
area of the wall at the molecular level
•
Van der Waals forces form between molecules on the surface of the
gecko's foot and molecules on the surface of the wall permits the
animal to climb vertical structures:
– The combined strength of all the van der Waals forces allows the gecko to
balance the pull of gravity
– When the gecko needs to move its foot, it peels the foot off at an angle and
reattaches it at another location on the wall
Properties of Water
• Water is also the single most abundant compound in
most living things
• Water is one of the few compounds that is a liquid at the
temperatures found over much of Earth's surface
• Unlike most substances, water expands as it freezes
• Thus, ice is less dense than liquid water, which explains
why ice floats on the surface of lakes and rivers
• If the ice sank to the bottom, the situation would be
disastrous for fish and plant life in regions with cold
winters, to say nothing of the sport of ice skating!
Water Molecule
• Like all molecules, a water molecule (H2O)
is neutral (BUT conducts electricity!!!!!)
– The positive charges on its 10 protons
balance out the negative charges on its 10
electrons
• However, there is more to the story
• WATER:
–70% of earth’s surface is covered by
water
–65% of your body mass is water
–Thousands of substances dissolve
(soluble) in water (UNIVERSAL
SOLVENT)
–Certain substances will not dissolve
in water (insoluble)
Polarity
•
•
With 8 protons in its nucleus, an
oxygen atom has a much stronger
attraction for electrons than does
the hydrogen atom with a single
proton in its nucleus
– Thus, at any moment, there is a
greater probability of finding
the shared electrons near the
oxygen atom than near the
hydrogen atom
Because the water molecule has a
bent shape, as shown to the right, the
oxygen atom is on one end of the
molecule and the hydrogen atoms
are on the other
– As a result, the oxygen end of
the molecule has a slight
negative charge and the
hydrogen end of the molecule
has a slight positive charge
Water Molecule
• WATER STRUCTURE:
–Two hydrogen atoms bond
covalently with one oxygen atom
(electrons are shared)
• Sharing is unequal
–Oxygen—slight negative charge
–Hydrogen—slight positive charge
Water Molecule
• The unequal sharing of electrons causes a
water molecule to be polar:
– The hydrogen end of the molecule is slightly
positive, and the oxygen end is slightly negative
• A molecule in which the charges are unevenly
distributed is called a polar molecule because
the molecule is like a magnet with poles
• A water molecule is polar because there is an
uneven distribution of electrons between the
oxygen and hydrogen atoms
– The negative pole is near the oxygen atom and the
positive pole is between the hydrogen atoms
Hydrogen Bonds
•
•
•
•
Because of their partial positive
and negative charges, polar
molecules such as water can
attract each other, as shown to
the right
The charges on a polar molecule
are written in parentheses, (−) or
(+), to show that they are weaker
than the charges on ions such as
Na+ and Cl−
The attraction between the
hydrogen atom on one water
molecule and the oxygen atom
on another water molecule is an
example of a hydrogen bond
Hydrogen bonds are not as
strong as covalent or ionic
bonds, but water's ability to
form multiple hydrogen bonds
is responsible for many of its
special properties
Hydrogen Bonds
Cohension
• A single water molecule may be involved in as many as
four hydrogen bonds at the same time:
– The ability of water to form multiple hydrogen bonds is
responsible for many of water's properties
• Cohesion is an attraction between molecules of the
same substance:
– Because of hydrogen bonding, water is extremely cohesive
– Water's cohesion causes molecules on the surface of water to be
drawn inward, which is why drops of water form beads on a
smooth surface
• Cohesion also explains why some insects and spiders
can walk on a pond's surface
ADHESION
• Adhesion is an attraction between molecules of
different substances
• Have you ever been told to read the volume in a
graduated cylinder at eye level?
– The surface of the water in the graduated cylinder dips
slightly in the center because the adhesion between water
molecules and glass molecules is stronger than the
cohesion between water molecules
– Adhesion between water and glass also causes water to rise
in a narrow tube against the force of gravity:
• This effect is called capillary action
• Capillary action is one of the forces that draw water out of the roots
of a plant and up into its stems and leaves
• Cohesion holds the column of water together as it rises
• METALLIC BONDS—outer electrons of the
atoms form a common electron cloud
– The electrons become the property of all the
atoms
– The positive nuclei of atoms of metals are
surrounded by free-moving electrons
that
are all attracted by the nuclei at the same
time
– Electrons are free to flow
– Excellent conductors of both heat and
electricity
– High melting points
• OXIDATION NUMBER:
– Describes the combining capacity of an atom
– Indicates the number of electrons an atom gains,
loses, or shares when it forms chemical bonds
– Positive number indicates a lose of electrons
– Negative number indicates a gain of electrons
– Number indicates how many electrons
– Used to predict how atoms will combine and what
the formula for the resulting compound will be
– The sum of the oxidation numbers of the atoms
in a compound must be ZERO
Solutions and Suspensions
• Water is not always pure—it is often found as
part of a mixture
– A mixture is a material composed of two or more
elements or compounds that are physically mixed
together but not chemically combined
• Salt and pepper stirred together constitute a mixture
• So do sugar and sand
• Earth's atmosphere is a mixture of gases
• Living things are in part composed of
mixtures involving water:
– Two types of mixtures that can be made with
water are solutions and suspensions
Solutions
• If a crystal of table salt is
placed in a glass of warm
water, sodium and chloride
ions on the surface of the
crystal are attracted to the
polar water molecules
• Ions break away from the
crystal and are surrounded
by water molecules, as
illustrated to the right
• The ions gradually become
dispersed in the water,
forming a type of mixture
called a solution
Solutions
• All the components of a
solution are evenly
distributed throughout the
solution
• In a salt–water solution, table
salt is the solute—the
substance that is dissolved
• Water is the solvent—the
substance in which the solute
dissolves
• Water's polarity gives it the
ability to dissolve both ionic
compounds and other polar
molecules, such as sugar
• Without exaggeration, water is
the greatest solvent on Earth.
