Transcript Chemistry of Life - Dr. Wilson`s Site

Anatomy & Physiology I
Chapter 2
Science that deals with composition and properties
of matter
Necessary to understand normal and abnormal
functioning of body
Biochemistry – the study of the molecules that
compose living organisms
carbohydrates, fats, proteins, and nucleic acids
Matter - Anything that has mass and occupies space.
All matter is composed of elements.
Elements
Make up ALL matter
Cannot be broken down by ordinary chemical means
Each has unique physical and chemical properties
92 occur in nature (24 elements have biological role)
Identified by names, chemical symbols or number
Described and organized in periodic table
Oxygen (O)
Carbon (C)
Hydrogen (H)
Nitrogen (N)
About 96% of body mass
About 3.9% of body mass:
Calcium (Ca), phosphorus (P), potassium (K), sulfur
(S), sodium (Na), chlorine (Cl), magnesium (Mg),
iodine (I), and iron (Fe)
Trace Elements (< 0.01% of body mass):
Part of enzymes, e.g., chromium (Cr), manganese
(Mn), and zinc (Zn)
Atoms
Unique building blocks for each element
Smallest complete units of matter
Cannot be broken down or changed by
ordinary chemical and physical means
Determined by numbers of subatomic particles
Protons (positive charge [+]); in nucleus
Neutrons (no charge); in nucleus
Electrons (negative charge [–]); in concentric clouds that
surround nucleus
determine the chemical properties of an atom
the atom is electrically neutral because number of electrons is
equal to the number of protons
valence electrons in the outermost shell
determine chemical bonding properties of an atom
Atoms of different elements contain different
numbers of subatomic particles
Compare hydrogen, helium and lithium (next slide)
Proton
Neutron
Electron
Hydrogen (H)
(1p+; 0n0; 1e–)
Helium (He)
(2p+; 2n0; 2e–)
Lithium (Li)
(3p+; 4n0; 3e–)
Isotopes – varieties of an element that differ
from one another only in the number of
neutrons and therefore in atomic mass
extra neutrons increase atomic weight
isotopes of an element are chemically similar
have same valence electrons
Atomic weight of an element accounts for the
fact that an element is a mixture of isotopes
2-10
Hydrogen (1H)
(1p+, 0n0, 1e–)
Deuterium (2H)
(1p+, 1n0, 1e–)
Key
= Proton
= Neutron
= Electron
Tritium (3H)
(1p+, 2n0, 1e–)
Isotopes
same chemical behavior, differ in physical behavior
breakdown (decay) to more stable isotope by giving off
radiation
Radioisotopes
unstable isotopes that give off radiation
every element has at least one radioisotope
Similar chemistry to stable isotopes
Can be detected with scanners
Radioactivity
radioisotopes decay to stable isotopes releasing radiation
we are all mildly radioactive
• Ions – charged particles with unequal number of protons
and electrons
Ionization - transfer of
electrons from one
atom to another
( stability of valence
shell)
11 protons
12 neutrons
11 electrons
Sodium
atom (Na)
17 protons
18 neutrons
17 electrons
Chlorine
atom (Cl)
Transfer of an electron from a sodium atom to a chlorine atom
• Anion
– atom that gained electrons (net negative charge)
• Cation
– atom that lost an electron (net positive charge)
• Ions with opposite charges are attracted to each other
+
–
11 protons
12 neutrons
10 electrons
Sodium
ion (Na+)
17 protons
18 neutrons
18 electrons
Chloride
ion (Cl–)
Sodium chloride
2 The charged sodium ion (Na+) and chloride ion (Cl–) that result
Salts that ionize in water and form solutions capable of
conducting an electric current.
Electrolyte importance
chemical reactivity
osmotic effects (influence water movement)
electrical effects on nerve and muscle tissue
Electrolyte balance is one of the most important
considerations in patient care.
