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Chapter 1: An Introduction to Physiology
-physiology: the study of the functions of living things
-the human body is comprised of 11 major organ systems:
- in studying these systems we can use two approaches:
1.
teleological
-emphasizes the purpose of a body process
-explains a function in terms of meeting a bodily need
-emphasizes the WHY
e.g. “why do I shiver?” – to warm up because
shivering generates heat
2.
mechanistic
-emphasizes the underlying mechanism by which
this process occurs
-view the body as a machine whose actions can be explained
in terms of cause and effect
-emphasize the HOW
e.g. “why do I shiver?” – detection of body temperature by sensory
receptors leads to activation of the somatic division of the nervous
system and trigger the involuntary contraction of skeletal muscles
1. skeletal
2. articular
3. muscular
4. digestive
5. respiratory
6. urinary
7. reproductive
8. circulatory
9. nervous
10. integumentary
11. endocrine
-physiology is closely related to anatomy – because structure and function are closely related
-physiological mechanisms are made possible because of the structure and design
of a body part
-some relationships are obvious – e.g. structure of the elbow as a hinge joint
-others are more subtle – e.g. interface between the air and the blood in the lungs
-air sac structure + capillary bed
- 300 million air sacs + associated capillaries provides
a total surface area for gas exchange = tennis court
Life: Levels of Organization
•Atoms
•Molecules
•Macromolecules
•Organelles
•Cells
•Tissues
•Organs
•Organ systems
•Organism
Organizational Levels
1. chemical or molecular: 4 major elements within the body
-99% of the total number of atoms within the body
-C, N, O and H
-molecular composition - 67% of our bodies is water
Atom = smallest unit of an element that still retains the chemical &
physical properties of that element
i.e. really, really, really tiny thing!
-composed of: protons = one positive charge, 1 atomic mass unit (1.673x10-24g)
electrons = one negative charge, no mass (9.109x10-28g)
neutrons = no charge, 1 atomic mass unit
-elements are grouped on a Periodic Table of Elements
-the elements are grouped according to physical and chemical
characteristics
-on the chart each element is associated with a letter, an atomic number
& an atomic mass
-each atom is comprised of a nucleus of protons and neutrons
+ orbiting electrons
Periodic Table of Elements
IA
IIA
http://www.chemicalelements.com/
http://periodic.lanl.gov/default.htm
IIIA
IVA VA VIA VIIA
VIII
atomic
symbol
atomic
mass (weight)
12
6
C
e.g. # protons (e-) = 6
# pr+6 + #No 6 = 12
atomic
number
7
3
Li
e.g. # protons (e-) = 3
# pr+3 + #No 4 = 7
Atomic mass = number of protons + neutrons
Atomic number = number of protons when the element is electrically
neutral
** when neutral, the number of protons and electrons are equal
Isotope:
• same atomic # (same pr+, same e-)
• differs only in # of neutrons
pr+:
e-:
No:
12C
13C
6
6
6
6
6
7
14C
6
6
8** radioactive
Radioactive isotope uses:
-radioactive isotopes have a high neutron to proton ratio
-the nuclear “glue” within the nucleus is not strong enough to hold the nucleus
together = radioactivity
1. carbon dating - 14C
2. radioactive imaging - e.g. PET scanning
-use of FDG – radioactive glucose tracer
-18F radioactive isotope (2-fluoro-deoxy-glucose)
3. cancer treatment - 60Co
Chemical bonds
-forces holding atoms together = chemical bonds
-different kinds of chemical bonds – but all involve the electrons of atoms
Two types of bonds:
1. Ionic
2. Covalent
Electron Configurations
• Bed check for electrons
• description on how are electrons organized around the nucleus
of protons and neutrons?
