Life

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CHAPTER 2

• Lifes Chemical Basis

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Atoms

• Element- fundamental substance consisting of only
  one type of atom. C, N, H, O are the most
  commonly occurring elements in the human body.
• There are 110 different elements, but only 92
  naturally occurring ones. Numbers 93-110 are
  very unstable and degrade quickly.

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You try it - Name the elements

• Sodium
• Nickel
• Cesium
• mercury

• Na
• Ni
• Cs
• Hg

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Vocabulary

• Matter anything that takes up space and has
  mass
• Subatomic particles electron, protons, neutrons

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Levels of Chemical Organization

• Subatomic particles
• Atom
• Molecules

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What Are Atoms?

• Smallest particles that retain properties of an
  element
• Made up of subatomic particles
• Protons ()
• Electrons (-)
• Neutrons (no charge)
• Note protons and neutrons themselves are now
  known to consist of still smaller particles
  called quarks.

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Structure of an Atom

• Nucleus
• Contains protons and neutrons
• Protons- subatomic particle with a small positive
  charge ()
• Neutrons- subatomic particle with no charge (0)

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Structure of an Atom
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Structure of Atoms
electron
proton
neutron
Hydrogen
Helium
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Structure of an Atom

• Shells around the nucleus of an atom contain
  electrons.
• Electron- subatomic particle with a small
  negative charge (-)
• An atom usually
• contains equal
• numbers of protons
• and electrons. When
• it doesnt its called
• an ion.

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How to Read the Periodic Table

• See Appendix IV in back of textbook for good
  periodic table.
• Atomic Number- number of protons in a nucleus of
  an atom. This defines which element the atom is.
  Always the smaller number.

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Atomic Number Protons

• Symbol No
• O 8
• Ca 20
• Na 11
• Cl 17
• K 19
• Fe 26
• N 7

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How to Read a Periodic Table

• Mass Number- total number of protons and neutrons
  in the atomic nucleus. The larger number.
• Mass Number Number of protons Number of
  neutrons
• http//www.dayah.com/periodic

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You Try It!

• Symbol Protons Neutrons Mass
• O 8 8 16
• Ca 20 ? 40
• Na ? 12 23
• Cl 17 ? 35
• K ? 20 39
• Fe 26 30 ?
• N 7 ? 14

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Putting Radioisotopes to work

• Isotope- atoms with different numbers of
  neutrons.
• Isotopes have different mass numbers
• Example
• Carbon 12 has 6 protons, 6 neutrons
• Carbon 14 has 6 protons, 8 neutrons

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Vocabulary

• Radioisotopes an isotope that spontaneously
  emits energy in the form of subatomic particles
  and x-rays when its nucleus disintegrates
• Radioactive decay isotopes that the nucleus is
  disintegrating and transforming one element into
  another spontaneously. (three types)
• Half life the amount of time for ½ of the mass
  of an radioactive isotope to break down

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Example

• Two isotopes of carbon
• Carbon 12 (12C)and Carbon 14 (14C)
• Carbon 14
• Half life -Decay time is 5700 yrs
• If you have 100 grams of Carbon 14
• after 5700 yrs you have 50 grams carbon 14 and
  50 grams Nitrogen 13

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Decay example- The nucleus captures an electron
which basically turns a proton into a neutron.
Here's a diagram of electron capture with
beryllium-7
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Radioisotopes

• Have an unstable nucleus that emits energy and
  particles
• Radioactive decay transforms radioisotope into a
  different element
• Decay occurs at a fixed rate (half-life)
• Carbon-14 becomes N-14 at a specific rate

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Radioisotopes can be used in medicine, ecology,
botany and many other scientific fields.

• A tracer is a molecule in which a radioisotope
  has been substituted for a more stable isotope
  and can then be traced by the energy it
  releases as it decays.

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Popcorn Chemistry

• When will each kernel pop?
• After it pops, can it even go back to being a
  kernel?
• What is released during the popping?

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• Like popcorn, there is no way to know which
  radioisotope will degrade first!
• Also like popcorn, once radioisotopes change,
  they can never return to their previous state.

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• When radioisotopes degrade, energy is released.
• Machines that can see this energy can be used
  to follow the path of the radioisotopes as they
  travel through a system, each one releasing
  energy as they degrade.

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Radioactive tracers can be used in

• PET scans
• Pharmaceutical research
• Metabolic studies
• X-rays

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Concentrations of radioactive tracer bound to
monoamine oxidase B (MAO B). Red shows the
highest concentration. Clearly, lower
concentrations are seen in the smoker. MAO B
helps regulate nerve function and blood pressure.

