Introductory Biochemistry

Instructors

• Dr. Nafez Abu Tarboush

• Dr. Mamoun Ahram

Recommended textbooks

• Biochemistry; Mary K. Campbell and Shawn O. Farrell, Brooks Cole; 6th edition

Recommended electronic web address

• NCBI Bookshelf:

(http://www.ncbi.nlm.nih.gov/entrez/query.fcgi?db=Books)

• The Medical Biochemistry Page: (http://web.indstate.edu/thcme/mwking/home.html)

• Biochemistry, Garret and Grishan, Second Ed.: http://web.virginia.edu/Heidi/home.htm

Outline

• Introduction (MA) • Acid, base, and pH (MA) • Macromolecules and carbohydrates (MA) • Lipids (MA) • Nucleic acids and nucleotides (MA) • Amino acids (NA) • Polypeptides and protein structure (NA) • Protein analysis (NA) • Protein structure-function relationship (NA) • Enzymes (NA) • Enzymes (cofactors) (MA)

Mid-term (preliminary)

Sunday, July 1

Introduction into biochemistry

&

Chemical composition of living organisms

Dr. Mamoun Ahram

Lecture 1

Reference

• Campbell and Farrell, Page 35-43

What is biochemistry?

• Biochemistry is the chemistry of living organisms

• It seeks to describe the structure, organization, and functions of living matter in molecular terms

Understanding life

• Know the chemical structures of biological molecules

• Understand the biological function of these molecules

• Understand interaction and organization of different molecules within individual cells and whole biological systems

• Understand bioenergetics (the study of energy flow in cells)

Biochemistry and medicine

• diagnose and monitor diseases

• design drugs (new antibiotics, chemotherapy agents)

• understand the molecular bases of diseases

The chemical elements

Chemical elements in living creatures

• Living organisms on Earth are composed mainly of 31 elements

Abundant elements

• Four primary elements: carbon, hydrogen, oxygen, and nitrogen

– 96.5% of an organism's weight

• The second groups includes sulfur and phosphorus

• Most biological compounds are made of only SIX elements: C, H, O, N, P, S

Others…

• Minor, but essential, elements

– Mostly metals

Dalton

• The atomic weight of an atom, or the molecular weight of a molecule, is its mass relative to that of a hydrogen atom

– Specified in Daltons

• One Dalton equals to the mass of a hydrogen atom

CHEMICAL BONDS

Types of chemical bonds

• There are two types of chemical bonds between atoms:

– an ionic bond is formed when electrons are donated by one atom to another (example: NaCl)

– a covalent bond is formed when two atoms share a pair of electrons

Important properties of bonds

• Bond strength (amount of energy that must be supplied to break a bond)

• Bond length: the distance between two nuclei

• Bond orientation: bond angles determining the overall geometry of atoms

The three-dimensional structures of molecules are specified by the bond angles and bond lengths for

each covalent linkage

COVALENT BONDS

Properties of covalent bonds

• Bond strength: The strongest bonds

• Bond length: variable

• Bond orientation: specific bond angles determining the overall geometry of atoms

The three-dimensional structures of molecules are specified by the bond angles and bond lengths for

each covalent linkage

Single and double bonds

• Most are single bonds

• Some are double bonds

Single vs. double bonds

• O, N, S, P, and C atom allow double bonds

• Double bonds are shorter and stronger

• A single covalent bond allows rotation of a molecule

Polarity of covalent bonds

• Covalent bonds in which the electrons are shared unequally in this way are known as polar covalent bonds

Examples

• Oxygen and hydrogen

• Nitrogen and hydrogen

• Not carbon and hydrogen

• Oxygen and nitrogen atoms are electronegative

• Water is an excellent example of polar molecules

NON-COVALENT INTERACTIONS

What are they?

• Reversible and relatively weak

• Electrostatic interactions, hydrogen bonds, and van der Waals interactions

Electrostatic interactions

(charge-charge interactions)

• Formed between two charged particles

• These forces are quite strong in the absence of water

Hydrogen bonds

• The hydrogen atom in a hydrogen bond is partly shared between two relatively electronegative atoms

Donor and acceptor

van der Waals interactions

• The distribution of electronic charge around an atom changes with time

• The strength of the attraction is affected by distance

Hydrophobic interactions

• Not true bonds

CARBON

Why is carbon important?

• It can form single, double, or triple bonds

• Different geometries

– Rotation

• Stable

• Internediate electronegativity

– Hydrophilic vs. hydrophobinc

• Chains and rings

– backbone

• Versatile three-dimensional structure

WATER

Polarity of water

• Water accounts for about 70% of a cell's weight

• In the water molecule, oxygen is more electronegative than hydrogen; therefore, the oxygen side of the molecule has a negative charge and the other side has a positive charge

Hydrogen bonds

• Each water molecule can form hydrogen bonds through its two H atoms to two other water molecules, producing a network

Properties of water

• Polar molecule

– Bent, not linear, the charge distribution is asymmetric

• An excellent solvent

– It weakens electrostatic forces and hydrogen bonding

– Small size

• Highly cohesive

– Networks of hydrogen bonds

• Reactive

– Nucleophile

• Ionization

ORGANIC COMPOUNDS AND

FUNCTIONAL GROUPS

Functional groups

• Groups of atoms attached to carbon skeleton – Usually hydrophilic

Functional groups

• Hydroxyl group (-OH)

– -Alcohols. eg. ethanol, sugars, phenol

– -Dissolve in water (sugars)

• Carbonyl group (C=O)

– aldehyde

– ketone

• Carboxyl group (-COOH)

– Carboxylic acids

• formic acid, acetic acid, amino acids

Functional groups

• Amino group (-NH2)

– Amines. eg. amino acids

• Sulfhydryl group (-SH)

– Thiols

• Phosphate group

– Phosphate is formed by dissociation of an acid called phosphoric acid (H3PO4)