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Human Biology
Chemical and the body
Dr. jassim Mohammed Abdo
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2
Human Biology
Chemical and the body

I. Elements:
– Substances that can not be broken down into
simpler substances by chemical reactions.
– There are 92 naturally occurring elements:
Oxygen, carbon, nitrogen, calcium, sodium, etc.
• Life requires about 25 of the 92 elements
• Chemical Symbols:
– Abbreviations for the name of each element.
– Usually one or two letters of the English or Latin
name of the element
– First letter upper case, second letter lower case.
Example: Helium (He), sodium (Na), potassium
(K), gold (Au).

• Main Elements: Over 98% of an organism’s mass is
made up of six elements.
– Oxygen (O): 65% body mass
• Cellular respiration, component of water, and most
organic compounds.
– Carbon (C): 18% of body mass.
• Backbone of all organic compounds.
– Hydrogen (H): 10% of body mass.
• Component of water and most organic compounds.
– Nitrogen (N): 3% of body mass.
• Component of proteins and nucleic acids (DNA/RNA)
– Calcium (Ca): 1.5% of body mass.
• Bones, teeth, clotting, muscle and nerve function.
– Phosphorus (P): 1% of body mass
• Bones, nucleic acids, energy transfer (ATP).

• Minor Elements: Found in low amounts. Between
1% and 0.01%.
– Potassium (K): Main positive ion inside cells.
• Nerve and muscle function.
– Sulfur (S): Component of most proteins.
– Sodium (Na): Main positive ion outside cells.
• Fluid balance, nerve function.
– Chlorine (Cl): Main negative ion outside cells.
• Fluid balance.
– Magnesium (Mg): Component of many
enzymes and chlorophyll.

• Trace elements: Less than 0.01% of mass:
– Boron (B)
– Chromium (Cr)
– Cobalt (Co)
– Copper (Cu)
– Iron (Fe)
– Fluorine (F)
– Iodine (I)
– Manganese (Mn)
– Molybdenum (Mo)
– Selenium (Se)
– Silicon (Si)
– Tin (Sn)
– Vanadium (V)
– Zinc (Zn)

II. Structure & Properties of Atoms
Atoms: Smallest particle of an element that
retains its chemical properties. Made up of three
main subatomic particles.
Particle Location Mass Charge
Proton (p+) In nucleus 1 +1
Neutron (no) In nucleus 1 0
Electron (e-) Outside nucleus 0 -1

Structure and Properties of Atoms
1. Atomic number = # protons
– The number of protons is unique for each element
– Each element has a fixed number of protons in its
nucleus. This number will never change for a given
element.
– Written as a subscript to left of element symbol.
Examples: 6C, 8O, 16S, 20Ca
– Because atoms are electrically neutral (no charge),
the number of electrons and protons are always the
same.
– In the periodic table elements are organized by
increasing atomic number.

Structure and Properties of Atoms:
2. Mass number = # protons + # neutrons
– Gives the mass of a specific atom.
– Written as a superscript to the left of the element
symbol.
Examples: 12C, 16O, 32S, 40Ca.
– The number of protons for an element is always the
same, but the number of neutrons may vary.
– The number of neutrons can be determined by:
# neutrons = Mass number - Atomic number

Structure and Properties of Atoms:
3. Isotopes: Variant forms of the same element.
– Isotopes have different numbers of neutrons
and therefore different masses.
– Isotopes have the same numbers of protons and
electrons.
– Example: In nature there are three forms or
isotopes of carbon (6C):
• 12C: About 99% of atoms. Have 6 p+, 6 no, and 6 e-.
• 13C: About 1% of atoms. Have 6 p+, 7 no, and 6 e-.
• 14C: Found in tiny quantities. Have 6 p+, 8 no, and 6 e-.
Radioactive form (unstable). Used for dating
fossils.