Solutions
• NaCl Dissolving in
Water
• When an ionic
compound such as
sodium chloride is
placed in water, water
molecules surround
and separate the
positive and
negative ions
Solutions
• WATER STRUCTURE:
– Molecule has oppositely charged ends
– Charged ends give the property of
Polarity:
• A force of attraction is set up between
the solute and solvent
–Separates the molecules of the solute,
causing the solute to dissolve
– NONPOLAR substances will not dissolve
in water but will dissolve in nonpolar
solvents
– LIKE DISSOLVES LIKE
Suspensions
• Some materials do not dissolve when placed in
water but separate into pieces so small that they do
not settle out
• The movement of water molecules keeps the small
particles suspended
• Such mixtures of water and nondissolved material are
known as suspensions
• Some of the most important biological fluids are both
solutions and suspensions
• The blood that circulates through your body is mostly
water, which contains many dissolved compounds
– However, blood also contains cells and other undissolved
particles that remain in suspension as the blood moves through
the body
Acids, Bases, and pH
•
•
•
A water molecule can react to form ions
This reaction can be summarized by a chemical equation in which double
arrows are used to show that the reaction can occur in either direction
How often does this happen?
– In pure water, about 1 water molecule in 550 million reacts and
forms ions
– Because the number of positive hydrogen ions produced is equal to the
number of negative hydroxide ions produced, water is neutral
The pH scale
•
•
•
•
•
•
•
•
•
Chemists devised a measurement system
called the pH scale to indicate the
concentration of H+ ions in solution
As the figure at right shows, the pH scale
ranges from 0 to 14
At a pH of 7, the concentration of H+ ions
and OH− ions is equal
Pure water has a pH of 7
Solutions with a pH below 7 are called
acidic because they have more H+ ions
than OH− ions
The lower the pH, the greater the acidity
Solutions with a pH above 7 are called
basic because they have more OH− ions
than H+ ions.
The higher the pH, the more basic the
solution.
Each step on the pH scale represents a
factor of 10:
–
Example:
•
liter of a solution with a pH of 4 has 10 times as
many H+ ions as a liter of a solution with a pH of
5
• pH SCALE
–Measure of the hydronium ion ( H3O+)
concentration (hydrogen ion H+ ion
concentration)
–Hydronium ion is formed by the
attraction between a hydrogen ion
(H+ ) from an acid and a water
molecule( H2O )
• pH
–Indicates HOW ACIDIC a solution
is
–Series of numbers 0 to 14
• middle is 7 (neutral point)
• below 7 (acid)
• above 7 (base)
The pH scale
• pH
– 10-7 MOLES OF H3O+ IONS IN 1 LITER OF H2O
– pH 7
– 0.0000001 moles H3O+ / liter H2O
• pH: MOLES H3O+ / LITER H2O
pH 1 ( 0.1 moles/liter ) ( 10 -1 moles/l)
pH 2 ( 0.01 moles/liter ) ( 10 -2 moles/l)
pH 3 ( 0.001 moles/liter ) ( 10 -3 moles/l)
pH 4 ( 0.0001 moles/liter ) ( 10 -4 moles/l)
pH 5 ( 0.00001 moles/liter ) ( 10 -5 moles/l)
pH 6 ( 0.000001 moles/liter ) ( 10 -6 moles/l)
pH 7 ( 0.0000001 moles/liter) ( 10 -7 moles/l)
pH 8 ( 0.00000001 moles/liter ) ( 10 -8 moles/l)
pH 9 ( 0.000000001 moles/liter ) ( 10 -9 moles/l)
pH 10 ( 0.0000000001 moles/liter ) ( 10 -10 moles/l)
pH 11 ( 0.00000000001 moles/liter ) ( 10 -11 moles/l)
pH 12 ( 0.000000000001 moles/liter ) ( 10 -12 moles/l)
pH 13 ( 0.0000000000001 moles/liter )( 10 -13 moles/l)
pH 14 ( 0.00000000000001 moles/liter )( 10 -14moles/l)
Acids
• Where do all those extra H+
ions in a low-pH solution come
from?
