Human Biology Chemical and the body

Dr. jassim Mohammed Abdo

-2-

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

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.

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)

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