Imbalances have ranging effects from muscle cramps,
brittle bones, to coma and cardiac arrest
Chemical particles with an odd number of electrons
Produced by
normal metabolic reactions, radiation, chemicals
Causes tissue damage
reactions that destroy molecules
causes cancer, death of heart tissue and aging
Antioxidants
neutralize free radicals
in body, superoxide dismutase (SOD)
in diet (Selenium, vitamin E, vitamin C, carotenoids)
Molecules
Formed when two or more atoms unite on the
basis of their electron structures
Can be made of like atoms or atoms of different
elements
Compounds
Composed of two or more elements
Smallest subunits of a compound are molecules
Most matter exists as mixtures
Two or more components physically
intermixed
Three types of mixtures
Solutions
Colloids
Suspensions
Homogeneous mixtures
Usually transparent, e.g., atmospheric air
or seawater
Solvent
Present in greatest amount, usually a liquid
Solute(s)
Present in smaller amounts
Colloids (emulsions)
Heterogeneous translucent mixtures, e.g.,
cytosol
Large solute particles that do not settle out
Undergo sol-gel transformations
Suspensions:
Heterogeneous mixtures, e.g., blood
Large visible solutes tend to settle out
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
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
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)
Stable and unreactive
Outermost energy level fully occupied or
contains eight electrons
(a)
Chemically inert elements
Outermost energy level (valence shell) complete
8e
2e
Helium (He)
(2p+; 2n0; 2e–)
2e
Neon (Ne)
(10p+; 10n0; 10e–)
Outermost energy level not fully occupied by
electrons
Tend to gain, lose, or share electrons (form
bonds) with other atoms to achieve stability
(b)
Chemically reactive elements
Outermost energy level (valence shell) incomplete
1e
Hydrogen (H)
(1p+; 0n0; 1e–)
6e
2e
Oxygen (O)
(8p+; 8n0; 8e–)
4e
2e
Carbon (C)
(6p+; 6n0; 6e–)
1e
8e
2e
Sodium (Na)
(11p+; 12n0; 11e–)
Ionic
Covalent
Hydrogen
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
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.
1.
2.
3.
(b) After electron transfer, the oppositely
charged ions formed attract each other.
An electron is transferred from the Na atom to Cl atom.
The sodium (Na) becomes a cation (Na+) and the chlorine (Cl) becomes an
anion (Cl-).
The opposite charges attract forming an ionic bond.
In step 2, why does sodium become positive and chlorine negative?
CI–
Na+
(c) Large numbers of Na+ and Cl– ions
associate to form salt (NaCl) crystals.
Formed by two atoms sharing valence electrons
Types of covalent bonds
single - sharing of single pair electrons
double - sharing of 2 pairs of electrons
nonpolar covalent bond
shared electrons spend approximately equal time around
each nucleus
strongest of all bonds
polar covalent bond
if shared electrons spend more time orbiting one nucleus
than they do the other, they lend their negative charge to the
area they spend most time
Reacting atoms
Resulting molecules
+
Hydrogen
atoms
or
Carbon
atom
Molecule of
methane gas (CH4)
Structural
formula
shows
single
bonds.
Formation of four single covalent bonds:
Each hydrogen atom shares its electron with carbon while carbon shares one of
its valence electrons with each hydrogen atom.
Reacting atoms
Resulting molecules
+
Oxygen
atom
or
Oxygen
atom
Molecule of
oxygen gas (O2)
Structural
formula
shows
double
bond.
Formation of a double covalent bond:
Each oxygen atom shares two electrons with the other oxygen atom.
Sharing of electrons may be equal or unequal
Equal sharing produces electrically balanced
nonpolar molecules
CO2
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
Attraction between a slightly positive hydrogen
atom in one molecule and a slightly negative
oxygen or nitrogen atom in another.
Water molecules are weakly attracted to each
other by hydrogen bonds
very important to physiology
protein structure
DNA structure++
H
O
H
H
H
O
O
H
H
O
H
Covalent bond
H
Hydrogen bond
O
Water molecule
H
H
Water’s polar covalent bonds and its V-shaped
molecule gives water a set of properties that account for
its ability to support life.