– They come from acids
• An acid is any compound that
forms H+ ions in solution:
– Acidic solutions contain
higher concentrations of H+
ions than pure water and
have pH values below 7
– Strong acids tend to have pH
values that range from 1 to 3
– The hydrochloric acid
produced by the stomach to
help digest food is a strong
acid.
• ACIDS:
–Physical property—sour taste
(never taste in lab)
–Affects color of indicators:
• compounds that show a
definite color change when
mixed with an acid or a base)
–litmus paper—red
–phenolphthalein—clear
(colorless)
• ACIDS:
–React with active metals to
produce hydrogen gas and a
metal compound (corrodes the
metal and produces a residue)
–Lab
–Car battery (danger)
• ACIDS:
– ALL contain HYDROGEN
• When dissolved in water, acids ionize to
produce positive (+) hydrogen ions (H+)
• Hydrogen ion is a PROTON
• Acids are defined as proton producers
• The attraction between a water (H2O)
molecule and a hydrogen ion (H+ )
results in the formation of a hydronium
ion (H3O+)
• ACIDS:
–PROTON DONOR
• STRONG ACIDS:
–Ionize to a high degree in
water and produce hydrogen
ions
–Strong electrolytes
• WEAK ACIDS:
–Do not ionize to a high degree
in water
–Produce few hydrogen ions
–Poor electrolytes
–Good BUFFERS
ACIDS
• H2SO4
• HCl
• HNO3
Bases
• A base is a compound
that produces hydroxide
ions (OH− ions) in
solution
• Basic, or alkaline,
solutions contain lower
concentrations of H+
ions than pure water
and have pH values
above 7
• Strong bases, such as
lye, tend to have pH
values ranging from 11 to
14.
• BASES:
–Physical property—bitter taste
(never taste in lab) and
slippery to the touch
–Can be poisonous and corrosive
• BASES:
–Affect color of indicators
( compounds that show a
definite color change when
mixed with an acid or a base)
•Litmus paper—blue
•Phenolphthalein—pink
• BASES:
–Emulsify, or dissolve fats and oils
–React with the fat or oil to form a
soap
• BASES:
–ALL contain the HYDROXIDE
ION (OH-)
• Since the hydroxide ion ( OH-)
can combine with a hydrogen
ion ( H+) and form water, a base
is often called a PROTON
ACCEPTOR
• STRONG BASES:
–Ionize to a high degree in
water and produce large
number of ions
–Good electrolytes
• WEAK BASES:
–Do not ionize to a high
degree in water
–Produce few ions
–Poor electrolytes
–Good BUFFERS
BASES
•
•
•
•
•
NaOH
LiOH
Ca(OH)2
Ba(OH)2
Al(OH)3
Buffers
• The pH of the fluids within most cells in the human body
must generally be kept between 6.5 and 7.5
• If the pH is lower or higher, it will affect the chemical
reactions that take place within the cells
– Thus, controlling pH is important for maintaining homeostasis
• One of the ways that the body controls pH is through
dissolved compounds called buffers
• Buffers are weak acids or bases that can react with
strong acids or bases to prevent sharp, sudden changes
in pH
ACID-BASE BALANCE
• Because of the abundance of hydrogen bonds in the
body’s functional proteins (enzymes, hemoglobin,
cytochromes, and others) they are strongly influenced
by hydrogen ion concentration:
– It follows then that nearly all biochemical reactions are
influenced by the pH of their fluid environment, and the
acid-base balance of body fluids is closely regulated
– Optimal pH varies from one body fluid to another:
• When arterial blood pH rises above 7.45, the body is in alkalosis
(alkalemia); when arterial pH falls below 7.35, the body is in
acidosis (acidemia)
– Between 7.0 and 7.35 is called physiological acidosis even though the
value is slightly basic
– Most hydrogen ions originate as metabolic by products,
although they can also enter the body via ingested foods
ACID-BASE BALANCE
• Dissociation of strong
and weak acids:
• (a): when added to water,
the strong acid HCl
dissociates completely
into its ions (H+ and Cl-)
• (b): dissociation of
H2CO3, a weak acid, is
very incomplete, and
some molecules of
H2CO3 remain
undissociated in
solution
COMPARISON OF DISSOCIATION OF
STRONG AND WEAK ACIDS
Chemical Buffer System
Bicarbonate Buffer Systems
•
•
A chemical buffer is a system of one or two molecules that acts to
resist changes in pH by binding H+ when the pH drops, or releasing H+
when the pH rises
The bicarbonate buffer system is the main buffer of the extracellular
fluid, and consists of carbonic acid and its salt, sodium bicarbonate:
– When a strong acid is added to the solution, carbonic acid is
mostly unchanged, but bicarbonate ions of the salt bind excess H+,
forming more carbonic acid:
• HCl
+ NaHCO3 → H2CO3
+ NaCl
• Strong acid + weak base → weak acid + salt
• pH lowered slightly
– When a strong base is added to solution, the sodium bicarbonate
remains relatively unaffected, but carbonic acid dissociates
further, donating more H+ to bind the excess hydroxide
• NaOH
+ H2CO3
→ NaHCO3 + H2O
• Strong base + weak acid → weak base + water
• pH rises very little
– Bicarbonates buffer system: sodium, potassium, and magnesium
Chemical Buffer System
Phosphate Buffer System
• The phosphate buffer system operates in the urine
and intracellular fluid similar to the bicarbonate
