The nature of molecules

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Chapter 2

• The nature of molecules

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The nature of atoms

• Matter Any substance that has mass and occupies
  space
• Atoms The smallest unit of an element that
  contains all the characteristics of that element
• Atoms are the building blocks of matter
• Understanding the structure of atoms is critical
  to understanding the nature of biological
  molecules
• Atomic structure developed by Niels Bohr (1913)
• Includes central nucleus and orbiting electrons

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Atomic structure

• Atoms are composed of
• Protons- positively charged particles
• Neutrons- neutral particles
• Electrons- negatively charged particles
• Protons and neutrons are located in the nucleus
• Electrons are found in orbitals outside of the
  nucleus

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Atomic Structure

• Atomic number- number of protons
• Mass umber of atom- number of protons number of
  neutrons
• This information can be given simply in the form
• How many protons and neutrons has this atom got?
• The atomic number counts the number of protons
  (9) the mass number counts protons neutrons
  (19). If there are 9 protons, there must be 10
  neutrons for the total to add up to 19.
• The atomic number is tied to the position of the
  element in the periodic table
• Therefore the number of protons defines what sort
  of element you are talking about
• Examples So if an atom has 8 protons (atomic
  number 8), it must be oxygen. If an atom has 12
  protons (atomic number 12), it must be
  magnesium.

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Basic structure of atoms
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Atomic Structure

• Every different atom has a characteristic number
  of protons in the nucleus.
• Atoms with the same atomic number have the same
  chemical properties and belong to the same
  element.
• An element is any substance that cannot be broken
  down to any other substance by ordinary chemical
  means
• Each proton and neutron has a mass of
  approximately 1 dalton
• Protons and neutrons don't in fact have exactly
  the same mass - neither of them has a mass of
  exactly 1 on the carbon-12 scale (the scale on
  which the relative masses of atoms are measured).
  On the carbon-12 scale, a proton has a mass of
  1.0073, and a neutron a mass of 1.0087.

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Isotopes

• Isotopes are atoms of the same element that have
  different atomic mass numbers due to different
  numbers of neutrons
• Radioactive isotopes-unstable and undergoes
  radioactive decay, releasing energy
• Half-life- time it takes for ½ of the atoms in a
  sample to decay
• Carbon-14 dating and tagging molecules and tract
  progress in both chemical reactions and in living
  tissues

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Atomic Structure

• Neutral atoms have the same number of protons and
  electrons
• Ions are simply charged particles or
• An atom or molecule which has lost or gained one
  or more valence electrons, giving it a positive
  or negative electrical charge.
• Cations are positively charged and have more
  protons than electrons
• Anions are negatively charged and have more
  electrons than protons

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Electrons

• Electrons are located in orbitals surrounding the
  nucleus.
• Orbitals
• Each orbital can contain only 2 electrons.
• Electrons possess potential energy
• Electrons far from the nucleus have the most
  energy

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Electron orbitals
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Oxidation and Reduction

• Oxidation is loss of an electron
• Reduction is gain of an electron
• Electrons can be transferred from one atom to
  another, while still retaining the energy of
  their position in the atom
• Redox reactions when an atom or molecule is
  oxidized while another is reduced.

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Elements found in living systems

• Valence electrons electrons in the outermost
  energy level of an atom
• An elements chemical properties are dependent on
  interactions among valence electrons of different
  atoms.

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Periodic Table

• The periodic table arranges all elements
  according to their atomic number.
• 90 elements occur naturally
• Each with different number of protons and
  different arrangement of electrons
• Mendeleev arranged table
• Pattern of chemical properties in groups of eight
• The periodic table identifies elements with
  similar chemical properties

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Periodic Table
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Elements

• Octet rule Atoms tend to establish completely
  full outer energy levels.
• Atoms with full energy levels are less reactive
  than atoms with unfilled energy levels
• Elements with 7 electrons (1 fewer than the max
  number of 8) are highly reactive
• Examples include F, Cl, Br
• Elements possessing all eight electrons in their
  outer energy level are inert or nonreactive
• For helium (he)- 2 in outer energy level
• He, Ne, Ar etc. are termed the noble gases

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Reactivity of He vs. N
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Question?

• Of the 90 naturally occurring elements, only 12
  are found in living systems in more than trace
  amounts ( or gt to 0.01), can you name this 12
  elements?
• Of this 12 elements, what four make up 96.3 of
  the weight of your body?