60%–80% of the volume of living cells
Most important inorganic compound in living organisms
because of its properties:
solvency
cohesion
adhesion
chemical reactivity
thermal stability
Solvency - ability to dissolve other chemicals
water is called the Universal Solvent
Hydrophilic – substances that dissolve in water
molecules must be polarized or charged
Hydrophobic - substances that do not dissolve in
water
molecules are non-polar or neutral (fat)
Virtually all metabolic reactions depend on the
solvency of water
A water strider can walk on a pond because of the high
surface tension of water, a result of the combined
strength of its hydrogen bonds.
is the ability to participate in chemical
reactions
water ionizes into H+ and OH-
water ionizes other chemicals (acids and salts)
water involved in hydrolysis and dehydration
synthesis 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
H + H  H2 (hydrogen gas)
(reactants)
(product)
4H + C  CH4 (methane)
Synthesis (combination) reactions
Decomposition reactions
Exchange reactions
Synthesis reactions
Smaller particles are bonded
together to form larger,
more complex molecules.
• A + B  AB
• Always involve bond
formation
• Anabolic
Example
Amino acids are joined together to
form a protein molecule.
Amino acid
Molecules
(monomers)
Protein
Molecule
(polymer)
Decomposition reactions
• AB  A + B
• Reverse synthesis
reactions
• Involve breaking of
bonds
Bonds are broken in larger
molecules, resulting in smaller,
less complex molecules.
Example
Glycogen is broken down to release
glucose units.
Glycogen
(polymer)
• Catabolic
Glucose
Molecules
(monomers)
AB + C  AC + B
Also called displacement reactions
Bonds are both made and broken
An acid is proton donor (releases H+ ions in
water)
A base is proton acceptor (accepts H+ ions)
releases OH- ions in water
pH – a measure derived from the molarity of H+
a pH of 7.0 is neutral pH
(H+ = OH-)
a pH of less than 7 is acidic solution (H+ > OH-)
a pH of greater than 7 is basic solution (OH- > H+ )
Acids and bases are electrolytes
pH - measurement of molarity of H+ [H+] on a logarithmic
scale
a change of one number on the pH scale represents a 10
fold change in H+ concentration
a solution with pH of 4.0 is 10 times as acidic as one with pH of
5.0
Our body uses buffers to resist changes in pH
slight pH disturbances can disrupt physiological functions and
alter drug actions
pH of blood ranges from 7.35 to 7.45
deviations from this range cause tremors, paralysis or even
death
Gastric juice
(0.9–3.0)
Lemon
juice
1M
(2.3)
Hydrochloric
Acid (0)
Wine,
Bananas,
vinegar
tomatoes
(2.4 -–3.5)
(4.7)
3
2
1
0
Bread,
black
coffee
(5.0)
4
Pure water
Milk,
Egg white
(7.0)
saliva
(8.0)
(6.3 -–6.6)
Household
bleach
(9.5)
Household
ammonia
(10.5 - 11.0)
Oven cleaner, lye
(13.4)
1 M sodium
hydroxide
(14)
5
6
7
Neutral
8
9
10
11
12
13
14
Acid solutions contain [H+]
As [H+] increases, acidity increases
Alkaline solutions contain bases (e.g., OH–)
As [H+] decreases (or as [OH–] increases), alkalinity
increases
Neutral solutions:
Pure water is pH neutral (contains equal numbers
of H+ and OH–)
All neutral solutions are pH 7
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
Mixture of compounds that resist pH changes
Convert strong acids or bases into weak ones
Carbonic acid-bicarbonate system
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
Must contain at least one carbon and one
hydrogen atom. (CO2 and CO are inorganic
because they have no hydrogen atom.)