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When atoms combine with atoms

• Atoms acquire, share, donate electrons.
• Whether one atom will bond with others depends on
  the number and arrangement of its electrons.
• The atoms of some elements do this quite easily
  and other do not.
• When an atom has one or more vacancies in
  orbitals, it interacts with other atoms by
  donating, accepting, or sharing electrons
  (forming chemical bonds).

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Electrons

• Electrons occupy orbital's, or defined volumes of
  space around an atoms nucleus.
• Successive orbital's correspond to levels of
  energy, which become higher with distance from
  the atomic nucleus.
• One or at most two electrons can occupy an
  orbital.
• The atoms with vacancies in orbitals at their
  highest level tend to interact and form bonds
  with one another.

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Vacancy vs. No vacancy
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Electron organization

• First shell-
• Lowest energy
• Holds up to 2
• electrons
• Second shell
• holds up to 8
• electrons

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A model atomic structure is a diagram with
successively larger circles, or shells, that keep
track of all electrons in the orbital at a given
energy level.

• Max number of electrons in each shell
• Shell Number electrons
• 1 2
• 2 8
• 3 18
• 4 32

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electron
proton
neutron
CHLORINE 17p , 17e-
SODIUM 11p , 11e-
CARBON 6p , 6e-
OXYGEN 8p , 8e-
NEON 10p , 10e-
HYDROGEN 1p , 1e-
HELIUM 2p , 2e-
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Electron Vacancies

• Unfilled shells make atoms likely to react
• Hydrogen, carbon, oxygen, and nitrogen all have
  vacancies in their outer shells

NITROGEN 7p , 7e-
CARBON 6p , 6e-
HYDROGEN 1p , 1e-
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You do it - Draw the shell configuration for each
of these

• Na
• Chlorine
• K
• Neon
• Ca

• 2, 8, 1
• 2, 8, 7
• 2, 8, 18, 1
• 2, 8
• 2, 8, 18, 2

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• A B C
• What are these elements?
• Hint Its protons equal its atomic number, and
  since it is an atom so do the number of
  electrons.

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What are these elements? A oxygen B
aluminum C Nitrogen
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Periodic Table

• Each element in a family or group on the Periodic
  Table has common properties
• Examples
• Valence electrons
• Ion formation

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Ions

• Lose electrons
• Valence electrons 1-4
• Na 1 valence 1 ion
• Be 2 valence 2 ion
• B 3 valence 3 ion
• C 4 valence 4 ion
• Gain electrons
• Valence electrons 4-7
• C 4 valence -4 ion
• N 5 valence -3 ion
• O 6 valence -2 ion
• F 7 valence -1 ion

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• When an atom or molecule loses electrons, it
  becomes positively charged.
• For example, when Na loses an electron it becomes
  Na.
• Positively charged ions are called cations.

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• When an atom or molecule gains electrons, it
  becomes negatively charged.
• For example when Cl gains an electron it becomes
  Cl-.
• Negatively charged ions are called anions.
• An atom or molecule can gain or lose more than
  one electron.

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Valence Electrons of 8A- Noble Gases or Inert
Gases

• Element Valance electrons
• He 2
• Ne 8
• Ar 8
• Kr 8

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New Terms

• Chemical bond joining one atom to another by
  electrons joining
• Compound are molecules that consist of two or
  more different elements in proportions combined
  by a chemical bond
• Molecule two or more atoms of the same element
  or different elements joined by a chemical bond

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New Terms

• Formula- the proportional arrangement and short
  hand for a chemical
• Chemical reaction- reacting two or more compounds
  and or molecules with one another
• Reactants the things that are mixed together in
  a chemical reaction
• Products the things that are produced in a
  chemical reaction

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Electron Vacancies

• Unfilled shells make atoms likely to react
• Hydrogen, carbon, oxygen, and nitrogen all have
  vacancies in their outer shells

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Chemical Bonds, Molecules, Compounds

• Bond is union between electron structures of
  atoms
• Atoms bond to form molecules
• Molecules may contain atoms of only one element -
  O2
• Molecules of compounds contain more than one
  element - H2O

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Formulas Chemical Bookkeeping

• Use symbols for elements when writing formulas
• (Need to remember!) Formula for glucose
  is C6H12O6
• 6 carbon
• 12 hydrogen
• 6 oxygen

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Chemical Bookkeeping

• Chemical equation shows reaction
• Reactants ---gt Products
• Equation for photosynthesis

REACTANTS
PRODUCTS
sunlight energy
6CO2
12H2O

C6H12O6
6H2O

6O2

---gt
CARBON DIOXIDE
WATER
WATER
OXYGEN
GLUCOSE
12 hydrogens 6 oxygens
12 oxygens
6 carbons 12 hydrogens 6 oxygens
6 carbons 12 oxygens
24 hydrogens 12 oxygens
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Major Bonds in Biological Molecules

• .
• The bonding behavior of biological molecules
  starts with the number and arrangement of
  electrons in each type of atom.
• Ionic, covalent, and hydrogen bonds are the main
  categories of bonds between atoms in biological
  molecules.