Electron Arrangements of Important Elements of Life
1 Valence electron 4 Valence electrons 5 Valence electrons 6 Valence electrons

III. How Atoms Form Molecules: Chemical
Bonds
Molecule: Two or more atoms combined chemically.
Compound: A substance with two or more elements
combined in a fixed ratio.
• Water (H2O)
• Hydrogen peroxide (H2O2)
• Carbon dioxide (CO2)
• Carbon monoxide (CO)
• Table salt (NaCl)
– Atoms are linked by chemical bonds.
Chemical Formula: Describes the chemical composition of
a molecule of a compound.
– Symbols indicate the type of atoms
– Subscripts indicate the number of atoms

How Atoms Form Molecules:
Chemical Bonds
Atoms can lose, gain, or share electrons to satisfy
octet rule (fill outermost shell).
Two main types of Chemical Bonds
A. Ionic bond: Atoms gain or lose electrons
B. Covalent bond: Atoms share electrons

A. Ionic Bond: Atoms gain or lose electrons.
Bonds are attractions between ions of opposite
charge.
Ionic compound: One consisting of ionic bonds.
Na + Cl ----------> Na+ Cl-
sodium chlorine Table salt
(Sodium chloride)
Two Types of Ions:
Anions: Negatively charged particle (Cl-)
Cations: Positively charged particle (Na+)

B. Covalent Bond: Involves the “sharing” of one
or more pairs of electrons between atoms.
Covalent compound: One consisting of
covalent bonds.
Example: Methane (CH4): Main component
of natural gas.
H
|
H---C---H
|
H
Each line represents on shared pair of electrons.
Octet rule is satisfied: Carbon has 8 electrons,
Hydrogen has 2 electrons

Electronegativity: A measure of an atom’s
ability to attract and hold onto a shared
pair of electrons.
Some atoms such as oxygen or nitrogen
have a much higher electronegativity than
others, such as carbon and hydrogen.
Two Types of Covalent Bonds: Polar and Nonpolar

Polar and Nonpolar Covalent Bonds
A. Nonpolar Covalent Bond: When the
atoms in a bond have equal or similar
attraction for the electrons
(electronegativity), they are shared equally.
Example: O2, H2, Cl2

Nonpolar Covalent Bonds:
Electrons are Shared Equally

B. Polar Covalent Bond: When the atoms in a
bond have different electronegativities, the
electrons are shared unequally.
Electrons are closer to the more
electronegative atom creating a polarity or
partial charge.
Example: H2O
Oxygen has a partial negative charge.
Hydrogens have partial positive charges.
Polar and Nonpolar Covalent Bonds

Other Bonds: Weak chemical bonds are important in the
chemistry of living things.
• Hydrogen bonds: Attraction between the partially positive H
of one molecule and a partially negative atom of another
– Hydrogen bonds are about 20 X easier to break than a
normal covalent bond.
– Responsible for many properties of water.
– Determine 3 dimensional shape of DNA and proteins.
– Chemical signaling (molecule to receptor).

– Living cells are 70-90% water
– Water covers 3/4 of earth’s surface
– Water is the ideal solvent for chemical
reactions
– On earth, water exists as gas, liquid,
and solid
Water: The Ideal Compound for Life

I. Polarity of water causes hydrogen bonding
– Water molecules are held together by H-
bonding
– Partially positive H attracted to partially
negative O atom.
• Individual H bond are weak, but the cumulative
effect of many H bonds is very strong.
• H bonds only last a fraction of a second, but at any
moment most molecules are hydrogen bonded to
others.

Unique properties of water caused by H-bonds
– Cohesion: Water molecules stick to each other.
This causes surface tension.
– Adhesion: Water sticks to many surfaces.
Capillary Action: Water tends to rise in narrow
tubes.

Unique properties of water caused by H-bonds
– Universal Solvent: Dissolves many (but not all) substances to
form solutions.
Solutions are homogeneous mixtures of two or more
substances (salt water, air, tap water).
All solutions have at least two components:
• Solvent: Dissolving substance (water, alcohol, oil).
– Aqueous solution: If solvent is water.
• Solute: Substance that is dissolved (salt, sugar, CO2).
– Water dissolves polar and ionic solutes well.
– Water does not dissolve nonpolar solvents well.