• Bohr model: Nils Bohr proposed electrons orbit around the
atom’s nucleus in specific energy levels or orbits (shells)
– these shells have a specific energy level – closer the electron is to the
nucleus the less energy it needs to “orbit”
– this model only works for smaller atoms
– larger atoms are described by quantum mechanics – orbitals have
energy, momentum/shape, spin and magnetic characteristics
– comprised of subshells
– 1st shell – closest to the nucleus only holds 2 electrons (s subshell only)
– 2nd shell can hold 8 (s and p subshells – 2 + 6 electrons)
– 3rd holds 18 (s, p and d subshells – 2 + 6 + 10 electrons)
– 4th holds 18
•an atom will always try to complete its outermost shell
•basis for bonding reactions
•the number of electrons the atom gains or loses to complete its outer shell =
valence
•chemists really only consider the electrons in the s and p orbitals as valence
electrons
Molecule:
•particle formed by the union of more than one atom
e.g. same kind of atom - O2
e.g. different types of atoms - H20
Molecules are held together by either covalent
or ionic bonds
-these bonds form through interactions between the
valence electrons
1. Ionic bond:
•
attraction between 2 oppositely charged atoms (ions)
e.g. Na+ and ClNaCl
e.g. Ca2+ and Cl-
CaCl2
-positively charged ions = cations
-negatively charged ions = anions
-these form as one atom transfers electrons to another atom
-ions may also be composed of more then one atom = polyatomic
-these are treated as the same as monoatomic ions
e.g. sulfate SO43-, nitrite NO2-, hydroxide OH-
Ionic Reaction
-most of your group I and group II
metals will form ionic bonds with the
group VIIA halogens
2. Covalent Bond:
•if it isn’t favorable for an atom to gain or lose an electron
it will have to share it with another
•covalent bond = bond in which atoms share electrons
e.g O2, , N3
e.g H20
-usually forms when one atom has to lose or gain three or more
electrons
e.g. carbon would have to gain 4 valence electrons to complete its outer shell
nitrogen would have to gain 3 valence electrons
-also form between two identical atoms – e.g. nitrogen, oxygen gas
Polar and Non-polar bonds
• the sharing of electrons does not have to be equal
• nonpolar covalent bond = equal sharing of electrons
– e.g. oxygen (O2), methane (CH4)
• polar covalent bond = uneven sharing of electrons leading
to a slight charge
– e.g. water = H20
d-
Water:
• 60-70% of body weight
• covalent bond
• POLAR molecule (uneven sharing of
electrons)
d+
O
H
H
d+
•polar compounds are attracted to other
•the bond between one oxygen and
the hydrogens of adjacent water
molecules = Hydrogen Bond
**HB = occurs between a covalently
bonded hydrogen and negatively
charged atom a distance away
Chemical reactions:
3 types:
1. Synthesis - A + B
AB (Anabolism reactions)
2. Decomposition - AB
3. Exchange - AB + CD
A + B (Catabolism reactions)
AD + BC
-these equations must be balanced
-Law of conservation of Mass or “chemical bookeeping”
-i.e. the number of atoms of each element is the same before
and after a chemical reaction
Chemical reactions
• are made up of reactants and creates products
• these reactions go on constantly within the human body
= metabolism
• each reaction involves changes in energy
– if the reaction requires energy = endothermic (anabolism)
– if it liberates energy = exothermic (catabolism)
• atoms, molecules and ions are continuously moving
and colliding with one another = kinetic energy
– this kinetic energy if big enough (i.e. collision is large
enough) can break a bond or cause a new one to form
– this collision energy = activation energy
– critical to the progression of all chemical reactions in our
body
– the more often a collision occurs the greater chance a bond
will form or break
Activation Energy & Catalysts
• activation energy = initial “energy
investment” required to start a
reaction
– the reactants must absorb enough
energy to cause their chemical bonds
to become unstable and created new
ones
– as these bonds form – energy is
released into the environment – if
more energy is released than
absorbed = heat (exothermic
reaction)
– two influences on AE – temperature
and concentration
• concentration – increasing this
increases the chance of collision
between atoms
• temperature – heating a reaction
increases the kinetic energy of the
reactants – collide more often
• catalysts = compounds that lower the
activation energy of a reaction
Molecules of Life:
• the chemicals used in metabolic reactions or those
that are produced by them can be classified into
2 groups:
1. Inorganic
2. Organic
Inorganic Compounds
 water
 oxygen,carbon dioxide
 inorganic salts
Water
• major component of blood, plasma, CSF etc…
• role in: transporting chemicals
transporting waste products
transporting & absorbing heat
• polar molecule - asymmetrical distribution of charge
• liquid at room temperature – we can drink it
•universal solvent for polar compounds – facilitates most chemical
reactions in the body
