Physiology 31 Lecture
Chapter 2 - Atoms, Ions & Molecules
I. Overview
A. Elements & Atoms
B. Molecules & Bonds
C. Solutions & Solutes
D. Acids, Bases & pH
E. Biomolecules
II. Elements & Atoms
A. Matter is anything that takes up space and has mass; all matter is
composed of elements
B. An element is a substance that contains just one type of atom
1. The most abundant elements (98% of wt.) in the body are:
a. Carbon
(C)
b. Hydrogen
(H)
c. Oxygen
(O)
d. Nitrogen
(N)
e. Calcium
(Ca)
f. Phosphorous
(P)
2. Lesser elements (minerals) make up about 0.8% of body wt.:
a. Sodium
(Na)
b. Potassium
(K)
c. Chlorine
(Cl)
d. Sulfur
(S)
e. Iron
(Fe)
f. Magnesium
(Mg)
3. Trace elements are only about 0.7%, but are also essential
a. They include chromium (Cr), Cobalt (Co), Copper (Cu),
Fluorine (F), Iodine (I), Lithium (Li), Manganese (Mn),
Molybdenum (Mo), Selenium (Se), Silicon (Si), Tin (Sn),
Vanadium (V), and Zinc (Z)
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b. Trace elements often serve as cofactors in enzymes
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C. Periodic Table of Elements Organization
1. Horizontal rows are arranged by atomic number (number of
protons) & mass (protons + neutrons)
2. Vertical columns are arranged by no. of electron orbitals
D. Atom - Smallest unit of matter involved in chemical reactions. An
atom consists of:
1. Nucleus - atom center composed of protons & neutrons
a. Proton (p+) has positive (+) charge; atomic mass of 1 amu
b. Neutron (no) has no charge; atomic mass of 1 amu
c. Atomic Number = number of protons
d. Atomic mass = number of protons + neutrons
2. Electron (e-) has a negative (-) charge and negligible weight
a. Electron Orbitals (shells)
1) The innermost orbital contains up to 2 electrons
2) Each outer orbital contains up to 8 electrons
3) Atoms are most stable when they have full electron shells
b. Highest energy electrons (valence electrons) are located in
orbitals farthest from the nucleus, and react with other atoms
c. An electrically neutral atom has # electrons = # protons
E. Isotope - atom with same # of protons and electrons as another
atom of the same type, but with a different # of neutrons (e.g., 14C
vs. 12C)
1. Radioactive isotopes break down (decay) and emit radioactive
particles
2. Radioisotopes are used in medicine to diagnose and treat some
diseases
F. Ions are atoms that have gained or lost electrons, giving them and
overall negative or positive charge
1. Anions gain electrons, thus have a negative charge (e.g., Cl-)
2. Cations lose electrons, thus have a positive charge (e.g., Na+)
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G. Electrolytes are molecules that ionize in water to form solutions
that can conduct electricity.
1. Human electrolytes include: Na+, Cl-, K+, Ca2+, Mg2+, PO43-, and
HCO32. Electrolytes are essential to nerve and muscle function
H. A Free Radical is a particle with an unpaired electron (e.g., O2-.)
1. Free radicals disrupt other molecules in the body by taking
electrons from them
2. Antioxidants, such as vitamins A, C, and E, neutralize free
radicals
III. Molecules & Bonds
A. Molecule - Two or more atoms covalently bonded together
(e.g., O2, N2)
B. Compound – molecule that contains more than one type of
element (e.g., CO2, H2O).
C. Four types of bonds join atoms: ionic, covalent, and hydrogen
bonds, and Van der Waals forces
D. Ionic Bond - strong attraction between negatively and positively
charged ions
1. In ionic reactions, atoms give or take electrons to get a full outer
electron orbital
2. Oppositely charged ions are strongly attracted to each other,
form ionic bonds, and are called salts or electrolytes
E. Covalent Bonds are strong chemical bonds between atoms that
result from the sharing of electrons in their outer orbitals.