Unique to living systems
Include carbohydrates, lipids, proteins, and
nucleic acids (Molecules of Life)
Many are polymers—chains of similar units
(monomers or building blocks)
Synthesized by dehydration synthesis
Broken down by hydrolysis reactions
Macromolecules - very large organic molecules
proteins, DNA
Polymers – molecules made of a repetitive series
of identical or similar subunits (monomers)
Monomers - an identical or similar subunits
joining monomers to form a polymer
dehydration synthesis (condensation) is how living cells
form polymers
a hydroxyl (-OH) group is removed from one monomer, and a
hydrogen (H+) from another
producing water as a by-product
hydrolysis – opposite of dehydration synthesis; breaks
down polymers to monomers
a water molecule ionizes into –OH and H+
the covalent bond linking one monomer to the other is broken
the –OH is added to one monomer
the H+ is added to the other
Monomers covalently bond together to form a
polymer with the removal of a water molecule
A hydroxyl group is removed from one monomer and a
hydrogen from the next
Dimer
Monomer 1
Monomer 2
OH
O
HO
H+ + OH–
(a) Dehydration synthesis
H2 O
Splitting a polymer (lysis) by the addition of a water molecule
(hydro)
a covalent bond is broken
All digestion reactions consists of hydrolysis reactions
Dimer
Monomer 1
OH
O
H2 O
(b) Hydrolysis
Monomer 2
H+ + OH–
HO
Sugars and starches
Contain C, H, and O [(CH20)n]
Three classes
Monosaccharides
Disaccharides
Polysaccharides
Building blocks of carbohydrates: monosaccharides
Simple sugars containing three to seven C
atoms (CH20)n
Glucose – blood sugar; major source of cellular
fuel
Fructose - found in fruits, vegetables and honey
Galactose – component of some glycolipids (i.e.
the ABO antigens that determine blood type)
Ribose in RNA & Deoxyribose in DNA
Double sugars (made of 2 monosaccharides)
Maltose – malt sugar (used to make beer)
made of 2 glucose monomers
Sucrose – table sugar (cane sugar)
made of glucose and fructose
Lactose – milk sugar
made of glucose and galactose
(b) Disaccharides
Consist of two linked monosaccharides
Example
Sucrose, maltose, and lactose
(these disaccharides are isomers)
Glucose
Fructose
Sucrose
Glucose
Glucose
Maltose
Galactose Glucose
Lactose
Polymers of many simple sugars
Starch – storage form of excess glucose in plants (made
of several hundred to several thousand glucose units)
Glycogen – storage form of excess glucose in animals
(made of several thousand glucose units)
The liver and skeletal muscles are major depots of glycogen.
Glycogen is broken back down into glucose when energy is needed
(a process called glycogenolysis).
Cellulose – cell walls of plants; provides dietary fiber,
roughage (made of approx. 10,000 glucose units)
(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
2-70
Contain C, H, O (less than in carbohydrates),
and sometimes P
Insoluble in water
Main types:
Triglycerides (neutral fats)
Phospholipids
Steroids
Eicosanoids
Building blocks of lipids: glycerol and fatty acids
Chain of 4 to 24 carbon atoms
carboxyl (acid) group on one end, methyl group on the other
and hydrogen bonded along the sides
Classified
saturated - carbon atoms saturated with hydrogen
unsaturated - contains C=C bonds without hydrogen
polyunsaturated – contains many C=C bonds
essential fatty acids – obtained from diet, body can not
synthesize
(saturated)
3 fatty acids covalently bonded to three carbon
alcohol, glycerol molecule
each bond formed by dehydration synthesis
broken down by hydrolysis
triglycerides at room temperature
when liquid called oils
often polyunsaturated fats from plants
when solid called fat
saturated fats from animals
Primary Function - energy storage, insulation and
shock absorption (adipose tissue)
similar to neutral fat except
that one fatty acid replaced
by a phosphate group
CH3
N+
CH3
Nitrogencontaining
group
(choline)
CH2
structural foundation
of cell membrane
CH2
O
–O
Amphiphilic
fatty acid “tails” are
hydrophobic
phosphate “head” is
hydrophilic
CH3
P
O
Phosphate
group
Hydrophilic region
(head)
O
Glycerol
O
CH2
CH
O
O
C
C
(CH2)5
CH
CH
(CH2)5
CH3
(CH2)12
CH3
CH2
O
Fatty acid
tails
Hydrophobic region
(tails)
Steroid – a lipid with 17 of its carbon atoms in
four rings