Solubility of a Solute Depends on its
Chemical Nature
Solubility: Ability of substance to dissolve in a given
solvent.
Two Types of Solutes:
A. Hydrophilic: “Water loving” dissolve easily in
water.
• Ionic compounds (e.g. salts)
• Polar compounds (molecules with polar regions)
• Examples: Compounds with -OH groups (alcohols).
• “Like dissolves in like”

Solubility of a Solute Depends on its
Chemical Nature
Two Types of Solutes:
B. Hydrophobic: “Water fearing” do not
dissolve in water
• Non-polar compounds (lack polar regions)
• Examples: Hydrocarbons with only C-H non-polar
bonds, oils, gasoline, waxes, fats, etc.

ACIDS, BASES, pH AND BUFFERS
A. Acid: A substance that donates protons (H+).
– Separate into one or more protons and an anion:
HCl (into H2O ) -------> H+ + Cl-
H2SO4 (into H2O ) --------> H+ + HSO4
-
– Acids INCREASE the relative [H+] of a solution.
– Water can also dissociate into ions, at low levels:
H2O <======> H+ + OH-

B. Base: A substance that accepts protons (H+).
– Many bases separate into one or more positive ions
(cations) and a hydroxyl group (OH- ).
– Bases DECREASE the relative [H+] of a solution ( and
increases the relative [OH-] ).
H2O <======> H+ + OH-
Directly NH3 + H+ <=------> NH4
+
Indirectly NaOH ---------> Na+ + OH-
( H+ + OH- <=====> H2O )

Strong acids and bases: Dissociation is almost complete
(99% or more of molecules).
HCl (aq) -------------> H+ + Cl-
NaOH (aq) -----------> Na+ + OH-
(L.T. 1% in this form) (G.T. 99% in dissociated form)
• A relatively small amount of a strong acid or base will
drastically affect the pH of solution.
Weak acids and bases: A small percentage of molecules
dissociate at a give time (1% or less)
H2CO3 <=====> H+ + HCO3
-
carbonic acid Bicarbonate ion
(G.T. 99% in this form) (L.T. 1% in dissociated form)

C. pH scale: [H+] and [OH-]
– pH scale is used to measure how basic or acidic a solution
is.
– Range of pH scale: 0 through 14.
• Neutral solution: pH is 7. [H+ ] = [OH-]
• Acidic solution: pH is less than 7. [H+ ] > [OH-]
• Basic solution: pH is greater than 7. [H+ ] < [OH-]
– As [H+] increases pH decreases (inversely proportional).
– Logarithmic scale: Each unit on the pH scale represents a
ten-fold change in [H+].

D. Buffers keep pH of solutions relatively constant
– Buffer: Substance which prevents sudden large changes
in pH when acids or bases are added.
– Buffers are biologically important because most of the
chemical reactions required for life can only take place
within narrow pH ranges.
– Example:
• Normal blood pH 7.35-7.45. Serious health problems will
arise if blood pH is not stable.

CHEMICAL REACTIONS
– A chemical change in which substances (reactants) are
joined, broken down, or rearranged to form new
substances (products).
– Involve the making and/or breaking of chemical bonds.
– Chemical equations are used to represent chemical
reactions.
Example:
2 H2 + O2 -----------> 2H2O
2 Hydrogen Oxygen 2 Water
Molecules Molecule Molecules

Organic Chemistry:
Carbon Based Compounds
A. Inorganic Compounds: Compounds without carbon.
B. Organic Compounds: Compounds synthesized by cells and containing
carbon (except for CO and CO2).
– Diverse group: Several million organic compounds are known
and more are identified every day.
– Common: After water, organic compounds are the most
common substances in cells.
• Over 98% of the dry weight of living cells is made up of organic
compounds.
• Less than 2% of the dry weight of living cells is made up of inorganic
compounds.