buffer system
• The components of the phosphate system are the:
– Sodium salts of dihydrogen phosphate (H2PO4-)
– Sodium salts of monohydrogen phosphate (HPO42-)
– NaH2PO4 acts as a weak acid
–
HCl
+
Na2HPO4 → NaH2PO4 + NaCl
• Strond acid + weak base
–
H+
→
weak acid
+
salt
released by strong acids is tied up in weak acids
•
NaOH
+ NaH2PO4
• Strong base
weak acid
→ Na2HPO4 + H2O
→ weak base + water
– Strong bases are converted to weak bases
Chemical Buffer System
The Protein Buffer System
•
Proteins in plasma and in cells are the body’s most plentiful and
powerful source of buffers, and constitute the protein buffer system:
– At least ¾ of all the buffering power of body fluids resides in cells, and most of
this reflects the buffering activity of intracellular proteins
•
Proteins are polymers of amino acids:
– Consists of organic acids containing carboxyl groups that dissociate to:
• Release H+ when the pH begins to rise
– R—COOH
→
R—COO-
+
H+
• Bind excess H+ when the pH declines
– R—COO-
+
H+ →
R—COOH
– Consists of an amide group that can act as a base and accept H+:
• The exposed –NH2 group can bind with hydrogen ions, becoming –NH3-
– R—NH2 + H+ → R—NH3+
» Because this removes free hydrogen ions from the solution, it prevents the solution from
becoming too acidic
» Consequently, a single protein molecule can function reversibly as either an acid
or a base depending on the pH of its environment
» Molecules with this ability are called amphoteric molecules
» Example: hemoglobin
Respiratory Regulation of H+
•
Carbon dioxide from cellular metabolism enters erythrocytes and is
converted to bicarbonate ions for transport in the plasma:
–
carbonic
–
anhydrase
– CO2 + H2O
↔
H2CO3
↔
H+ +
HCO3–
carbonic acid
bicarbonate ion
–
– When hypercapnia (increased amount of carbon dioxide in the blood)
occurs, blood pH drops, activating medullary respiratory centers,
resulting in increased rate and depth of breathing and increased
unloading of CO2 in the lungs
• The reaction is pushed to the right
– A rising plasma H+ concentration resulting from any metabolic process
excites the respiratory center indirectly (peripheral chemoreceptors) to
stimulate deeper, more rapid respiration
• As ventilation increases, more CO2 is removed from the blood,
pushing the reaction to the left and reducing the H+
concentration
Carbon Compounds
• Until the early 1800s, many chemists thought that
compounds created by organisms—organic
compounds—were distinctly different from
compounds in nonliving things
• In 1828, a German chemist was able to synthesize the
organic compound urea from a mineral called
ammonium cyanate
• Chemists soon realized that the principles governing
the chemistry of nonliving things could be applied to
living things
• Scientists still use the term organic chemistry, but now
it describes something a little different
– Today, organic chemistry is the study of all compounds that
contain bonds between carbon atoms
The Chemistry of Carbon
• Carbon atoms have four valence
electrons
– Each electron can join with an electron
from another atom to form a strong
covalent bond
• Carbon can bond with many elements,
including hydrogen, oxygen, phosphorus,
sulfur, and nitrogen
The Chemistry of Carbon
•
•
•
•
•
Carbon atoms can bond to other carbon atoms, which gives carbon the ability
to form chains that are almost unlimited in length
These carbon-carbon bonds can be single, double, or triple covalent bonds
Chains of carbon atoms can even close upon themselves to form rings, as shown
below
Carbon has the ability to form millions of different large and complex structures
No other element even comes close to matching carbon's versatility
The Chemistry of Carbon
• CARBON
– 90% of all known compounds contain
carbon
– Forms an important family of
compounds called ORGANIC
COMPOUNDS
– Forms covalent bonds (single, double,
triple) with other carbon atoms
– Straight chains, branched chains, single
rings, or joined rings
• Single bond—2 electrons
• Double bond—4 electrons
• Triple bond—6 electrons
CARBON
• A carbon atom has 6 protons
– Therefore, 6 electrons
• 2 electrons in the 1st energy level
• 4 electrons in the 2nd energy level
– Can form 4 covalent bonds
Carbon Compounds
• Bonding diagram: shows how the electrons in
the outer energy level form covalent bonds
• Molecular formula: tells the number of each
kind of atom in a molecule of the compound
• Structural formula: shows the bonds
connecting the atoms and the arrangement of
the atoms within each molecule
• Space-filling model: shows how the atoms in
the molecule are arranged in space (notice that
the methane molecule is not flat but is shaped
like a pyramid)
• CARBON:
–STRUCTURAL FORMULAS
• Shows the kind, number, and
arrangement of atoms in a
molecule
• A dash ( — ) is used to represent
the pair of shared electrons
forming the covalent bond
• CARBON:
–ISOMERS:
• Compounds with the same
molecular formula but different
structures
• Can have different physical and
chemical properties
• As the number of carbon atoms
