1. The document discusses the key chemical elements that make up the human body, including oxygen, carbon, hydrogen, nitrogen, calcium, and phosphorus. 2. It also explains the structure of atoms and how they form bonds through ionic and covalent interactions. Water is highlighted as being an ideal solvent for life due to its unique properties arising from hydrogen bonding. 3. The roles of acids, bases, and pH in solutions are covered. Buffers are described as substances that help maintain a stable pH when acids or bases are introduced.

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

2. 2
Human Biology
Chemical and the body

3. 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).

4. • 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).

5. • 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.

6. • 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)

7. 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

10. 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.

11. 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

12. 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.

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

14. 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

15. 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

16. 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+)

19. 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

21. 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

22. 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

23. Nonpolar Covalent Bonds:
Electrons are Shared Equally

24. 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

25. 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).

26. – 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

27. 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.

28. 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.

29. 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.

30. 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”

31. 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.

32. 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-

33. 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 )

34. 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)

35. 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+].

36. 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.

37. 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

46. 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.

47. 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

49. 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).

50. 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.

51. 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

52. 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.

54. 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.

56. 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)

58. – 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

59. 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

60. Lipids: Fats, phospholipids, and steroids
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.

61. Lipids: Fats, phospholipids, and steroids
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)

62. 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

64. 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.

65. 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

66. 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

67. 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.

68. 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)

71. 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.

72. 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.

73. 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