Carbon: unique element for basic building
block of molecules of life
• Carbon has 4 valence electrons: Can form
four covalent bonds
– Can form single , double, triple bonds.
– Can form large, complex, branching
molecules and rings.
– Carbon atoms easily bond to C, N, O, H, P,
S.
• Huge variety of molecules can be formed
based on simple bonding rules of basic
chemistry

Diversity of Organic Compounds
• Hydrocarbons:
– Organic molecules that contain C and H only.
– Good fuels, but not biologically important.
– Undergo combustion (burn in presence of oxygen).
– In general they are chemically stable.
– Nonpolar: Do not dissolve in water (Hydrophobic).
Examples:
• (1C) Methane: CH4 (Natural gas).
• (2C) Ethane: CH3CH3
• (3C) Propane: CH3CH2CH3 (Gas grills).
• (4C) Butane: CH3CH2CH2CH3 (Lighters).

Relatively few monomers are used by cells to make
a huge variety of macromolecules
Macromolecule Monomers or Subunits
1. Carbohydrates 20-30 monosaccharides
or simple sugars
2. Proteins 20 amino acids
3. Nucleic acids (DNA/RNA) 4 nucleotides
(A,G,C,T/U)
4. Lipids (fats and oils) ~ 20 different fatty acids
and glycerol.

III. Carbohydrates: Molecules that store energy and are used
as building materials
– General Formula: (CH2O)n
– Simple sugars and their polymers.
– Diverse group includes sugars, starches, cellulose.
– Biological Functions:
– Fuels, energy storage
– Structural component (cell walls)
– DNA/RNA component
– Three types of carbohydrates:
A. Monosaccharides
B. Disaccharides
C. Polysaccharides

A. Monosaccharides: “Mono” single & “sacchar” sugar
– Preferred source of chemical energy for cells (glucose)
– Can be synthesized by plants from light, H2O and CO2.
– Store energy in chemical bonds.
– Carbon skeletons used to synthesize other molecules.
Characteristics:
1. May have 3-8 carbons. -OH on each carbon; one with C=0
2. Names end in -ose. Based on number of carbons:
• 5 carbon sugar: pentose
• 6 carbon sugar: hexose.
3. Can exist in linear or ring forms
4. Isomers: Many molecules with the same molecular
formula, but different atomic arrangement.
• Example: Glucose and fructose are both C6H12O6.
Fructose is sweeter than glucose.

B. Disaccharides: “Di” double & “sacchar” sugar
 Covalent bond formed by condensation reaction between 2
monosaccharides.
Examples:
1. Maltose: Glucose + Glucose.
• Energy storage in seeds.
• Used to make beer.
2. Lactose: Glucose + Galactose.
• Found in milk.
• Lactose intolerance is common among adults.
• May cause gas, cramping, bloating, diarrhea, etc.
3. Sucrose: Glucose + Fructose.
• Most common disaccharide (table sugar).
• Found in plant sap.

C. Polysaccharides: “Poly” many (8 to 1000)
Functions: Storage of chemical energy and structure.
– Storage polysaccharides: Cells can store simple sugars in
polysacharides and hydrolyze them when needed.
1. Starch: Glucose polymer (Helical)
• Form of glucose storage in plants (amylose)
• Stored in plant cell organelles called plastids
2. Glycogen: Glucose polymer (Branched)
• Form of glucose storage in animals (muscle and liver cells)

– Structural Polysaccharides: Used as structural
components of cells and tissues.
1. Cellulose: Glucose polymer.
• The major component of plant cell walls.
• CANNOT be digested by animal enzymes.
• Only microbes have enzymes to hydrolyze.
2. Chitin: Polymer of an amino sugar (with NH2 group)
• Forms exoskeleton of arthropods (insects)
• Found in cell walls of some fungi

Lipids: Fats, phospholipids, and steroids
Diverse groups of compounds.
Composition of Lipids:
– C, H, and small amounts of O.
Functions of Lipids:
– Biological fuels
– Energy storage
– Insulation
– Structural components of cell membranes
– Hormones

1. Simple Lipids: Contain C, H, and O only.
A. Fats (Triglycerides).
• Glycerol : Three carbon molecule with three hydroxyls.
• Fatty Acids: Carboxyl group and long hydrocarbon
chains.
– Characteristics of fats:
• Most abundant lipids in living organisms.
• Hydrophobic (insoluble in water) because nonpolar.
• Economical form of energy storage (provide 2X the
energy/weight than carbohydrates).
• Greasy or oily appearance.

