Chemistry Review

Atoms, Ions, and Molecules Atom – the basic unit of matter that consists of a nucleus and outer energy levels, and an equal number of electrons and protons. Ex. Ca, Na, O Ion – A charged atom or groups of atom. Ex. Ca+2, Na+1, O-2 Molecule – two or more atoms joined by one to three covalent bonds. Ex. O=O (O2), C6H12O6

Isotopes Isotopes are different forms of the same element that have the same number of protons but a different number of neutrons. Some isotopes are radioactive (naturally break down into other elements). Ex. Carbon-12 (12C): The most prevalent form of carbon. Carbon-14 (14C): A radioactive form of carbon that is used for carbon dating dead organic material. C-14 half life is 5,730±40 years.

Some Common Isotopes Used in Medicine Iron-59 (46 d): Used in studies of iron metabolism in the spleen. Potassium-42 (12 h): Used for the determination of exchangeable potassium in coronary blood flow. Rhenium-186 (3.8 d): Used for pain relief in bone cancer. Beta emitter with weak gamma for imaging. Rhenium-188 (17 h): Used to beta irradiate coronary arteries from an angioplasty balloon. Cobalt-57 (272 d): Used as a marker to estimate organ size and for in-vitro diagnostic kits. Gallium-67 (78 h): Used for tumour imaging and localisation of inflammatory lesions (infections).

Types of Bonding Intramolecular Bonding: Strong bonds that form WITHIN compounds. Covalent Bonds – bonds between atoms in which outermost energy level electrons are shared between atoms. Result: Formation of molecules. Polar Covalent Bonds – the electrons are NOT equally shared between the atoms. Result: dipole formation. Nonpolar Covalent Bonds – The electrons are equally shared between the two bonding atoms. Ionic Bonds – electrostatic forces of attraction between positive and negative ions. Result: Ionic Compound formation.

Types of Bonding Intermolecular Bonding: forces of attraction that form BETWEEN different compounds. *** Not as strong as intramolecular bonding. Hydrogen Bonds – H atoms are shared. Requirement: The H atom MUST be bonded to either F, N, or O (highly electronegative) and are strongly attracted to an F, N, or O from another compound. Strongest type of intermolecular bonding Dipole Forces – Attraction between the positive dipole in one molecule to the negative dipole in another molecule. London Forces (van der Waals) – the weakest attractions between oppositely polarized electron clouds in two different compounds. See table 2.1, page 23

“Like Dissolves Like” The major solvent in living system is water. Water is a polar molecule. Polar molecules and ions dissolve in water. We call these substances hydrophilic (“water loving”). Nonpolar solvents dissolve nonpolar molecules. Nonpolar molecules do not dissolve in water. We call them hydrophobic (“water fearing”). **NOTE: Hydrocarbons – compounds composed of carbon and hydrogen only – are nonpolar and therefore hydrophobic!

The Special Properties of Water At room temperature, water is a liquid. This is due to its ability to form hydrogen bonds with other water molceules. Water also is able to form hydrogen bonds with other molecules that have a hydrogen bonded to nitrogen, fluorine, or oxygen (highly electronegative atoms).

The Special Properties of Water A water molecule with its dipoles. A water molecule is able to form 4 hydrogen bonds with other water molecules. Note the positive dipole (of the hydrogens) are attracted to the negative dipole (of oxygen) on the other molecules.

The Special Properties of Water The hydrogen bonding of water and its small size enables it to have the following special properties: Cohesion Adhesion High specific heat capacity Lower density as a solid than a liquid. Universal solvent High heat of vaporization

H2O (l)  H+(aq) + OH-(aq) Acids and Bases The acidity or basicity of a solution is dependent upon the number of H+ ions (H3O+) in comparison to the number of OH- ions. H2O (l)  H+(aq) + OH-(aq) An acid has more H+ ions than OH- ions. A base has more OH- ions than H+ ions.

Acids and Bases The pH scale measures the hydrogen ion concentration of a solution: pH = - log [H+] 0-6.9  acidic 7.1 – 14.0  basic

Buffers The pH within cells and the extracellular fluid (ECF) surrounding cells of multicellular organisms remains constant (~pH = 7). Enzymes will denature and not work (or work well) if the pH changes even slightly . Buffers are substances found within living organisms to help them resist pH changes.

Buffers How do buffers maintain pH? Buffers act in a way to keep [H+] constant. They consist of an acid and base pair. A buffer will release H+ ions when a base is added. A buffer will absorb H+ ions when an acid is added.

Buffers Example of an Acid/Base Buffer Pair: A Blood Buffer Carbonic acid (H2CO3)  acid Bicarbonate (HCO3-)  base

Buffers This figure shows the major organs that help control the blood concentrations of CO2 and HCO3-, and thus help control the pH of the blood. Removing CO2 from the blood helps increase the pH. Removing HCO3- from the blood helps lower the pH.