increases, the number of isomers
increases
• HYDROCARBONS — contain only
hydrogen and carbon
–Most abundant source is Petroleum
–SATURATED: all bonds between
carbon atoms are single covalent
bonds
–UNSATURATED: one or more of
the bonds between carbon atoms
is a double covalent or triple
covalent bond
Macromolecules
• Many of the molecules in living cells are so
large that they are known as
macromolecules, which means “giant
molecules”
• Macromolecules are made from
thousands or even hundreds of
thousands of smaller molecules
Macromolecules
•
•
•
•
Macromolecules are formed by a process known as polymerization (pahlih-mur-ih-ZAY-shun), in which large compounds are built by joining smaller
ones together
The smaller units, or monomers, join together to form polymers
The monomers in a polymer may be identical, like the links on a metal
watch band; or the monomers may be different, like the beads in a
multicolored necklace
The figure below illustrates the formation of a polymer from more than
one type of monomer
Macromolecules
Macromolecules
• Four groups of organic compounds
found in living things are:
– Carbohydrates
– Lipids
– Nucleic acids
– Proteins
• Sometimes these organic compounds are
referred to as biomolecules
Carbohydrates
• Compounds made up of carbon, hydrogen, and
oxygen atoms, usually in a ratio of 1 : 2 : 1
• Living things use carbohydrates as their main
source of energy
• Plants and some animals also use carbohydrates for
structural purposes
• The breakdown of sugars, such as glucose, supplies
immediate energy for all cell activities
• Living things store extra sugar as complex
carbohydrates known as starches
• The monomers in starch polymers are sugar
molecules
CARBOHYDRATES
• Composed of carbon, hydrogen, and
oxygen
• Generalized formula: C x H 2 O
• Types:
– Monosaccharides
– Disaccharides
– Polysaccharides
CARBOHYDRATES
• Monosaccharides:
– Simple sugars
– Ratio: C H 2 O
– Most common (glucose, fructose, galactose) are isomers
• All have the same chemical molecular formula ( C 6 H 12 O 6 ) but
different structural formulas
– Glucose (dextrose):
» Produced by plants (photosynthesis)
» Main source of energy in plants and animals
» Metabolized in cellular respiration releasing energy
– Fructose:
» Found in fruits
» Sweetest of the monosaccharides
– Galactose:
» Found in milk
» Usually in combination with glucose and fructose making
disaccharides
CARBOHYDRATES
• Disaccharides:
– Double sugar
– Combination of two monosaccharides
• Formed by the chemical linking of two monosaccharides in a
condensation reaction (dehydration synthesis)
– Examples:
• Sucrose:
– Found in sugarcane and sugar beets
– Composed of fructose and glucose
• Lactose:
– Found in milk
– Composed of glucose and galactose
• Maltose:
– Malt sugar
CARBOHYDRATES
•
Polysaccharides:
– Complex molecule composed of three or more monosaccharides
– Formed by the chemical linking of three or more monosaccharides in
condensation reactions (dehydration synthesis)
– Examples:
• Starch:
– Storage form of glucose in plants
– Two basic forms:
» Long unbranched chains that coil like a telephone cord
» Highly branched like glycogen
• Glycogen:
– Storage form of glucose in animals
– Called animal starch
– Composed of hundreds of glucose molecules in a highly branched
chain
• Cellulose:
– Gives strength and rigidity to the plant cell
– Thousands of glucose monomers are linked in long, straight chains
CARBOHYDRATES
• The large macromolecules formed from
monosaccharides are known as polysaccharides
• Many animals store excess sugar in a polysaccharide
called glycogen, or animal starch
• When the level of glucose in your blood runs low,
glycogen is released from your liver
– Must be broken down to glucose before the body can utilize
the energy stored:
• Hydrolysis: splitting by the addition of water
• The glycogen stored in your muscles supplies the
energy for muscle contraction and, thus, for movement:
– Must be broken down to glucose before the body can utilize
the energy stored:
• Hydrolysis: splitting by the addition of water
CARBOHYDRATES
• Plants use a slightly different polysaccharide,
called plant starch, to store excess sugar:
– Must be broken down to glucose before the body
can utilize the energy stored:
• Hydrolysis: splitting by the addition of water
• Plants also make another important
polysaccharide called cellulose
• Tough, flexible cellulose fibers give plants much
of their strength and rigidity
• Cellulose is the major component of both wood
and paper, so you are actually looking at
cellulose when you are reading a textbook
Lipids
• Lipids are a large and varied group of biological
molecules that are generally not soluble in water
• Lipids are made mostly from carbon and hydrogen
atoms
• The common categories of lipids are fats, oils, and
waxes
• Lipids can be used to store energy
• Some lipids are important parts of biological
membranes and waterproof coverings
• Steroids are lipids as well
– Many steroids serve as chemical messengers
LIPIDS
• Fatty compound made up of a large number of carbon