Molecules with covalent bonds are represented 3 ways:
1. Electron-dot structures (e.g.: :O::C::O:) in which each dot
represents an electron in an atom’s outer orbital
2. Structural formulas in which each pair of shared electrons is
represented by a line (e.g.: O=C=O).
3. Chemical formulas that show only the number of each type of
atom in the molecule (e.g.: CO2)
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4. Chemical formulas can be used to calculate molecular weight
by multiplying the atomic mass of each element in the molecule
by the number of atoms in the molecule, then adding the results.
5. Nonpolar molecules are those whose atoms share their electrons
equally, thus do not have an overall charge (e.g., CO2)
6. Polar molecules are those in which some atoms in the molecule
attract the shared electrons more strongly, thus the molecule has
partial positive and negative charges (e.g., H2O)
E. Hydrogen Bonds are weak bonds between a slightly positive
hydrogen atom and a slightly negative oxygen (e.g., in H2O) or
nitrogen atom (e.g., in DNA)
F. Van der Waals forces are very weak bonds that result from the
attraction of one atom’s nucleus to the electrons of another atom
IV. Solutions & Solutes
A. A solution consists of minute particles (solute) dissolved in a
solvent (usually water). The solute may be gas, liquid, or solid.
B. The body is composed of about 60% water, which has important
functions
1. Polar (hydrophilic) molecules and salts dissolve in water
a. The degree to which a molecule dissolves is its solubility
b. The more readily a molecule dissolves, the higher its solubility
2. Nonpolar (hydrophobic) molecules (e.g., lipids) do not dissolve
in water
3. Water’s surface tension is important in lung function
4. Water has a high heat capacity (amount of heat needed to raise
the temperature of 1 gram of water by 1C = 1cal/g). Thus, body
temp. rises and falls slowly, preventing sudden changes
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C. Concentrations of solutions can be expressed several ways
1. Moles & Molarity
a. Mole – a large number of particles equal to 6.02 x 1023
b. Gram molecular weight is the atomic mass or molecular
weight of a substance in grams. Example: the molecular wt. of
glucose (C6H12O6) is 180 amu, so one mole glucose = 180 g
c. Molarity is the number of moles of solute in a liter of solution
(mol/L or M).
1) A one molar solution of glucose would be made by
dissolving 180 g (one mole) of glucose in enough water to
make 1 liter total
2) Biological solutions are so dilute, they are usually expressed
in millimoles per liter (mmol/L or mM)
2. Equivalent (eq) – the molarity of an ion times the number of
charges it carries.
a. Examples: Na+ would have 1 equivalent per mole; Ca2+ would
have 2 equivalents per mole
b. Concentration of ions in the blood are often expressed in
milliequivalents per liter (meq/L)
3. Weight per Volume and Percent Solutions
a. Percent solution - the percent of solute in the total solution.
b. If the solute is solid, the percent solution is expressed as
weight/volume. Example: a 5% solution of glucose is made
by mixing 5 g glucose with enough water to make 100 ml
c. If the solute is liquid, the solution is made using
volume/volume measurements. Example: to make 0.1% HCl
sol’n, mix 0.1 ml concentrated HCl with water to make 100 ml
d. Concentrations of drugs and chemicals in the body are often
given in milligrams per deciliter (mg/dl). A deciliter is 100 ml.
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V. Acids, Bases, & pH
A. Liquid water dissociates into equal numbers of H+ & -OH ions, in
addition to a majority of H2O molecules (e.g.: H2O  H+ + -OH)
B. Acids - dissociate in water and release H+ ions (protons) (e.g.:
HCl  H+ + Cl-); the more H+ ions released, the stronger the acid
C. Bases - dissociate in water and take up H+ ions or release -OH
(hydroxide) ions (e.g.: NaOH  Na+ + -OH)
