THE CHEMISTRY OF LIFE

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THE CHEMISTRY OF LIFE

• Chapter 02 03

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• A. Matter
• Material that takes up space.
• 1. Elements
• Pure chemical substances composed of atoms.
• Examples?
• How many elements exist?
• How many of these elements are essential to life?

109 elements have thus far been named (92 are
naturally occurring 17 are synthetic). 25
elements are essential to life.
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Periodic Table of Elements
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Bulk elements needed in large quantities. H,
C, N, O, P, S, Cl, K, Ca, etc Trace
elements needed in small quantities. ex.
Iodine is required for the thymus gland to
produce its hormones Zinc is required to
produce chlorophyll A few elements (arsenic,
bromine tin) are toxic in large amounts, but
may be vital in very small amounts.
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• 2. Atom
• The smallest piece of an element that retains
  the characteristics of that element.
• Composed of 3 subatomic particles
• Protons
• Neutrons
• Electrons

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Characteristics of Subatomic Particles
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• Atomic number
• protons in nucleus of an atom (establishes
  identity of the atom)
• Since most atoms are electrically neutral, atomic
  number indicates of electrons as well.
• Atomic mass
• protons plus neutrons in nucleus of an atom

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• How can we determine the number of neutrons in an
  atom?
• neutrons atomic mass - atomic
• Determine neutrons in a carbon atom (atomic
  mass 12 atomic 6).
• neutrons 12 - 6 6
• Do all carbon atoms have the same number of
  protons?
• Do all carbon atoms have the same number of
  neutrons?

All carbon atoms have the same number of protons,
but they may have differing numbers of neutrons.
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• Isotopes
• Atoms having the same number of protons, but
  differing numbers of neutrons.
• Ex. Carbon isotopes
• carbon 12 (12C) ? 6 neutrons
• carbon 13 (13C) ? 7 neutrons
• carbon 14 (14C) ? 8 neutrons

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Since isotopes have differing numbers of
neutrons, they have differing masses. Neutrons
determine nuclear stability. Atoms with equal
numbers of protons neutrons are most
stable. Unstable isotopes tend to break down into
more stable forms. When they break down they
release radioactive energy. Often one isotope of
an element is very abundant others are rare
(99 of carbon isotopes have 6 neutrons).
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Periodic table information on carbon
Atomic mass given in table is average mass of all
the elements isotopes.
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• 3. Compound
• A pure substance formed when atoms of different
  elements bond.
• The number of atoms of each element is written as
  a subscript.
• Examples
• CO2 carbon dioxide
• H2O water
• CH4 methane
• C6H12O6 glucose

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• 4. Molecule
• Smallest piece of a compound that retains
  characteristics of that compound.
• The number of molecules is written as a
  coefficient.
• Examples
• 4CO2 4 molecules of carbon dioxide
• 2C6H12O6 2 molecules of glucose
• 6O2 6 molecules of oxygen

NOTE some molecules are diatomic, consisting
of two atoms of the same element. Diatomic
molecules include oxygen (O2), nitrogen (N2),
hydrogen (H2) and chlorine (Cl2). These diatomic
molecules are not considered to be compounds by
definition.
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• 5. Chemical Bonds
• Type of bond formed is determined by the number
  of valence electrons in the interacting atoms
  octet rule.
• a) Covalent bonds - form when atoms share
  electron pairs.
• strongest type of bond
• tend to form when atoms have 3, 4 or 5 valence
  electrons
• can be nonpolar or polar

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Lines between the interacting atoms indicate a
shared pair of electrons. The compounds seen here
are commonly called hydrocarbons because they
contain C H only.
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• Nonpolar covalent bonds - electrons are shared
  equally between atoms.
• Ex. methane