and hydrogen atoms but a smaller number of oxygen
atoms
• Fats, oil, waxes, triglycerides, steroids
• Not soluble in water (insoluble)
• Major component of cell (plasma) membrane forming a
barrier between the internal and external aqueous
environments
• Store energy efficiently
• Large number of carbon-hydrogen bonds that store
more energy than carbon-oxygen bonds
LIPIDS
• Fatty acids:
– Monomers that make of lipids
– Long-straight hydrocarbon chain with a
carboxyl (COOH) group attached at one end
• Carboxyl end is polar which attracts water
which is polar (hydrophilic)
• Hydrogen end is nonpolar which tends to repel
water (hydrophobic)
• Cell membrane: the hydrophilic ends are
oriented to the aqueous side and the
hydrophobic ends are oriented to the center
LIPIDS
• Fatty acids:
– Saturated:
• All single bonds between carbon atoms
• No double bonds
• Maximum possible number of hydrogen atoms
bonded to each carbon atom
• Molecule is saturated with hydrogen
– Unsaturated:
• One or more double bonds between carbon atoms
• Fewer hydrogen atoms (not saturated)
Lipids
• Many lipids are formed when a glycerol molecule combines with
compounds called fatty acids, as shown below
• If each carbon atom in a lipid's fatty acid chains is joined to another
carbon atom by a single bond, the lipid is said to be saturated
• The term saturated is used because the fatty acids contain the
maximum possible number of hydrogen atoms
• The lipid represented has double bonds: unsaturated
Lipids
• If there is at least one carbon-carbon double bond in
a fatty acid, the fatty acid is said to be unsaturated
• Lipids whose fatty acids contain more than one
double bond are said to be polyunsaturated
• If the terms saturated and polyunsaturated seem familiar,
you have probably seen them on food package labels
• Lipids such as olive oil, which contains unsaturated fatty
acids, tend to be liquid at room temperature
• Cooking oils, such as corn oil, sesame oil, canola oil,
and peanut oil, contain polyunsaturated lipids
Lipids
LIPIDS
• Triglycerides:
– Lipid in which the macromolecule is composed of three
molecules of fatty acids joined by chemical condensation
reactions (dehydration synthesis) to one molecule of
glycerol
– Two main types:
• Oils: liquid triglyceride at room temperature
– Found mainly in plants (seeds)
» Source of stored energy
• Fats: solid triglycerides at room temperature
– Found mainly in animals
» Source of stored energy
LIPIDS
• Wax:
– Consists of long fatty acid chain joined to a
long alcohol chain
– Highly waterproof
– In plants, forms a protective covering on the
outer surfaces
– In animals, forms protective layer
• E.g. earwax: barrier that keeps microorganisms
from entering the middle ear
LIPIDS
• Steroid:
– Composed of four carbon rings
– No fatty acids
– Considered a lipid because they do not
dissolve in water (insoluble in water)
– Some hormones, nerve tissue, toad venoms,
and plant poisons
Nucleic Acids
•
•
•
•
Macromolecules containing hydrogen, oxygen, nitrogen, carbon, and
phosphorus
Nucleic acids are polymers assembled from individual monomers
known as nucleotides
Nucleotides consist of three parts: a 5-carbon sugar, a phosphate
group, and a nitrogenous base, as shown in the figure below
Individual nucleotides can be joined by covalent bonds to form a
polynucleotide, or nucleic acid
Nucleotide
Nucleic Acids
• The monomers that make up a nucleic acid are
nucleotides
• Each nucleotide has a 5-carbon sugar, a
phosphate group, and a nitrogenous base
• Nucleic acids store and transmit hereditary,
or genetic, information
• There are two kinds of nucleic acids:
– Ribonucleic acid (RNA): contains the sugar ribose
– Deoxyribonucleic acid (DNA): contains the sugar
deoxyribose
NUCLEIC ACIDS
• Complex organic molecules that store important
information in the cell
• Composed of thousands of monomers called
nucleotides
– Three components: phosphate group, five carbon
sugar, and a ring-shaped nitrogen base
• Two types:
– DNA (Deoxyribonucleic Acid)
• Stores information that is essential for almost all cell activities
• Replicated in cell division
– RNA (Ribonucleic Acid)
• Stores and transfers information that is essential for the
manufacturing of proteins
Proteins
• Macromolecules that contain nitrogen as
well as carbon, hydrogen, and oxygen
• Proteins are polymers of molecules
called amino acids
• Amino acids are compounds with an
amino group (−NH2) on one end and a
carboxyl group (−COOH) on the other
end
PROTEINS
• Organic compounds composed mainly of
hydrogen, oxygen, carbon, and nitrogen
• Formed from the linkage of monomers
called amino acids in a chemical
reaction of condensation (dehydration
synthesis)
• Structural and functional compounds of
importance in cells (plants and animals)
PROTEINS
• Amino acids
–
–
–
–
20 different types
Monomers that form proteins
Share the same basic structure
Each amino acid contains a central carbon atom to
which four other atoms or groups of atoms bond
covalently
•
•
•
•
A single hydrogen atom bonds at one site
