Biology Chapter 2 Chemistry of Life

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					                          Chapter 2 Chemistry of Life
                                     Section 1
                            Composition of Matter
•   Matter is anything that occupies space and has mass.
•   Mass is the quantity of matter an object has.
•   Mass and Weight are not the same
•   Weight is the force produced by gravity acting on mass
•   This is why the same mass has less weight than on the moon
•   Chemical changes in matter are essential to all life processes
•   Elements and Atoms
•    Elements are made of a single kind of atom and cannot be broken down by
    chemical means into simpler substances.
•   More than 100 elements have been identified
•   30 are important to living things
•    More than 90% of the mass of living things is composed of combinations of 4
    elements: oxygen, carbon, hydrogen, and nitrogen
•   Information about elements is summarized on the periodic table
•   The simplest particle of an element is an atom
•   Atoms are composed of protons, neutrons, and electrons.
•   The Nucleus
•   Protons and neutrons make up the nucleus of the atom.
•   Protons have a positive charge
•   Neutrons have no charge
•   The atomic number is the number of protons in an atom
•    The atomic mass number is the total number of protons and neutrons of an
    atom
•   Electrons
•   Electrons move about the nucleus in orbitals.
•   Electrons have a negative (-) charge and little mass
•    An orbital is a three-dimensional region around a nucleus that indicates the
    probable location of an electron.
•    Electrons in orbitals farther away from the nucleus have greater energy than
    electrons that are in orbitals closer to the nucleus
•    Orbitals correspond to specific energy levels (each orbital can hold to
    electrons)
•   The 1st level can hold 2 electrons
•   The 2nd level can hold 8 electrons and so on.
•    The net electrical charge of an atom is zero (equal number of protons and
    electrons)


•   Isotopes
•    Isotopes are atoms of the same element that have a different number of
    neutrons
•   This changes the mass of the element
•   Most elements are made up of a mixture of isotopes
•    The average atomic mass of an element takes this into account (on periodic
    table)
•   Compounds
•    Consist of atoms of two or more elements that are joined by chemical bonds
    in a fixed proportion.
•   Example is water (H2O)
•    How elements combine and form compounds depends on the number and
    arrangement of electrons in their orbitals
•   An atom is chemically stable when orbitals in its highest energy level are filled
    with the maximum number of electrons
•   Examples are the noble gases (Helium)
•   Chemical bonds are the attractive forces that hold atoms together
•   Covalent Bonds
•   A covalent bond is formed when two atoms share electrons. (Figure 2-4)
•   Example is water
•    A molecule is the simplest part of a substance that retains all of the
    properties of that substance and can exist in a free state
•   Examples include a water molecule or a molecule of oxygen gas (O2)
•   Ionic Bonds
•    Formed when one atom gives up an electron to another. The positive ion is
    then attracted to a negative ion to form the ionic bond. (Figure 2-3)
•   An ion is an atom or molecule with an electrical charge
•   Example is salt, sodium chloride (NaCl)
•   Na+ and Cl-
                                        Section 4
                                         Energy
    Energy and Matter
•   Energy is the ability to do work
•    Numerous forms of energy (solar, electrical, chemical, mechanical, thermal,
    geothermal and radiant)
•   Forms of energy important to biological systems are chemical, thermal,
    mechanical, and electrical
•   States of Matter
•   Examples are solid, liquid, gas, or gel
•   A solid maintains a fixed volume and shape
•    A liquid maintains a fixed volume, but the particles move freely which allows
    it to flow and conform to the shape of containers
•    A gas has little or no attraction by its particles and fill a volume of a container
    they occupy
•    Addition of energy to a substance can cause its state to change from a solid to
    a liquid and from a liquid to a gas.


