The Chemistry of Life

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The Chemistry of Life
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Elements

• Only 90 naturally occurring elements and only 25
  of those are essential to living organisms.
• Four elements make up more than 96 of the human
  body Carbon, hydrogen, oxygen, and nitrogen.
• (trace amounts of P and SCHONPS)

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Periodic Table

• Rows are called periods and are arranged by the
  number of orbits they contain. (2, 8, 8, 16)
• Columns are called families and like families
  they have similar bonding qualities.
• You can take any element and tell how many orbits
  it will have and how many valence electrons it
  will have based on its row and column.
• A zig-zag line separates the metals from the
  non-metals.

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Atoms

• An element has only one
• type of atom.
• From the Greek word atomos, which means unable
  to be cut
• A million atoms placed side by side would make a
  row only 1 centimeter long
• Contain subatomic particles
• Protons, neutrons electrons

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Nucleus

• At the center of the atom
• Contain positively charged protons and neutral
  neutrons that are held together by strong forces.
• Protons neutrons are about the same size and
  make up all the mass of the atom.

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Electrons

• Negatively charged electrons orbit around the
  nucleus much like the earth orbits around the sun
  
• They are much smaller than protons and neutrons
  about 1/1840 the mass of a proton, so they do not
  add to the atomic mass.
• Electrons are constantly in motion.
• They are attracted to the positively charged
  nucleus.

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The overall charge of an atom is NEUTRAL

• Atoms have the same number of positively charged
  protons as they do negatively charged electrons.
• Even though their masses are different, their
  charges still count equally.
• 2 protons() and 2 electrons(-) would equal a
  neutral charge.

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Isotopes

• An element is identified by its number of
  protons, called the atomic number.
• The atomic mass is the number of protons plus the
  number of neutrons in the nucleus.
• An element can have different numbers of
  neutrons, but not different number of protons or
  it would be a different element.
• These different versions of the same element are
  called isotopes.
• Isotopes will have the same chemical properties.

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Atomic mass

• The atomic mass listed on the periodic table is
  the average of the isotopes of the element found
  on earth multiplied by the percentage at which
  they are found.
• The mass is calculated this way in order to more
  accurately represent the way the element is found
  in nature.

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Radioactive Isotopes

• Some isotopes of certain elements are unstable
  and will break down over a period of time.
• This means they will lose neutrons at a constant
  rate and change into other isotopes of the same
  element.
• Since they will not lose protons, they are still
  the same element, just different versions of that
  element.

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Uses of radioactive isotopes

• Carbon dating using Carbon 14, which breaks down
  to Carbon 13 Carbon 12.
• As isotopes break down, they give off radiation
  and can be used to treat cancer, as tracers to
  follow the movement of substances in organisms,
  and to kill bacteria that cause food to spoil.

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Compounds

• When 2 or more elements form a stable union, it
  is called a compound or a molecule.
• A compound is a combination in definite
  proportions such as H2O is always water, H2O2 is
  hydrogen peroxide.
• A molecule is just the smallest form of a
  compound and the two terms are used
  interchangeably.
• They can be same element or different elements.

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• Every molecule has unique characteristics that
  differentiate it from every other molecule.
• Even a small change can make a big difference
• Glucose and Fructose are both C6H12O6, but just
  where their atoms are located make them different
  and we can even taste the difference.

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Molecules are important in life

• DNA and RNA are distinguished only by the removal
  of one oxygen atom.
• DNA is double stranded
• and more stable than the
• single stranded RNA

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Chemical Reactions

• Compounds and molecules are created when bonds
  are formed between elements.
• REACTANTS -------? PRODUCTS
• Na I2 ? 2 Na I
• Balancing Equations

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Bonds

• Chemical bonds have different strengths.
• Ionic Bonds are strong and form crystals, which
  take a lot of energy to break apart, but they
  easily dissolve in water.
• Covalent bonds are the most common.
• Hydrogen bonds are the weakest, but are strong
  when there are many of them together.

Bond Animation
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Ionic Bonds

• Ionic bonds occur between metals and non-metals
  on the periodic table.
• The valence electrons are transferred from the
  metal to the nonmetal.
• An ion is a molecule that creates ions and the
  positive and negative charges created draw the
  atoms very close together.
• The atoms involved have a strong need to fill or
  empty an orbit.

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Table Salt Crystal
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Ionic Bonding

• Turn to your periodic table and examine the two
  columns headed by Li (ignore hydrogen), Be.
• These columns provide most (not all) of the
  positive partners involved in ionic bonding that
  a high school kid will be held responsible for.