Cholesterol - the ‘parent’ steroid from which the
other steroids are synthesized
cortisol, progesterone, estrogens, testosterone and
bile acids
Cholesterol
synthesized only by animals
especially liver cells
15% from diet, 85% internally synthesized
important component of cell membranes
one kind of cholesterol
does far more good than harm
‘good’ and ‘bad’ cholesterol actually refers to droplets
of lipoprotein in the blood
complexes of cholesterol, fat, phospholipid, and protein
HDL – high-density lipoprotein – “good” cholesterol
lower ratio of lipid to protein
may help to prevent cardiovascular disease
LDL – low-density lipoprotein – “bad” cholesterol
high ratio of lipid to protein
contributes to cardiovascular disease
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)
20 carbon compounds derived from a fatty acid called
arachidonic acid
hormone-like chemical signals between cells
includes prostaglandins – produced in all tissues
role in inflammation, blood clotting, hormone action, labor
contractions, blood vessel diameter
O
COOH
OH
OH
2-79
Polymers of amino acids (20 types)
Joined by peptide bonds
Contain C, H, O, N, and sometimes S and P
Building blocks of proteins: amino acids
peptide – any molecule composed of two or more amino
acids joined by peptide bonds
peptide bond – joins two amino acids
formed by dehydration synthesis
Peptides named for the number of amino acids
dipeptides have 2
tripeptides have 3
oligopeptides have fewer than 10 to 15
polypeptides have more than 15
proteins have more than 50
Amine
group
Acid
group
(a) Generalized
structure of all
amino acids.
(b) Glycine
is the simplest
amino acid.
(c) Aspartic acid
(d) Lysine
(an acidic amino acid)
(a basic amino acid)
has an acid group
has an amine group
(—COOH) in the
(–NH2) in the R group.
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.
• There are 20 naturally occurring amino acids.
• All 20 amino acids are identical except for their R group. (notice
the structures highlighted green)
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
Hydrolysis: Peptide
bonds linking amino
acids together are
broken when water is
added to the bond.
Dipeptide
Amino acid
Amino acid
Amino acid
Amino acid
Amino acid
The sequence of amino acids forms the polypeptide chain.
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)
Enzymes - proteins that function as biological
catalysts
permit reactions to occur rapidly at normal body
temperature
Substrate - substance an enzyme acts upon
Naming Convention
named for substrate with -ase as the suffix
amylase enzyme digests starch (amylose)
All enzymes are proteins
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)
Enzyme-substrate
complex (E-S)
1 Substrates bind
2 Internal
at active site.
rearrangements
Enzyme changes
leading to
shape to hold
catalysis occur.
substrates in
proper position.
Enzyme (E)
3
Product is
released. Enzyme
returns to original
shape and is
available to catalyze
another reaction.
DNA and RNA
Largest molecules in the body
Contain C, O, H, N, and P
Building block = nucleotide, composed of
nitrogenous base, a pentose sugar, and a
phosphate group
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
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)
(c) Computer-generated image of a DNA molecule
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
Adenine-containing RNA nucleotide with two
additional phosphate groups
High-energy phosphate
bonds can be hydrolyzed
to release energy.
Adenine
Phosphate groups
Ribose
Adenosine
Adenosine monophosphate (AMP)
Adenosine diphosphate (ADP)
Adenosine triphosphate (ATP)
Phosphorylation:
Terminal phosphates are enzymatically transferred
to and energize other molecules
Such “primed” molecules perform cellular work
(life processes) using the phosphate bond energy
Solute
+
Membrane
protein
(a) Transport work: ATP phosphorylates transport
proteins, activating them to transport solutes
(ions, for example) across cell membranes.
+
Relaxed smooth
muscle cell
Contracted smooth
muscle cell
(b) Mechanical work: ATP phosphorylates
contractile proteins in muscle cells so the
cells can shorten.
+
(c) Chemical work: ATP phosphorylates key
reactants, providing energy to drive
energy-absorbing chemical reactions.