•   Energy and Chemical Reactions
•    A chemical reaction is when one or more substances change to produce one
    or more different substances
•    Energy is absorbed or released when chemical bonds are broken and new
    ones are formed
•   Figure 2-6 (Co2 + H2O ↔ H2CO3)
•   Reactants are substances that enter chemical reactions. (left side of reaction)
•    Products are substances produced by chemical reactions. (right side of
    reaction)
•    Energy your body needs is provided by sugars, proteins, and fats found in
    food
•   Metabolism is all of the chemical reactions that occur in an organism
•   Activation Energy (Figure 2-19)
•   The amount of energy needed to start a reaction
•    Catalysts (enzymes) lower the amount of activation energy necessary for a
    reaction to begin in living systems
•   Enzymes are proteins or RNA molecules that speed up metabolic reactions
    without being permanently changed or destroyed
•   Oxidation Reduction Reactions
•    A chemical reaction in which electrons are exchanged between atoms (also
    called redox reactions)
•    An oxidation reaction is when a reactant loses one or more electrons,
    becoming more positive in charge (Sodium in the NaCl reaction)
•    A reduction reaction is when a reactant gains one or more electrons,
    becoming more negative in charge (Chlorine in the NaCl reaction)
•   Redox reactions always occur together
                                       Section 2
                                 Water and Solutions
    Polarity
•    Water is considered to be a polar molecule due to an uneven distribution of
    charge. (Figure 2-6)
•    The electrons in a water molecule are shared unevenly between hydrogen
    and oxygen.
•   Solubility of Water
•    The polarity of water makes it effective at dissolving other polar substances
    such as sugars, ionic compounds, and some proteins.
•    Water doesn’t dissolve nonpolar substances, like oil because of a weaker
    attraction between polar and nonpolar molecules
•    Dissolved or dissociated ions are present in all of the aqueous solution in
    living things and are important in maintain normal body functions
•   Hydrogen Bonding
•   Waters’ polar nature causes water molecules to be attracted to each other
•    A hydrogen bond is the force of attraction between a hydrogen molecule with
    a partial positive charge and another atom or molecule with a partial or full
    negative charge. (figure 2-7)
•   Hydrogen bonds in water exert an attractive force strong enough to allow
    water to “cling” to itself and other substances
•   Hydrogen bonds form, break, and reform with great frequency (States of
    water)
•   Cohesion and Adhesion
•    Cohesion is an attractive force that holds molecules of a single substance
    together, such as water molecules.
•   This allows the upward movement of water from plant roots to their leaves
•    Surface tension of water is due to the surface molecules being pulled
    downward into the water and forms a thin “skin” on the surface
•    Adhesion is the attractive force between two particles of different
    substances, such as water molecules and glass molecules.
•    Capillarity is the attraction between molecules that results in the rise of the
    surface of a liquid when in contact with a solid
•    Cohesion, Adhesion, and Capillarity help water rise through narrow tubes
    against the force of gravity (Figure 2-8)
•   Temperature Moderation
•    Water has the ability to absorb a relatively large amount of energy as heat
    and the ability to cool surfaces through evaporation.
•   Evaporation prevents organisms from overheating
•   Density of Ice
•   Solid water is less dense than liquid water due to the shape of the water
    molecule and hydrogen bonding
•   This is why ice cubes float in a glass of water
•   Bodies of water Lakes and ponds) freeze from top to bottom
•   Solutions
•    A mixture in which one or more substances are uniformly distributed in
    another substance
•   Solutions can be mixtures of liquids, solids, or gases
•   Example is plasma
•   A solution consists of a solute dissolved in a solvent.
•   A solute is a substance dissolved in the solvent
•   The solvent is the substance in which the solute is dissolved
•    The concentration of a solution is the amount of solute dissolved in a fixed
    amount of the solution
•   The more solute dissolved, the greater the concentration of the solution
•   A saturated solution is one in which no more solute can dissolve
•    Aqueous solutions are solutions in which water is the solvent (important to
    living things)
•   Acids and Bases
•   Ionization of Water
•   Water ionizes into hydronium ions (H3O+) and hydroxide ions (OH–).
•   H2O ↔ H+ + OH- (Hydroxide ion)
•   H+ + H2O ↔ H3O+ (Hydronium ion)
•   Acids
•   Acidic solutions contain more hydronium ions than hydroxide ions.
•   HCL ↔ H+ + Cl-
•    The free hydrogen ions combine with water molecules to form hydronium
    ions
•   Acids are sour tasting and highly corrosive
•   Bases
•   Basic solutions contain more hydroxide ions than hydronium ions.
•   NaOH ↔ Na+ + OH-
•   Alkaline refers to bases
•    Bases taste bitter and feel slippery (the OH- ions react with the oil in our skin
    to form a soap)
•    pH (figure 2-10) “The potential of hydrogen” developed by SØren SØrensen
•     Scientists have developed a scale for comparing the relative concentrations of
     hydronium ions and hydroxide ions in a solution. This scale is called the pH
     scale, and it ranges from 0 to 14.
•    A solution with a pH of 0 to 6.9 is acidic
•    A solution with a pH of 7 is neutral
•    A solution with a pH of 7.1 to 14 is basic
•    A solution’s pH is measured on a logarithmic scale
•    The change of one pH unit reflects a 10-fold change in the acidity or alkalinity
•     pH can be measured by litmus paper, a chemical indicator that changes color,
     or by electronic device
•    Buffers
•     Chemical substances that neutralize small amounts of either an acid or base
     added to a solution
•    Buffering systems maintain pH values in a normal healthy body


