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Chemistry in Physiology
Chemistry
Lecture Text Chapter 2
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 Periodic Table
of The Elements
• Physiology requires some familiarity
with basic chemistry
– atomic and molecular structure
– chemical bonds
– pH
– organic compounds (next week)
The Nucleus
• Nucleus = central body
– Contains protons and
neutrons
• number of protons
determines the element
– Fundamental type of matter
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
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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
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
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
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
Types of Chemical Bonds
1. Covalent bonds
2. Ionic bonds
3. Hydrogen bonds
Ionic Bonds
• Between metal and nonmetal
• One or more valence
electrons completely
transferred from one
atom to another
• Forms ions
– atoms or molecules with
unequal numbers of
protons and electrons
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Ionic Bonds
Dissociation of Ionic Compounds
• Cations
– Positive charge
– More protons than electrons
– Metals
• Anions
– Negative charge
– More electrons than protons
– Non Metals
• Attract each other
• ionic bonds tend to be
weak
– Can dissociate in water
– Water attracted
electrostatically
– forms hydration spheres
around molecules
– form ionic compound
Water Solubility
• Hydration sphere
formation determines
water solubility
• Hydrophilic
– Water soluble
– Polar molecules and ions
• Hydrophobic
– Water insoluble
– Nonpolar molecules
Acidity and Alkalinity
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
• In pure water, equal amounts of H+ and OH- are formed
• 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-
– 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
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pH
• Index of [H+] in a solution
• Quantify acidity or alkalinity of a solution
pH = log(1/[H+])
pH
• Solutions w/ pH = 7.0 are neutral
• Solutions w/ pH < 7 are acidic
– [H+] > 1x10-7 M
• Solutions w/ pH > 7 are alkaline
• Example: for pure water [H+] = 1 x 10-7M
– [H+] < 1x10-7 M
pH = log (1/0.0000001)
= log (10,000,000)
=7
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
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