D. The pH Scale is used to measure acidity and basicity
1. pH number indicates degree of acidity or basicity of solution
2. pH = power of hydrogen - the negative logarithm of the
hydrogen ion concentration of a solution (pH 3 = -log 3 [H+] =
10-3 M H+). pH is an exponential scale of H+ concentration
3. Acid is less than pH 7; the lower the number, the more H+, the
more acidic. A decrease of 1 pH increases acidity 10 fold
4. Neutral is pH 7. In pure water (pH7) H+ = -OH
5. Base is more than pH 7; the higher the number, the less H+, the
more basic. An increase of 1 pH increases basicity 10 fold
E. Blood pH must be maintained within 7.35-7.45 pH.
1. H+ ion concentrations lower or higher than this range can
interfere with hydrogen bonding and Van der Waals forces
holding organic molecules together
F. Buffers - chemicals that take up excess H+ or -OH ions and keep
the pH constant
1. Example: carbon dioxide and water in the blood combine to
form weak carbonic acid, which dissociates into hydrogen ions
and bicarbonate ions. The bicarbonate is a buffer that can take
up excess H+ ions. (CO2 + H2O ↔ H2CO3  H+ + HCO3-)
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VII. Biomolecules are organic molecules that compose living things
A. Organic molecules always contain carbon and hydrogen
1. Carbon has 4 electrons in its outer shell, therefore can form
up to 4 covalent bonds with other atoms (i.e., H, O, N, S)
2. This property allows carbon to become the backbone of
organic molecules
3. Functional groups, such as carboxyl (-COOH = acid) and
hydroxyl (-OH = base), methyl (-CH3), amine (-NH2), and
phosphate (-H2PO4) can be bonded to carbon groups
B. Four classes of organic polymers are: carbohydrates, lipids,
proteins, and nucleic acids
VIII. Carbohydrates – hydrophilic molecules that contain C, H, and O
in a ratio of about 1-C:2-H:1-O; function: primary energy source &
structural role in cells.
A. Monosaccharides - simple sugars with 3-7 carbon atoms in
[CH2O]n form; the basic building blocks of carbohydrates.
Pentose sugars have 5 carbons, hexose sugars have 6 carbons.
1. Glucose (C6H12O6) - primary energy source for cells (blood
sugar)
2. Fructose (C6H12O6) - fruit sugar
3. Galactose (C6H12O6) - milk sugar
B. Disaccharides - two monosaccharides bond together
1. Sucrose - glucose + fructose
2. Maltose - glucose + glucose
3. Lactose - glucose + galactose
C. Polysaccharides - polymers of many monosaccharides (usually
glucose molecules) covalently bonded together
1. Glycogen - storage form of glucose in animals; stored in our
liver and muscle cells, broken down to glucose when needed
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2. Starch - storage form of glucose in plants; stored in starch
granules in plant cells, digested to glucose in our bodies
3. Cellulose - main component of plant cell walls; also known as
fiber; we can’t digest this because the covalent bonds between
the glucose molecule are slightly different than those of starch
and glycogen
4. Most carbohydrates are broken down to glucose, which is used
in cellular respiration to produce ATP energy
IX. Lipids - molecules composed of CHO in a variable ratio; functions:
long-term energy storage, adipose tissue insulates and cushions
organs. Lipids are nonpolar, do not mix with water. Types of lipids
include:
A. Fatty acid - hydrocarbon chain with carboxylic acid group at
one end and a methyl group at the other end
1. Saturated - have no double bonds between carbons;
unhealthy fats, mostly from animal sources and
hydrogenated plant oils; solid at room temp.
2. Unsaturated - have one or more double bonds between
carbons; healthier oils, from plant sources, liquid at room
temp.
B. Triglyceride - three fatty acids bonded to one glycerol; nonpolar
molecule; excess in blood contributes to heart disease
C. Phospholipid - two fatty acids (nonpolar) and a phosphate group
(polar) bonded to one glycerol; found in cell membranes
D. Steroids - have a backbone of 4 fused carbon rings with different
functional groups; includes cholesterol and several hormones
1. Cholesterol - manufactured in animal livers, used in cell
membranes, bile, and steroid hormones. Circulates in the blood
in two main forms:
a. LDL (Low Density Lipoprotein) - “bad” cholesterol, excess
in blood clogs arteries
b. HDL (High Density Lipoprotein) - “good” cholesterol,
recycles cholesterol back to liver
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2. Steroid hormones include the sex hormones estrogen,
progesterone, and testosterone
E. Eicosanoids are 20 carbon compounds derived from the fatty acid
arachadonic acid
1. Include leukotrienes, thromboxanes, and prostaglandins
2. These are signaling molecules in inflammation, blood clotting,
and control of blood vessel diameter
X. Proteins - large organic molecules consisting of chains of 50 or more
amino acids composed of carbon, hydrogen, oxygen, nitrogen, and
sometimes sulfur.