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Elements

• 90 naturally occurring elements
• Only 12 elements are found in living organisms in
  substantial amounts
• Carbon, Oxygen, Hydrogen, Nitrogen, Sodium,
  Chlorine, Calcium, Phosphorus, Potassium, Sulfur,
  Iron, Magnesium
• Four elements make up 96.3 of human body weight
• Carbon,, Hydrogen, Oxygen, and Nitrogen (CHON)
• Majority of molecules that make up living things
  are compounds of carbon
• Organic compounds
• Trace elements such as Zinc and Iodine play
  crucial roles in cellular processes
• Iodine deficiency can cause thyroid dieases

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The nature of chemical bonds

• Molecule group of atoms held together in a
  stable association
• Compounds molecules containing more than one
  type of element
• Chemical bonds bonds that hold atoms together in
  molecules or compounds
• Ionic bonds a type of chemical bond that is
  formed by the attraction of oppositely charged
  ions.
• Covalent bonds a type of chemical bond that is
  formed when atoms share two or more valence
  electrons
• The strength of covalent bonds depends on the
  number of electron pairs shared by the atoms
• STRENGTH Triple bonds gt Double bondsgt Single
  bonds

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Salt An example of an ionic bond
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Examples of Covalent Bonds
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Hydrogen bonds

• The molecule formed by the two hydrogen bonds is
  stable for 3 reasons
• It has no net charge
• The octet rule is satisfied
• It has no unpaired electrons

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Electronegativity

• Electronegativity is an atoms affinity for
  electrons
• Increases left to right in a row on the periodic
  table and decreases down the columns of the table
• Elements in upper-right corner have the highest
  electronegativity

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Polar vs. Non-polar covalent bonds

• Differences in electronegativity dictate how
  electrons are distributed in polar bonds
• Nonpolar covalent bonds is equal sharing of
  electrons
• Polar covalent bonds is the unequal sharing of
  electrons

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Chemical bonds

• Chemical reactions involve the formation or
  breaking of chemical bonds.
• Whether a chemical reaction occurs is influenced
  by
• Temperature
• Concentration of reactants and products
• Availability of a catalyst
• A catalyst is a substance that increases the rate
  of a reaction
• Shortens the time needed to reach equilibrium
• Living systems proteins called enzymes act as
  catalysts

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Chemical bonds

• Chemical reactions are written with the reactants
  first, followed by the products
• 6H2O 6CO2 ? C6H12O6 6O2
• reactants products
• Chemical reactions are often reversible.
• C6H12O6 6O2 ? ? 6H2O 6CO2

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Bonds and Interactions- Table 2.1
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Water a vital compound

• ¾ of Earth is water
• All living things are dependent upon water
• 2/3 of any organisms body is composed of water
• Forms of Water
• Solid
• Liquid
• Gas

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Structure of water

• The structure of water is the basis for its
  unique properties.
• Each water molecule is composed of 1 oxygen atom
  and 2 hydrogen atoms
• Oxygen atom shares 1 electron with each hydrogen
  atom
• Water is a polar molecule 2 partial negative
  charges near the O atom and 2 partial positive
  charges on each H atom
• The most important property of water is the
  ability to form hydrogen bonds.

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Polarity of Water

• Forms basis for water chemistry and chemistry of
  life
• Bonds between oxygen and hydrogen are highly
  polar
• Partial electrical charges develop
• O is partially negative
• H is partially positive
• Negative charges occupy more space than positive
  charges causing the O-H bond is compressed
  (109.5?)

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Hydrogen Bonds

• Hydrogen bonds weak attractions between the d- O
  of one water molecule and the d H of a different
  water molecule
• Hydrogen bonds can form between water molecules
  or between water and another charged molecule.

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Water is both cohesive adhesive

• What causes water to be both cohesive and
  adhesive?
• Its polarity
• Cohesion water molecules stick or bind to other
  water molecules by hydrogen bonds
• Adhesion water molecules stick or bind to other
  polar molecules by hydrogen bonding

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Properties of water

• Water is cohesive
• Water has a high specific heat
• A large amount of energy is required to change
  the temperature of water
• Water has a high heat of vaporization
• The evaporation of water from the surface causes
  cooling of that surface
• Solid water is less dense than liquid water
• Bodies of water freeze from the top down
• Water is a good solvent
• Dissolves polar molecules and ions

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Properties of Water

• 6.) Water organizes nonpolar molecules.
• hydrophilic water-loving
• hydrophobic water-fearing
• water causes hydrophobic molecules to aggregate
  or assume specific shapes.
• 7.) Water can form ions.
• H2O ? OH-1 H1
• hydroxide ion hydrogen ion

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Acids and bases

• Hydrogen ion (H1) is the basis of the pH scale.
• Greater H1 concentration?lower pH (acidic)
• Lower H1 concentration?higher pH (basic)
• pH scale measures H1 concentration
• pH-log H1
• A pH value of 7 is neutral
• gt7 is basic
• lt7 is acidic

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pH Scale
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Acids and bases

• Acid a chemical that releases H1 ions
• Base a chemical that accepts H1 ions
• Buffer a chemical that accepts/releases H1 as
  necessary to keep pH constant

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Acids and bases

• Most biological buffers consist of a pair of
  molecules, one an acid and one a base.

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Buffers minimize changes in pH