                            Section 2-3 Biochemistry

I.       Carbon Bonding

        A. Compounds are classified into two categories

                1. Inorganic compounds- do not contain carbon

                     a. Examples are salt (NaCl), water (H2O) and carbon
                        dioxide (CO2) the exception to the rule

                2. Organic Compounds – contain carbon

                     a. Most matter in living organisms is made of organic
                        compounds
            b. Examples are proteins, lipids, carbohydrates, and
               nucleic acids

B. Carbon atoms

       1. Have 4 electrons in their valence shell (outer electron shell)

       2. It readily forms 4 covalent bonds with other elemental
          atoms

       3. It will bond with other carbon atoms to form straight or
          branched chains, or rings (figure 2-11) (great variety)

       4. A single bond forms when 2 atoms share 1 pair of electrons

       5. A double bond forms when atoms share 2 pairs of electrons

       6. A triple bond forms when atoms share 3 pairs of electrons

C. Functional Groups (Table 3-1 in class)

       1. Clusters of atoms that influence the characteristics of the
          molecules they compose and the chemical reactions the
          molecules undergo

       2. Example is the hydroxyl group (-OH) it makes the molecule
          it attaches to polar

       3. Polar molecules are hydrophilic (soluble in water) ex.
          Alcohol


D. Large Carbon Molecules

       1. Monomers are small, simple molecules that make up carbon
          compounds

       2. Polymers are molecules that are made up of repeated,
          linked monomers
      3. Macromolecules are large polymers composed of 100’s to
         1000’s of atoms

      4. The 4 main types of macromolecules
          a. Carbohydrates
          b. Lipids
          c. Proteins
          d. Nucleic acids

      5. Condensation reactions are chemical reactions that link
         monomers together to form polymers

      6. When a monomer is added to a polymer, a water molecule is
         released

          a. Glucose + Fructose ↔ Sucrose (table sugar) + H2O

      7. Hydrolysis is a reaction that breaks down polymers, by
         adding a water molecule to break a bond linking 2
         monomers

          a. Sucrose + H2O ↔ Glucose + Fructose

E. Energy Currency

      1. Life processes require energy

      2. Adenosine triphosphate (ATP) is a compound that stores a
         large amount of energy

      3. It is made up of
          a. The sugar - ribose
          b. The nitrogen base – adenine
          c. Three phosphate groups

      4. When one of the bonds between the phosphate groups is
         broken, energy is released


                     Molecules of Life
I.   Carbohydrates (Figure 2-13)

     A. Organic compounds composed of carbon, hydrogen, and oxygen

     B. They are always in the ratio of 1 carbon atom to 2 hydrogen atoms
        to 1 oxygen atom

     C. Serve as source of energy and as structural materials

     D. Three types of carbohydrates:

            1. Monosaccharides
            2. Disaccharides
            3. Polysaccharides

     E. Monosaccharides

            1. A monomer of a carbohydrate, referred to as a simple sugar

            2. Contains carbon, hydrogen, and oxygen in a ratio of 1:2:1

            3. General formula is (CH2O)n where N = any whole number
               from 3 to 8

            4. Three most common monosaccharides:

                 a. Glucose (C6H12O6) (main energy source for cells)
                 b. Fructose (In fruits and is the sweetest of the 3)
                 c. Galactose (Found in milk)