-water molecules are cohesive – therefore they cling together
(because of hydrogen bonding)
-this allows the even distribution of dissolved substances
throughout our system – so water is an excellent transport
medium
-the temperature of water rises and falls slowly – prevents
sudden and drastic changes of temperature in our bodies
-water requires high heat to evaporate – it cools our bodies
-frozen water is less dense than liquid water – ice floats
-water freezes from the top down – allows aquatic organisms
to survive winters
d-
O
H
d+
H
d+
Water:
-excellent solvent for dissolution of polar and ionic substances
e.g. H20 + NaCl
-in a solution – a solvent dissolves another substance called
a solute
-water is a versatile solvent because of its polar covalent bonds and its bent
shape which allows it to interact with its neighbours very well
-water + salt: the electronegative O of
water attracts the +ve sodium in salt, the
electropositive H attracts the –ve chlorine
-the salt becomes surrounded by water
molecules and the crystal lattice of salt
is broken up
•ions and molecules that react with water =
•ions and molecules that don’t react =
Solutions, Colloids and
Suspensions
• solution = homogenous (same) mixtures containing a relatively large
amount of one compound (solvent)
– e.g. sugar + water
– the mixture is the same no matter where you sample it
• colloid = solution of larger components called dispersed-phase
particles
– these particles all carry the same charge (repel each other)
– their particles are larger than that of solutions
• e.g. plasma proteins within the blood
• suspension = solution of larger components called dispersed-phase
particles
– larger particles than that of colloid
– if left undisturbed these particles will settle out to form a solid
• e.g. red blood cells within blood
• mixture = two or more types of elements or molecules physically
blended together without the formation of physical bonds between
them
– these individual compounds can be separated by physical or chemical
means
– types of mixtures: combination of solutions, suspensions and colloids
Electrolytes
• substances that release ions when they react with
water
-these ions will conduct electricity = Electrolytes
-created through the decomposition of ionic substances
e.g. salt
-although polar compounds can also liberate electrolytes
Inorganic salts:
• abundant in body fluids
• source of ions eg. Na+, Ca2+, K+, Mg2+
 ions play a role in: maintaining water concentration
maintaining pH
bone development
muscle function
nerve function
 ions must be maintained in a certain concentration
to maintain homeostasis
Inorganic Acids & Bases
•3 types:
1. release H+
e.g. HCl
Acids
H+ + Cl-
2. release ions to combine with H+
e.g. NaOH
Na+ + OH3. acids + bases
e.g. HCl + NaCl
Salts
H20 + NaCl
Bases
•one must consider acids & bases
in light of how they mix with water
e.g. HCl when mixed in water
dissociates into H+ ions and Clions
•if a base such as NaOH is added – it will
dissociate into Na+ ions and OH- ions
- the Na+ ions will combine with the Cl- ions to
form NaCl, the H+ ions will combine with the
OH- ions to reform water.
pH Scale
• measures concentration of H+ ions in a solution
e.g. pH 6 = 1 x 10-6
pH 7 = 1 x 10-7
pH 8 = 1 x 10-8
Buffers:
• chemical or compound that keeps the pH of a solution within
a normal range
• resists pH change by taking up excess H+ or OH- ions
H20
H2CO3
e.g. blood = pH 7.4
“Bicarbonate buffering system”
- our blood contains a small amount of
carbonic acid
H+ +
HCO3-
carbonic acid
excess OH-
excess H+
H2CO3
Organic Compounds
• always contain carbon, oxygen and hydrogen
• carbon can form 4 covalent bonds with other atoms
e.g. methane
H
symmetrical charge
H C H
H
• carbon can also form covalent bonds with itself
forming long chain hydrocarbons
H H H H H H
hydrophobic
H C C C C C C H
(non-polar)
H H H H H H
or a ring structure
H
H
C C
H
C
H
H
H
H
C C
H
H
CH
Organic
compounds
• the skeleton of carbon and hydrogen are frequently combined with other
atoms and molecules = functional groups
H H H
O
H C C C C
O
H
H H
these groups confer a specific property
to the organic compound
e.g. amino acid vs. nucleotide
carboxyl
group = polar group
hydrocarbon + carboxyl group
“hydrophilic”
Organic substances:
1. carbohydrates
2. lipids
3. proteins
4. nucleic acids
1. Carbohydrates:
• provide energy to cells
• supply materials to build certain cell structures
• stored as reserve energy supply (humans = glycogen)
• water soluble
• characterized H - C - OH (ratio C:H 1:2)
e.g. glucose C6H12O6
sucrose C12H22O12
monosaccharides
• classified by size: simple - sugars
disaccharides
complex – polysaccharides
-see Table 2-6 (Tortora)
A. Simple carbohydrates
• monosaccharides = single sugar in which the # of carbon
atoms is low - from 3 to 7
e.g. pentose - 5 carbon sugar
hexose - 6 carbon sugar
 hexose sugars: glucose
galactose
fructose
 pentose sugars: ribose
deoxyribose
-three ways to represent the structure of glucose
1. Molecular form
2.