A carboxyl group (COOH) bonds at a second site
A amine group (NH 2) bonds at the third site
A “R” group bonds at the fourth site
– The difference between amino acids results from different R
groups
Proteins
• The figure below shows one reason why proteins are among
the most diverse macromolecules
• More than 20 different amino acids are found in nature
• All amino acids are identical in the regions where they may be
joined together by covalent bonds
• This uniformity allows any amino acid to be joined to any other
amino acid — by bonding an amino group to a carboxyl group
Proteins
Proteins
• The portion of each amino acid that is different is a
side chain called an R-group
– Some R-groups are acidic and some are basic
– Some are polar and some are nonpolar
– Some contain carbon rings
• The instructions for arranging amino acids into
many different proteins are stored in DNA
• Each protein has a specific role
• Some proteins control the rate of reactions and
regulate cell processes
• Some are used to form bones and muscles
• Others transport substances into or out of cells or
help to fight disease
PROTEIN
• Dipeptides:
– Two amino acids chemically bonding together in a
condensation reaction (dehydration synthesis)
• The amino group of one amino acid releases a hydrogen ion
(H+) and the carboxyl group of the second amino acid
releases a hydroxide ion (OH -)
– Producing HOH (H2O)
• The nitrogen atom from the amine group and the carbon
atom from the carboxyl group bond covalently
• Covalent bond between the amine group of one amino acid
and the carboxyl group of another amino acids forms a
peptide bond
Proteins
•
Proteins can have up to four
levels of organization
– First: is the sequence of amino
acids in a protein chain
– Second: the amino acids within a
chain can be twisted or folded
– Third: the chain itself is folded
– Fourth: If a protein has more
than one chain
• Each chain has a specific
arrangement in space as shown
by the red and blue structures
in the figure at right
•
•
Van der Waals forces and
hydrogen bonds help maintain a
protein's shape
In the next section, you will learn
why a protein's shape is so
important
Proteins
Chemical Reactions and Enzymes
• Chemical Reactions:
– Process that changes one set of chemicals into another set
of chemicals
– Some chemical reactions occur slowly, such as the
combination of iron and oxygen to form an iron oxide called
rust
– Other reactions occur quickly
– When hydrogen gas is ignited in the presence of oxygen,
the reaction is rapid and explosive
– The elements or compounds that enter into a chemical
reaction are known as reactants
– The elements or compounds produced by a chemical
reaction are known as products.
– Chemical reactions always involve the breaking of bonds in
reactants and the formation of new bonds in products
Chemical Reactions
• One example of an important chemical reaction that
occurs in your body involves carbon dioxide
• Your cells constantly produce carbon dioxide as a
normal part of their activity
• This carbon dioxide is carried to your lungs through the
bloodstream, and then is eliminated as you exhale
• However, carbon dioxide is not very soluble in water. The
bloodstream could not possibly dissolve enough carbon
dioxide to carry it away from your tissues were it not for a
chemical reaction
• As it enters the blood, carbon dioxide reacts with water
to produce a highly soluble compound called carbonic
acid, H2CO3
Chemical Reactions
• CO2
+
H 2O
→
H2CO3
• The reaction shown above enables the
bloodstream to carry carbon dioxide to
the lungs
• In the lungs, the reaction is reversed:
• H2CO3 → CO2
+
H2O
– This reverse reaction produces carbon
dioxide gas, which is released as you
exhale
Energy in Reactions
• Energy is released or absorbed whenever
chemical bonds form or are broken
• Because chemical reactions involve
breaking and forming bonds, they involve
changes in energy
Energy Changes
• Some chemical reactions release energy (exothermic),
and other reactions absorb energy (endothermic)
– Energy changes are one of the most important factors in
determining whether a chemical reaction will occur
• Chemical reactions that release energy often occur
spontaneously
• Chemical reactions that absorb energy will not occur
without a source of energy
• An example of an energy-releasing reaction is hydrogen
gas burning, or reacting, with oxygen to produce water
vapor :
– 2H2
+
O2
→
2H2O
• The energy is released in the form of heat, and sometimes—when
hydrogen gas explodes—light and sound
Energy Changes
• 2H2O → 2H2 + O2
• The reverse reaction, in which water is changed
into hydrogen and oxygen gas, absorbs so
much energy that it generally doesn't occur
by itself
• In fact, the only practical way to reverse the
reaction is to pass an electrical current through
water to decompose water into hydrogen gas
and oxygen gas
– Thus, in one direction the reaction produces
energy, and in the other direction the reaction
requires energy
Energy Changes
• What significance do these energy changes have for
living things?
