Basic Chemistry Background

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If you are a member of this course, you should have taken Chemistry as a pre-requisite.  In this unit,  we review some basic concepts that relate to Physiology 3.  Members of this class are expected to read Chapter 2 for this information and more.  As tempting as it is, and as busy as we are, do not skip the chapter reading.

Matter

  • Matter is anything that occupies space and has mass.  Mass is a measure of how much matter there is
  • Units of mass are usually in kilograms.  There are 92 different kinds of natural matter, called elements.
  • The major elements are H, O, C and also Nitrogen, but others are crucial, Ca, Na, Cl.

Atomic Structure

An atom is the smallest division of an element that has all the chemical and physical properties of that element. 

  • Subatomic particles:
    Protons + charge;
    Electrons – charge;
    Neutrons are neutral
  • The overall charge on an atom is neutral
  • Hydrogen – has one proton and one electron
  • Carbon – has six protons and six neutrons and six electrons
  • Atomic number = number of protons in the nucleus;
    • Atomic number of Hydrogen is 1;  Hydrogen just has one proton and one electron but has no neutrons
    • Atomic number of Carbon is 6
  • Atomic weight = AKA atomic mass - the number of protons + the number of neutrons (electrons have no mass);
    • atomic wt of Carbon is 12. 
  • Orbitals (or shells) are energy levels that surround the nucleus of an atom. 
    • The first shell holds 2 electrons.
    • Each shell thereafter holds 8 electrons. OCTET RULE
    • Atoms are most stable when the outer shell is filled.
    • Electrons in unfilled outer shells participate in bonding; they are called valence electrons.  This is the driving force for bond formation.

Isotopes 

  • They have the same number of protons but a different number of neutrons.
  • Some isotopes are radioactive but many are stable.

Covalent Bonds 

  • Are formed when electrons are shared between two atoms.
  • Hydrogen will only be stable if it has two electrons in its outer shell, so it must pair up with another electron to achieve stability.

Structural Representations of Molecules 

  • there are 3 different modes of drawing out a molecule:
    • space filling models,
    • showing orbitals, or
    • just letters and lines to represent bonds.

Polar Covalent Bonds 

  • form when there is unequal sharing of electrons
  •  Due to electronegativity OF OXYGEN
  • This also happens with nitrogen and fluorine and phosphate
  • With water: oxygen pulls negative charge towards itself leaving Hydrogen positive
  • Polar molecules make great solvents.  Sodium chloride dissolves well in it. Physiology

Ionic Bonds

  • An ion is an atom that has gained or lost an electron
  • ionic bonds form. when atoms give an atom to another atom, and when that other atom accepts the electron
  • For example, Sodium gives up an electron to Chlorine so they are both stable.

Solubility and Polarity

  • Sodium Chloride dissolves in water:
    • The positive sodium bonds with the negative oxygen of water
    • The negative Chloride bonds with the positive hydrogen
  • Hydrogen spheres make a molecule soluble in water; the ability to make hydrogen spheres indicates that a molecule is polar.
  • Polar solutes are charged. This makes them hydrophilic.
  •  Polarity is due to 1) uneven electron distribution due to electronegative atoms. 2) Geometric asymmetry in the shape of the chemical, 3) lots of hydroxyl groups in a molecule will make it polar, b/c the OH groups are attracted to water.
  • “Ionic solvents are polarity taken to the extreme, so that we actually have two charged ions”
  • Nonpolar solutes, such as fats, are hydrophobic, AFRAID of water. 
    lipids float on water, vinegar and oil.

Hydrogen Bonds 

  • An attraction between the hydrogen on one water molecule and an oxygen on another water molecule.
  • They form between electronegative elements (F, N, O) and hydrogen.
  • Found in DNA and in proteins
  • Broken by temperature

Acids, Bases and the pH Scale 

  • Pure water ionizes to make equal amounts of H+ and OH-.  This solution is neutral.
  • Acid – A solution that gives off H+
  • Bases – a solution that accepts H+
  • For example, NaOH is a base because it removes H+ from solution:
  • NaOH + H+ --> Na+ + H2O
  • The pH scale indicates [H+]
  • pH is the negative log of the [H+]
  • [H+] is .0000001 moles/liter, or 10-7 M
  • - log[H+] = -log[10-7] = 7.0
  • pH values are ten-fold different

Strength of Acids

  • Strong acids – when you put them in water, they dissociate completely 100%.
    e.g. HCl --> H+ and Cl-
  • Weak acid – when you put them in water, they will dissociate but only weakly.
  • Strong bases and weak bases also exist.  We do not have strong bases in the body.

Buffers pH 

  • A buffer is a chemical that minimizes changes in pH.  It reduces the excess acidity or alkalinity.
  • HCl is a strong acid, but when it combines with sodium bicarbonate, a buffer, carbonic acid (a weak acid) forms, along with a salt.
  • Therefore, sodium bicarbonate is a buffer b/c it removes a strong acid

  • Bicarbonate buffer in the blood

    • Blood pH is maintained between  pH 7.35 and 7.45
    • In blood plasma, the pH is stabilized by this reversible reaction:
      • Bicarbonate ion (HCO3-)  + H+   Carbonic acid (H2CO3)
    •  A strong acid, H+, is made into a weak acid, H2CO3

Organic Molecules

  • Any molecule that has carbon and hydrogen is an organic molecule.
  • They form because Carbon can make up to four bonds so it can make chains, circles, double bonds
  • Therefore CARBON makes relatively inactive “backbone” molecular structures to which “functional groups” may be added.
  • There are many organic acids, such as amino acids, which make up proteins.

  • Functional Groups

    • These are reactive chemical groups which are responsible for the unique chemical properties of the molecule.

    • Classes of organic compounds can be named according to their functional groups.abbreviated COOH

Stereoisomers

  • Molecules with the same sequence of atoms, which differ from each other because of the way their atoms are arranged three-dimensionally in space.
  • Example: 1) CIS AND TRANS
  • Example 2) Enantiomers (optical isomers) - mirror images of each other.
    • They are like left- and right-handed gloves: if the palms are facing the same direction, they cannot be superimposed on each other. 
  • If you buy a vitamin which is the wrong stereoisomer, it will not work for you.  There are laws regulating this.

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