3. Simplest form
A. Simple carbohydrates
• disaccharide = two 6-carbon monosaccharides
-form by a dehydration synthesis reaction
-broken up by a hydrolysis reaction
e.g. glucose + glucose = maltose
e.g. glucose + fructose = sucrose
e.g. glucose + galactose = lactose
B. Complex carbohydrates:
• built of simple carbohydrates
e.g. glycogen
starch
cellulose
• multiple, repeating monomers or
“building blocks”
polymer
Starch & Glycogen
• starch = storage form of glucose found in plants
-hydrolyzed into glucose
• glycogen = storage form of glucose found in animals
-hydrolyzed into glucose (in liver)
Cellulose
• polysaccharide found in cell walls in plants
• linkage between glucose monomers differs from starch
• indigestible
Lipids
• many types
– 1. triglycerides = fats and oils
– 2. phospholipids
– 3. steroids
•
•
•
•
•
cholesterol – animal cell membranes, basis for steroid hormones
bile salts - digestion
vitamin D – calcium regulation
Adrenocorticosteroid hormones
Sex hormones
– 4. Eicanosoids
• prostaglandins
• leukotrienes
– 5. Others
•
•
•
•
•
fatty acids
carotenes – synthesis of vitamin A
vitamin E – wound healing
vitamin K – blood clotting
lipoproteins – HDL and LDL
2. Lipids
A. Fats
• energy supply
• most plentiful lipids in your body
• composed of C, H and O
• “building blocks” = 3 fatty acid chains (hydrocarbons)
1 glycerol molecule
fatty acid
fatty acid
fatty acid
glycerol portion
fatty acid portion
• fatty acids - carboxyl at end
-differ in chain length with each fat
-differ in carbon bonding
1. single C bonds - saturated
carboxyl gp
2. double C bonds - unsaturated
monounsaturated:
1 double bond
polyunsaturated:
2 or more double bonds
-some fatty acids cannot
be made by the body
and must be taken in
through food = essential
fatty acids
e.g omega-3 fatty acids
-polyunsaturated fatty acids
-important in regulating cholesterol
levels -lower LDL levels in the
blood
-increase calcium utilization by
body – stronger bones & teeth
-reduce inflammation – arthritis
-promote wound healing
B. Phospholipids
• similar to fat molecules - glycerol + 2 fatty acids
+ a phosphate group
• phosphate gp
hydrophilic “head”
• fatty acid gps
hydrophobic “tails”
• form the majority of the cell
membrane = lipid bilayer
C. Steroids
• backbone = 4 fused carbon rings
• diversity through attached functional groups
e.g. cholesterol
testosterone, estrogen
aldosterone
3. Proteins
• roles: structural
energy source
chemical messengers
combine with carbohydrates = glycoproteins
receptors
antibodies
metabolic role - enzymes
• building blocks = amino acids
R
HO C
C
H
N
O
H
H
carboxyl gp
H
amino gp
a.a. = amino group at 1 end,
carboxyl at the other
between is a single C
atom bound to: 1. H atom
2. R group
some amino acids:
asparagine
alanine
arginine
aspartic acid
cysteine
glutamic acid
glycine
histidine
leucine
lysine
phenylalanine
proline
serine
thymine
tyrosine
tryptophan
valine
• amino acids joined together by a condensation reaction
forming a peptide bond = between the NH2 of 1 a.a. and
the COOH of the next
peptide bond = polar
R
R
HO C
C
H
N
C
C
H
N
O
H
H
O
H
H
d+
d-
H
2 a.a.
dipeptide
3 a.a.
tripeptide
4 or more a.a.
polypeptide
• polypeptides have 4 types of structures or conformations
which affect their ultimate function
Protein conformation:
1. primary - a.a. sequence of polypeptides
2. secondary - a.a. chain folds into a-helical coils
or b-pleated sheets
3. tertiary - coiled a.a. helix
folds into a unique
3-D shape
4. quaternary = joining of 2
or more polypeptides
Enzymes
4. Nucleic acids
• make up DNA, RNA
• C,H,O,N,P
• building blocks = nucleotides
• nucleotide: 5 carbon sugar (pentose)
•this pentose sugar has precise numbering of its carbons
phosphate group (negative charge)
organic base - 4 types: adenine (A)
cytosine (C)
guanine (G)
thymine (T)
uracil (U)
• polynucleotide chain - formed by a phosphate bond between
the phosphate (5’) of 1 n.t. and the sugar of the
next (3’)
• two major types of nucleic acids:
1. RNA sugar = ribose
2. DNA sugar = deoxyribose
HOCH2
O
OH
H
H
OH
OH
ribose
HOCH2 O
OH
H
H
OH
H
deoxyribose
• so a DNA/RNA chain “grows” in one direction only
-5’ to 3’
A. RNA
single p.p chain
bases: A, C, G and uracil
(U) in place of T
3 types: mRNA
tRNA
Figure 2.17
rRNA
B. DNA
double p.p chain = double helix
sense strand (5’ to 3’) anti-sense strand
2 chains held by hydrogen bonds
between the bases
bases pair up in a complementary fashion
A=T
C G
c. ATP
individual n.t’s can have metabolic functions
e.g. adenosine = adenine + ribose
-adenine modified by adding three phosphates
major source of ATP = breakdown of glucose
1 glucose molecule
glycolysis
Kreb’s cycle
oxidative phosphorylation
36 ATP