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Ionic Bonds

• The first column (called the alkali metals) has
  Li, Na, K, Rb, Cs, and Fr
• All these guys go 1 in ionic bonding.
• The second column (called the alkaline earth
  metals) has Be, Mg, Ca, Sr, Ba and Ra. All go 2
  in ionic bonding

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Ionic Bonds

• Look to the column headed by F and below it,
  you'll see Cl, Br, I and At
• These elements will all gain one electron in
  ionic bonding and will therefore be negative one.
  
• The next column to the left is headed by O. The
  most common examples used from this column are O
  and S. Se and Te get used sparingly
• Gaining two electrons makes these atoms become a
  negative two charge in ionic bonding.

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Electronegativity

• Electronegativity is a measure of the tendency of
  an atom to attract a bonding pair of electrons.
• The Pauling scale is the most commonly used.
  Fluorine (the most electronegative element) is
  assigned a value of 4.0, and values range down to
  caesium and francium which are the least
  electronegative at 0.7.
• If two elements have an electronegativity greater
  than 1.9, they will form an ionic bond.

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Electronegativity

• Metals have low electronegativities (less than
  2.0), while non-metals have high
  electronegativities (above 2.0)
• Thus, nonmetals from the right side of the chart
  will ionically bond with metals from the left
  side of the chart.
• The only way you would know electronegativity is
  if you were given a chart.

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Covalent Bonds

• A covalent bond forms when electrons are shared
  instead of transferred between atoms.
• Atoms can share one electron for a single
  covalent bond, and even 6 electrons for a triple
  bond.
• The water molecule is a great example of covalent
  bonding.

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H2O
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Methane CH4
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Hydrogen Bonds

• The positive hydrogen in polar molecules forms
  weak bonds with the negative poles of other polar
  molecules.
• These weak bonds are called HYDROGEN BONDS and
  are very important in biological molecules.
• Hydrogen bonds allow DNA to be stable, yet easily
  pulled apart to be read for instructions.

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van der Waals forces

• Intermolecular forces that hold molecules
  together based on their slight polar charges.
• Little fibers on the foot of a Gecko allows it to
  climb walls using the van der Waals forces.

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Properties of Water
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Biologically Important

• ¾ of the earths surface is covered with water.
• Water makes up about 70 of the mass of the human
  body.
• Life as we know it could not exist without water.

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Types of Chemical Bonds
Diff. In electronegativity Type of Bond
Less than 0.5 Non-polar covalent
0.5 to 1.9 Polar Covalent
Greater than 1.9 Ionic
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• Shape of the water molecule

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• Electrons are not shared equally.
• Polar Covalent with partial positive and partial
  negative charges.

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• Hydrogen bonds form between water molecules.
• Individual bonds are relatively weak
• But multiple bonding holds molecules close
  together making it overall fairly strong.

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Hydrogen Bonds

• An intermolecular force of attraction
• A relatively weak bond (about 1/10 the strength
  of a covalent bond).
• Occurs between a hydrogen atom and an oxygen,
  nitrogen, fluorine.

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Thermal Properties of Water

• Unusually high melting and boiling points
  compared to other molecules of similar mass.
• Resists phase change
• Indication of how strongly the molecules of water
  are attracted to each other.

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High heat capacity

• calorieThe amount of energy required to raise a
  defined amount of a substance by one degree.

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Effect on Body Temperature

• Large change in surrounding temperature produces
  a small change in temperature of water.
• Helps maintain a fairly constant body temperature
  (37o C)
• Essential for normal biochemical processes

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High heat of vaporization

• Changing from liquid to gas
• Evaporation of water from the body surface is
  very efficient cooling system.
• Large amount of heat can be carried away by a
  relatively small amount of water (perspiration)
  being vaporized.

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Water is the universal solvent

• Many vital substances can be dissolved in water
• A solution is a mixture in which one or more
  substances (solutes) are distributed evenly in
  another substance (solvent).
• One substance is dissolved evenly in another
  substance and will not settle out.
• Example kool-aid

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pH value

• The pH value of a substance is directly related
  to the ratio of the hydrogen ion and hydroxyl ion
  concentrations.
• If the H concentration is higher than OH- the
  material is acidic.
• If the OH- concentration is higher than H the
  material is basic.
• 7 is neutral, lt is acidic, gt7 is basic

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• Acids taste sour, are corrosive to metals, change
  litmus (a dye extracted from lichens) red, and
  become less acidic when mixed with bases.
• Bases feel slippery, change litmus blue, and
  become less basic when mixed with acids.

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