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The Chemistry of Life

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The Chemistry of Life Hydrogen Bonds The positive hydrogen in polar molecules forms weak bonds with the negative poles of other polar molecules. – PowerPoint PPT presentation

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Title: The Chemistry of Life


1
The Chemistry of Life
2
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)

3
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

6
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.

7
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.

8
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.

9
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.

12
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.

13
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.

14
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.

15
  • 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.

16
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

17
Chemical Reactions
  • Compounds and molecules are created when bonds
    are formed between elements.
  • REACTANTS -------? PRODUCTS
  • Na I2 ? 2 Na I
  • Balancing Equations

18
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.

22
Table Salt Crystal
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25
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.

26
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

27
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.

28
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.

29
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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31
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.

32
H2O
33
Methane CH4
34
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.

35
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.

36
Properties of Water
37
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.

38
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
39
  • 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.

46
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.

49
High heat capacity
  • calorieThe amount of energy required to raise a
    defined amount of a substance by one degree.

50
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

51
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.

52
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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