A. Structural proteins - give strength and shape to different tissues
(e.g.: keratin in hair & nails, collagen in connective tissues, muscle
proteins, etc.)
B. Functional proteins - such as enzymes, hormones, membrane
proteins, antibodies, etc., are proteins that facilitate metabolic
reactions.
C. Amino Acids - basic building blocks of proteins; there are 20
different amino acids, each of which has the following bonded to a
central carbon atom (see book for the 20 amino acid names)
1. Amino group (-NH2)
2. Carboxyl group (-COOH)
3. Variable “R” (radical) group; some “R” groups are polar
(hydrophilic), and some are nonpolar (hydrophobic)
D. Amino acids may be essential or nonessential
1. Nonessential AAs are the 11 that are synthesized by the body
2. Essential AAs are the 9 that cannot be synthesized by the body,
thus must be obtained through the diet
E. Peptide – two to nine amino acids linked together by peptide
bonds, formed between the amine group of one amino acid and the
carboxyl group of another (HOOC-NH2)
1. Polypeptide – 10-100 bonded amino acids
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2. Protein – polypeptide of more than 100 amino acids
F. Levels of Protein Organization
1. Primary Structure - linear sequence of amino acids bonded
together. All polypeptides have specific amino acid sequences
2. Secondary Structure - twisting or folding of the polypeptide
chain (e.g.: alpha helix or beta-pleated sheet), formed by
hydrogen bonding between negative C=O and positive N-H
3. Tertiary Structure – folding and coiling of a polypeptide (e.g.:
to fibrous or globular) due to interactions among R groups and
surrounding water
4. Quaternary Structure – association of 2 or more polypeptide
chains together (e.g.: hemoglobin has 4 polypeptide chains)
5. Proteins must fold into precise shapes to function properly
6. Proteins can be denatured (nonfunctional change in shape) by
heat or strong pH
G. Proteins, lipids, and carbohydrates can combine to form
glycoproteins, glycolipids, or lipoproteins
XI. Nucleic Acids - include DNA, RNA, and ATP; functions include
cellular reproduction (DNA), protein manufacturing instructions
(DNA & RNA), and cellular energy source (ATP)
A. Nucleotide - basic building block of nucleic acids; composed of a
phosphate group, a pentose sugar, and a nitrogenous base
B. DNA (Deoxyribonucleic Acid)
1. Found in chromosomes, mostly in the cell nucleus
2. Forms genes (specific sequences of DNA) that regulate the
growth & reproduction of cells and direct the order in which
amino acids are put together to form proteins
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3. DNA is a double-stranded helix (like a twisted ladder) held
together by hydrogen bonds and is composed of:
a. Deoxyribose sugar and phosphate groups form the side rails
of the “ladder”
b. Nitrogenous bases held together by hydrogen bonds form
the ladder steps. These bases include:
1) Thymine (T)
2) Adenine (A)
3) Cytosine (C)
4) Guanine (G)
C. RNA (Ribonucleic Acid)
1. Works with DNA for protein synthesis
2. Is mostly a single-stranded molecule
3. Has ribose sugar
4. Has the same nitrogenous bases as DNA, except it has uracil
instead of thymine
D. ATP (Adenosine Triphosphate)
1. Is a nucleotide composed of adenine, a ribose sugar, and 3
phosphate groups held together with high energy bonds
2. ATP is formed via the process of cellular respiration, during
which glucose bond energy is transferred to ATP bonds
3. Functions as the main energy source in cells due to energy in
the phosphate bonds
4. Other molecules related to ATP include cAMP, NAD+, and
FADH
a. cAMP is a cell signaling molecule
b. NAD+ & FADH are energy transfer molecules in cell
respiration