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Polar covalent bonds - electrons are drawn more
strongly to 1 atoms nucleus than the other. Form
when less electronegative atoms bond with more
highly electronegative atoms. Ex. water
Electronegativity tendency of an atom to
attract electrons. Oxygen is highly electronegat
ive. Hydrogen has a low electronegativity. The
oxygen atom pulls hydrogens negatively
charged electrons away from hydrogens nucleus
an towards its own. Thus, the oxygen atom has a
slight negative charge, while both hydrogens
have a slight positive charge.
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• b) Ionic bonds - form when oppositely charged
  ions are attracted to each other.
• weaker than covalent bonds
• atoms with 1, 2 or 3 valence electrons give up
  electrons to atoms with 7, 6 or 5 valence
  electrons
• form salts
• Ex. NaCl

Sodium ion (Na) sodium atom that has lost an
electron. Chlorine ion (Cl-) chlorine atom that
has gained an electron.
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• c) Hydrogen bonds - form when opposite charges on
  two molecules are attracted to each other.
• weakest type of bond
• Ex. DNA

H2O
Hydrogen bonds are individually weak, but a
large number of them together will provide a
great deal of strength, like the teeth of a
zipper. Note Hydrogen bonds are located between
separate water molecules (here we see 4 hydrogen
bonds) however, within each water molecule, the
H atoms are bonded to the C by covalent bonds!
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• B. The Importance of Water
• 1. Properties
• Cohesion - the attraction of water molecules for
  each other.
• Adhesion - the attraction of water molecules for
  other compounds.
• High heat capacity takes a great deal of heat
  to raise the temperature of water.

Movement of water from a plants roots to its
highest leaves depends upon cohesion of water
molecules their adhesion to the plants
water-conducting tubes. Waters high heat
capacity allows organisms to be exposed to
extremes of temperature. Their body fluids heat
or cool slowly.
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• High heat of vaporization - a lot of heat is
  required to evaporate water.
• Exists as solid, liquid or gas - solid (ice) is
  less dense than liquid.
• 2. Solutions
• A solution is a mixture of one or more solutes
  dissolved in a solvent.
• If solvent is water, then it is an aqueous
  solution.
• Water is a strong solvent because it separates
  charged atoms or molecules.

Because water has a high heat of vaporization,
the body cools off quickly as sweat evaporates
from its surface. Because ice is less dense than
water, it floats on the surface, forming a cap
that retains heat in the water below.
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• 3. Acids Bases
• Acids - substances that add H to a solution.
• Bases - substances that remove H from solution.
• pH scale is measure of acidity/alkalinity based
  on H concentration.

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• Acids have pH values that range below 7.0,
• with a pH of zero being the most acidic.
• Bases have pH values that range above 7.0,
• with a pH of 14 being the most alkaline.
• Pure water has a pH of 7.0 and is considered
• to be neutral.
• Each unit on the pH scale represents a 10
• fold change in H concentration. Thus, the
• H concentration of soap (pH10.0) is 100
• times greater than that of household ammonia
• (pH12.0).

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• C. Major Organic Molecules
• Molecules that contain carbon in combination with
  hydrogen.
• 1. Carbohydrates
• contain C, H O ? C ? ? O
• function to store energy provide support
• building blocks (monomers) are monosaccharides

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• Monosaccharides
• simple sugars containing 3 - 7 carbons.
• C, H, O ratio is 121

• Common monosaccharides include
• glyceraldehyde (3 carbon sugar)
• ribose (5 carbon sugar)
• deoxyribose (5 carbon sugar)
• glucose - blood sugar (6 carbon sugar)
• fructose - fruit sugar (6 carbon sugar)
• galactose (6 carbon sugar)
• Note glucose fructose galactose have the same
  molecular formula
• C6H12O6. Thus, they are isomers.

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• Disaccharides
• simple sugars composed of 2 monosaccharides
  linked together by dehydration synthesis.