Types of Fats
– Saturated fats: Hydrocarbons saturated with H. Lack -
C=C- double bonds.
• Solid at room temp (butter, animal fat, lard)
– Unsaturated fats: Contain -C=C- double bonds.
• Usually liquid at room temp (corn, peanut, olive oils)

2. Complex Lipids: In addition to C, H, and O, also contain
other elements, such as phosphorus, nitrogen, and sulfur.
A. Phospholipids: Are composed of:
• Glycerol
• 2 fatty acid
• Phosphate group
– Amphipathic Molecule
• Hydrophobic fatty acid “tails”.
• Hydrophilic phosphate “head”.
Function: Primary component of the plasma membrane
of cells

B. Steroids: Lipids with four fused carbon rings
Includes cholesterol, bile salts, reproductive, and adrenal
hormones.
• Cholesterol: The basic steroid found in animals
– Common component of animal cell membranes.
– Precursor to make sex hormones (estrogen, testosterone)
– Generally only soluble in other fats (not in water)
– Too much increases chance of atherosclerosis.
C. Waxes: One fatty acid linked to an alcohol.
• Very hydrophobic.
• Found in cell walls of certain bacteria, plant and insect
coats. Help prevent water loss.

Proteins: Large three-dimensional
macromolecules responsible for most cellular
functions
– Polypeptide chains: Polymers of amino acids
linked by peptide bonds in a SPECIFIC linear
sequence
– Protein: Macromolecule composed of one or
more polypeptide chains folded into SPECIFIC
3-D conformations

Polypeptide: Polymer of amino acids connected in a
specific sequence
A. Amino acid: The monomer of
polypeptides
• Central carbon
– H atom
– Carboxyl group
– Amino group
– Variable R-group

Protein Function is dependent upon Protein Structure (Conformation)
CONFORMATION: The 3-D shape of a protein is determined by
its amino acid sequence.
Four Levels of Protein Structure
1. Primary structure: Linear amino acid sequence,
determined by gene for that protein.
2. Secondary structure: Regular coiling/folding of
polypeptide.
• Alpha helix or beta sheet.
• Caused by H-bonds between amino acids.

3. Tertiary structure: Overall 3-D shape of a polypeptide
chain.
4. Quaternary structure: Only in proteins with 2 or more
polypeptides. Overall 3-D shape of all chains.
• Example: Hemoglobin (2 alpha and 2 beta
polypeptides)

Nucleic acids store and transmit hereditary information for all living things
There are two types of nucleic acids in living things:
A. Deoxyribonucleic Acid (DNA)
• Contains genetic information of all living organisms.
• Has segments called genes which provide information to make
each and every protein in a cell
• Double-stranded molecule which replicates each time a cell
divides.
B. Ribonucleic Acid (RNA)
• Three main types called mRNA, tRNA, rRNA
• RNA molecules are copied from DNA and used to make gene
products (proteins).
• Usually exists in single-stranded form.

DNA and RNA are polymers of nucleotides that determine the primary
structure of proteins
• Nucleotide: Subunits of DNA or RNA.
Nucleotides have three components:
1. Pentose sugar (ribose or deoxyribose)
2. Phosphate group to link nucleotides (-PO4)
3. Nitrogenous base (A,G,C,T or U)
• Purines: Have 2 rings.
Adenine (A) and guanine (G)
• Pyrimidines: Have one ring.
Cytosine (C), thymine (T) in DNA or uracil (U) in RNA.

James Watson and Francis Crick Determined the 3-D Shape of DNA in
1953
– Double helix: The DNA molecule is a double helix.
– Antiparallel: The two DNA strands run in opposite directions.
• Strand 1: 5’ to 3’ direction (------------>)
• Strand 2: 3’ to 5’ direction (<------------)
– Complementary Base Pairing: A & T (U) and G & C.
• A on one strand hydrogen bonds to T (or U in RNA).
• G on one strand hydrogen bonds to C.
– Replication: The double-stranded DNA molecule can easily
replicate based on A=T and G=C pairing.---
– SEQUENCE of nucleotides in a DNA molecule dictate the amino
acid SEQUENCE of polypeptides

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