Chemistry Lecture Text Chapter 2. Chemistry in Physiology Physiology requires some familiarity with...
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Transcript of Chemistry Lecture Text Chapter 2. Chemistry in Physiology Physiology requires some familiarity with...
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