Chemistry

Lecture Text Chapter 2

Chemistry in Physiology

• Physiology requires some familiarity with basic chemistry– atomic and molecular structure– chemical bonds– pH– organic compounds (next week)

Atoms

• smallest units of matter that can undergo chemical change

• made up of three basic subatomic particles– protons – positively charged

– neutrons – neutrally charged

– electrons – negatively charged particles

The Nucleus

• Nucleus = central body– Contains protons and

neutrons

• number of protons determines the element– Fundamental type of matter

The Periodic Table of The Elements

Atomic Number and Mass

• Atomic number– number of protons in an atom

• Atomic mass (weight)– the total number of protons

and neutrons found within an atom

– Isotopes = atoms of the same element with different atomic masses

Electrons

• Revolve around the nucleus in certain volumes of space called orbitals

• Several such orbitals: – innermost can hold two

electrons– second layer can hold eight

electrons– valence electrons = electrons

in the outer shell

Electrons and the Periodic Table

• Elements are arranged in columns by the # of valence electrons

• atoms are most stable when the outermost orbital is full

• most elements do not have full sets of valence electrons

Chemical Bonds

• Atoms may give, take or share electrons in order to achieve full outer shell – link two or more atoms

together through chemical bonds

– molecules – structures consisting of atoms bound together by chemical bonds

Types of Chemical Bonds

1. Covalent bonds

2. Ionic bonds

3. Hydrogen bonds

Covalent bonds

• two or more atoms share their valence electrons

• Nonpolar molecules– atoms share electrons

equally

• Polar molecules– Unequal sharing of

electrons– Unequal charge between

different regions of the molecule

Ionic Bonds

• Between metal and non-metal

• One or more valence electrons completely transferred from one atom to another

• Forms ions – atoms or molecules with

unequal numbers of protons and electrons

Ionic Bonds• Cations

– Positive charge– More protons than electrons– Metals

• Anions– Negative charge– More electrons than protons– Non Metals

• Attract each other– form ionic compound

Dissociation of Ionic Compounds

• ionic bonds tend to be weak– Can dissociate in water– Water attracted

electrostatically– forms hydration spheres

around molecules

Water Solubility

• Hydration sphere formation determines water solubility

• Hydrophilic – Water soluble

– Polar molecules and ions

• Hydrophobic – Water insoluble

– Nonpolar molecules

Hydrogen bonds

• Polar molecules have weak electrostatic attraction for one another– Slight negative end to

slight positive end

• Responsible for water properties, protein shape, DNA structure, etc.

Acidity and Alkalinity

• Sometimes water molecules will split– Covalent bond between

oxygen and a hydrogen will be broken

– Form H+ (hydrogen ion) and OH- (hydroxide ion)

– H2O H+ + OH-

Acidity and Alkalinity

• In pure water, equal amounts of H+ and OH- are formed– Generally, [H+] = 1 x 10-7 M (= 0.0000001 M)

– Neutral solution

• Some solutes (acids) release H+ when mixed with water [H+] above [OH-]

– Acidic solution

• Some solutes (bases) bind H+ or release OH- when mixed with water [H+] below [OH-]

– Alkaline or Basic solution

pH

• Index of [H+] in a solution• Quantify acidity or alkalinity of a solution

pH = log(1/[H+])

• Example: for pure water [H+] = 1 x 10-7MpH = log (1/0.0000001) = log (10,000,000) = 7

pH

• Solutions w/ pH = 7.0 are neutral

• Solutions w/ pH < 7 are acidic– [H+] > 1x10-7 M

• Solutions w/ pH > 7 are alkaline– [H+] < 1x10-7 M

pH

• pH can range from 0 to 14• As pH increases, [H+]

decreases• A difference of 1.0 in pH

means a 10x difference in [H+]– A solution of pH 7 has 10x

the [H+] of a pH 8 solution