Other common disaccharides maltose (seed sugar)
lactose (milk sugar).
Glucose Fructose ? Sucrose (table sugar)
water Because water is formed, sucrose will have
2 less hydrogen atoms 1 less oxygen atom that
glucose fructose combined. Note hydrolysis
(splitting by adding water) breaks disaccharides
apart yielding component monosacchrides. Glucose
Glucose ? Maltose Glucose Galactose ?
Lactose
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• Polysaccharides
• complex carbohydrates made up of hundreds of
  monomers linked by dehydration synthesis.

• Common polysaccharides include
• cellulose - forms wood parts of
• plant cell walls.
• starch - energy storage form in plants.
• glycogen - short term energy storage
• form in animals.
• cellulose, starch glycogen are long
• chains of glucose units differ in branching
  patterns chitin - forms the
• exoskeletons of arthropods cell wall in many
  types of fungi.
• Cellulose is most common organic compound in
  nature.
• Chitin is second most common polysaccharide in
  nature.

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• 2. Lipids
• contain C, H, O ? C gtgt ? O
• do not dissolve in water
• Triglycerides (fats)
• composed of glycerol linked to 3 fatty acid
  chains by dehydration synthesis.
• function to cushion organs, as insulation in
  long-term energy storage (adipose tissue).

Number of carbon atoms always much greater than
number of oxygen atoms.
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Note many more C atoms than O atoms. Saturated
fats - fatty acids are saturated with H atoms.
There are no double-bonded carbons. Tend to be
liquid at room temperature. Unsaturated - fatty
acids are not completely saturated with H atoms
(thus have 1 or more double bonded
carbons). Double bonds cause kinks in the fatty
acid tails. Tend to be solid at room
temperature. Lipids in plants are less
saturated than those in animals.
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• Phospholipids
• lipid bonded to a phosphate group
• major component of cell membranes

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• Sterols
• lipids that have 4 interconnected carbon rings
• Ex. Vitamin D, cortisone, estrogen cholesterol

• Waxes
• fatty acids combined with hydrocarbons
• help waterproof fur, feathers, leaves fruits

Cortisone is a steroid hormone. Cholesterol is a
key component of cell membranes.
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• 3. Proteins
• contain C, H, O, N, (S)
• monomers are amino acids

• An amino acid contains
• a central carbon atom
• bonded to
• a hydrogen atom
• a carboxyl group (COOH)
• an amino group (NH2)
• an R group - differs for
• each of the 20 biologically important amino
  acids.

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• Proteins have a 3-dimensional shape
  (conformation)
• primary (1o) structure - amino acid sequence of
  polypeptide chain
• secondary (2o) structure - coiling folding
  produced by hydrogen bonds
• tertiary (3o) structure - shape created by
  interactions between R groups
• quarternary (4o) structure - shape created by
  interactions between two or more polypeptides

Conformation of protein is critical to its
function.
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• Examples
• Antibodies Function in immunity.
• Hemoglobin protein in RBCs that transports
  oxygen.
• insulin glucagon protein hormones that
  regulate levels of glucose in the bloodstream.
• Keratin structural protein found in hair, nails,
  hooves.
• fibrin thrombin proteins involved in blood
  clotting.
• spider silk (strongest natural fiber known)
• Enzymes (maltase, pepsin, lipase)
• Enzymes are protein catalysts (speed up chemical
  reactions without being altered in the process
  thus they are reusable).

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• 4. Nucleic Acids
• contain C, H, O, N, P
• monomers are nucleotides

• Each nucleotide is
• composed of
• a 5 carbon sugar (ribose or
• deoxyribose)
• a phosphate group
• a nitrogenous base (guanine,
• cytosine, thymine, adenine or uracil).

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• DNA (deoxyribonucleic acid)
• 5-carbon sugar is deoxyribose
• nitrogenous bases are A, G, C T
• double-stranded helix held together by hydrogen
  bonds
• is the genetic material

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• RNA (ribonucleic acid)
• 5-carbon sugar is ribose
• nitrogenous bases are A, G, C U
• single-stranded
• enables information in DNA to be expressed