– In order to stay alive, organisms need to carry out
reactions that require energy
• Thus, every organism must have a source of
energy to carry out chemical reactions
• Plants get their energy by trapping and storing the
energy from sunlight in energy-rich compounds
(Photosynthesis)
• Animals get their energy when they consume plants or
other animals (Respiration)
• Humans release the energy needed to grow tall, to
breathe, to think, and even to dream through the
chemical reactions that occur when humans metabolize,
or break down, digested food
Activation Energy
•
Even chemical reactions that release
energy do not always occur
spontaneously
–
•
The cellulose in paper burns in the
presence of oxygen and releases heat
and light
–
•
•
That's a good thing because if they did,
the pages in textbooks might burst into
flames
However, the cellulose will burn only if
you light it with a match, which supplies
enough energy to get the reaction
started
Chemists call the energy that is
needed to get a reaction started the
activation energy
As the figure to the right shows,
activation energy is a factor in
whether the overall chemical
reaction releases energy or absorbs
energy
Activation Energy
Activation Energy
• Chemical reactions that
release energy often occur
spontaneously
• Chemical reactions that
absorb energy will occur
only with a source of energy
• The peak of each graph
represents the energy
needed for the reaction to go
forward
• The difference between this
required energy and the
energy of the reactants is
the activation energy
Enzymes
• Some chemical reactions that make life
possible are too slow or have activation
energies that are too high to make them
practical for living tissue
• These chemical reactions are made possible
by a process that would make any chemist
proud—cells make catalysts
– A catalyst is a substance that speeds up the rate
of a chemical reaction
– Catalysts work by lowering a reaction's activation
energy
PROTEIN
• Enzymes
– Organic chemical catalyst
• Speeds up chemical reactions without being affected
(consumed) by the reactions themselves
– Proteins that act as catalysts in intermediary
metabolism
– Essential for the functioning of any cell
– Speed up a chemical reaction by using an alternate
way to change reactant to product
• Reaction catalyzed by the enzyme has a lower activation
energy than the reaction without the enzyme
ENZYMES
• Proteins that act as
biological catalysts
• Enzymes speed up
chemical reactions that
take place in cells
• Like other catalysts,
enzymes act by
lowering the activation
energies, as illustrated
by the graph to the right
ENZYMES
ENZYMES
• Lowering the activation energy has a
dramatic effect on how quickly the reaction
is completed
• How big an effect does it have?
– Consider the reaction in which carbon dioxide
combines with water to produce carbonic
acid:
• CO2
+
H 2O
→
H2CO3
Effect of Enzymes
•
•
Enzymes speed up chemical
reactions that take place in cells
Notice how the addition of an
enzyme lowers the activation
energy in this reaction
– This action speeds up the
reaction:
•
•
– CO2
+
H2O
→
H2CO3
Left to itself, this reaction is so
slow that carbon dioxide might
build up in the body faster than
the bloodstream could remove
it
Your bloodstream contains an
enzyme called carbonic
anhydrase that speeds up the
reaction by a factor of 10 million
Effect of Enzymes
• With carbonic anhydrase on the job, the
reaction takes place immediately and carbon
dioxide is removed from the blood quickly
• Enzymes are very specific, generally
catalyzing only one chemical reaction
– For this reason, part of an enzyme's name is usually
derived from the reaction it catalyzes
– Carbonic anhydrase gets its name because it
catalyzes the reaction that removes water from
carbonic acid
Enzyme Action
• How do enzymes do their jobs?
• For a chemical reaction to take place,
the reactants must collide with enough
energy so that existing bonds will be
broken and new bonds will be formed
– If the reactants do not have enough energy,
they will be unchanged after the collision
PROTEIN
• Enzyme
– Shape of a particular enzyme allows it to
hook up with a specific molecule
(substrate) (reactant)
• Linkage of the enzyme and the substrate
weakens some chemical bonds in the substrate
lowering the activation energy producing the
products which the enzymes releases
• Enzyme is unaltered and able to combine with
another substrate continuing the chemical
reactions
The Enzyme-Substrate Complex
• Enzymes provide a site where reactants
can be brought together to react
• Such a site reduces the energy needed
for reaction
• The reactants of enzyme-catalyzed
reactions are known as substrates
The Enzyme-Substrate Complex
•
•
•
•
•
•
•
•
The figure at right provides an
example of an enzyme-catalyzed
reaction
The enzyme is hexokinase
The substrates are glucose and ATP
During the reaction, a phosphate
group is transferred from ATP to the
glucose molecule
Recall that each protein has a
specific, complex shape
The substrates bind to a site on the
enzyme called the active site
The active site and the substrates
have complementary shapes
The fit is so precise that the active
site and substrates are often
compared to a lock and key
The Enzyme-Substrate Complex
The Enzyme-Substrate Complex
• The figure to the right shows
a substrate fitting into an
active site on an enzyme
• The enzyme and substrate are
bound together by
intermolecular forces and
form an enzyme-substrate
complex
• They remain bound together
until the reaction is done
• Once the reaction is over,
the products of the reaction
are released and the enzyme
is free to start the process
again
The Enzyme-Substrate Complex
Regulation of Enzyme Activity
• Because they are catalysts for reactions,
enzymes can be affected by any variable that
influences a chemical reaction
• Enzymes, including those that help digest food,
work best at certain pH values
• Many enzymes are affected by changes in
temperature
• Not surprisingly, those enzymes produced by
human cells generally work best at
temperatures close to 37°C, the normal
temperature of the human body
Regulation of Enzyme Activity
• Cells can regulate the activities of
enzymes in a variety of ways
• Most cells contain proteins that help to
turn key enzymes “on” or “off” at critical
stages in the life of the cell
• Enzymes play essential roles in regulating
chemical pathways, making materials that
cells need, releasing energy, and
transferring information