            5. Isomers are compounds with a single chemical formula but
               different structural forms

     F. Disaccharides

            1. Two monosaccharides combined in a condensation reaction
               to form a double sugar

            2. Example is Sucrose (fructose + glucose)

     G. Polysaccharides
       1. Made up of 3 or more monosaccharides

       2. Glycogen

            a. Stored form of glucose in animals (100’s of glucose
               molecule chains) (quick energy)

       3. Starch

            a. Stored form of glucose in plants

       4. Cellulose

            a. Another stormed form of glucose in plants

            b. Gives strength and rigidity to plant cells

            c. Hard to break down

H. Proteins (Figure 2-16)

       1. Organic compounds composed of carbon, hydrogen, oxygen,
          and nitrogen

       2. Formed by the linkage of monomers called amino acids

       3. Hair, skin, muscles, and enzymes are made of proteins

I. Amino Acids

       1. 20 different amino acids

       2. Each amino acid contains:

            a.   A single hydrogen atom
            b.   A carboxyl group
            c.   An amino group
            d.   The R group

       3. The main difference between amino acids is the R group
       4. This gives proteins different shapes to carry out different
          activities

J. Dipeptides and Polypeptides

       1. Two amino acids bond to form a dipeptide

       2. It’s called a peptide bond during the condensation reaction

       3. Polypeptides are long chains of amino acids

       4. Proteins are composed of polypeptides

K. Enzymes

       1. RNA or protein molecules that act as biological catalysts

       2. Enzyme reaction depend on a physical fit between the
          enzyme and the specific substrate (section 2-4, figure 2-21)

       3. A substrate is the reactant being catalyzed

       4. The active site is the region of the protein where the
          substrate fits onto the enzyme

       5. The enzyme breaks bonds and reduces the activation energy
          of the reaction

       6. The enzyme releases the products and is unchanged during
          the process, so it can be used many times

       7. Enzyme don’t work if the environment changes

            a. Temperature
            b. pH

L. Lipids (Figure 2-14)

       1. Large nonpolar organic molecules (Don’t dissolve in water)
       2. Examples:

           a.    Triglycerides
           b.    Phospholipids
           c.    Steroids
           d.    Waxes
           e.    Pigments


       3. Have a high ratio of carbon and hydrogen atoms to oxygen
          atoms

       4. Because of this lipids store more energy per gram

M. Fatty Acids

       1. Unbranched carbon chains that make up most lipids

       2. One end of the fatty-acid molecule (carboxyl group) is
          hydrophilic (attracted to water)

       3. The other end (hydrocarbon) is hydrophobic (fears water)

       4. Saturated fatty acids have a carbon atom covalently bonded
          to 4 atoms

       5. Unsaturated fatty acids carbon atoms that are not bonded to
          the maximum number of 4

N. Triglycerides

       1. Composed of 3 molecules of fatty acids joined to one
          molecule of glycerol (an alcohol)

       2. Saturated fats have high melting points and are hard at
          room temperature (butter and fats in red meat)

       3. Unsaturated fats are soft or liquid at room temperature (in
          plant seeds)

O. Phospholipids
       1. Composed of 2 molecules of fatty acids joined to a molecule
          of glycerol

       2. The cell membrane is made of two layers of phospholipids
          called the lipid bilayer

       3. Since lipids don’t dissolve in water, the membrane forms a
          barrier between the inside and outside of the cell

P. Waxes

       1. A structural lipid

       2. Waxes are waterproof (coating on plants)

       3. Waxes form protective layers on animals (earwax)

Q. Steroids

       1. Composed of 4 fused carbon rings with functional groups
          attached to them

       2. Examples are hormones (testosterone) and cholesterol

R. Nucleic Acids (Figure 2-15)

       1. Organic molecules that store and transfer information in
          the cell

       2. The two types of nucleic acids:

              a. DNA (deoxyribonucleic acid)
              b. RNA (ribonucleic acid)

       3. DNA contains information that determines the
          characteristics of organisms and directs its cell activities

       4. RNA stores and transfers information from DNA that is
          used to manufacture proteins
5. Some RNA molecules act as enzymes

6. Nucleotides are the basic building block of nucleic acids

7. They are composed of 3 parts:

    a. Phosphate group
    b. Five-carbon sugar
    c. Nitrogen base

8. DNA and RNA are polymers, composed of 1000’s of
   nucleotides

				
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