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New Terms

• Ionic Bonds gain or lose control over e-
• Covalent Bonds share e-
• Hydrogen Bonds between H

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Ion Formation

• Atom has equal number of electrons and protons -
  no net charge
• Atom loses electron(s), becomes positively
  charged ion
• Atom gains electron(s), becomes negatively
  charged ion

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Ionic Bonding

• One atom loses electrons, becomes positively
  charged ion
• Another atom gains these electrons, becomes
  negatively charged ion
• Charge difference attracts the two ions to each
  other

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Formation of NaCl

• Sodium atom (Na) (1 ion)
• Outer shell has one electron
• Chlorine atom (Cl) ( -1 ion)
• Outer shell has seven electrons
• Na transfers electron to Cl, forming Na and Cl-
• Ions remain together as NaCl (no charge)

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Ionic Bond

• The electron from sodium atom is lost to the
  chlorine atom which gains the electron.

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Covalent Bonding

• Atoms share a pair or pairs of electrons to fill
  outermost shell
• Example below is a single bond

• Single covalent bond
• Double covalent bond
• Triple covalent bond

Molecular hydrogen
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• Single bonds between H and carbon . Methane
  (lower right) has 4 single bonds formed with
  hydrogen.

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Triple bonds

• Share 6 electrons for three bonds (triple)

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Polar and non-polar covalent bond

• Non Polar share electrons equally
• Polar do not equally share electrons
• Electronegativity is a measure of the tendency of
  an atom to attract a bonding pair of electrons.

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Water Polar Covalent molecule

• Oxygen is much more electronegative than hydrogen
  
• shared electrons spend more time with the oxygen
  part of the molecule than with the hydrogen part
• Unequal sharing of electrons results in the
  oxygen having a partial negative charge and the
  hydrogen atoms having a partial positive charge.

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Water A Polar Covalent Molecule

• Molecule has no net charge
• Oxygen end has a slight negative charge
• Hydrogen end has a slight positive charge

O
H
H
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Covalent Polar or Nonpolar Bonds

• Non polar if atoms share electrons equally
• Hydrogen gas (H - H)
• Polar if electrons spend more time near nucleus
  with most protons
• Water
• Electrons more attracted to O nucleus than to H
  nuclei

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Hydrogen Bonding

• Atom in one polar covalent molecule is attracted
  to oppositely charged atom in another such
  molecule or in same molecule
• Hydrogen bonds are weak.

Water molecule
Ammonia molecule
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Hydrogen Bonding

• Important role in the structure and function of
  biological compounds.
• Found in DNA
• Gives DNA unique properties

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Properties of Water

• Bonds to hydrophilic substances
• Water loving
• Repels hydrophobic ones
• Water hating
• Temperature stabilizing
• Expands floats when it freezes

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Water cont.

• Capacity to dissolve substances
• Universal solvent
• Evaporation heat energy converts liquid water
  to a gaseous state
• Form skin evaporation can help cool body
• Water boils at sea level at 100C or 212F.
• Water freezes at sea level at 0C or 32F.

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Water cont.

• Cohesion has the ability to resist rupturing
  when placed under tension
• Cohesion causes high surface tension
• Shows a capacity to resist rupturing when
  stretched
• Helps to absorb nutrient laden water to grow
• Water rise in tubes
• Some evaporates in leaves
• Pull other water molecules to fill in behind
  those evaporated

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Ice
Liquid water

• This slide shows hydrogen bonding (white dotted
  line)
• Solids are more organized in shape.

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Why Ice Floats

• In ice, hydrogen bonds lock molecules in a
  lattice
• Water molecules in lattice are spaced farther
  apart then those in liquid water
• Ice is less dense than water

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Hydrophilic HydrophobicSubstances

• Hydrophilic substances
• Polar
• Hydrogen bond with water
• Example sugar
• Hydrophobic substances
• Nonpolar
• Repelled by water
• Example oil

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Temperature-Stabilizing Effects

• Liquid water can absorb much heat before its
  temperature rises
• Why?
• Much of the added energy disrupts hydrogen
  bonding rather than increasing the movement of
  molecules

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Evaporation of Water

• Large energy input can cause individual molecules
  of water to break free into air
• As molecules break free, they carry away some
  energy (lower temperature)
• Evaporative water loss is used by mammals to
  lower body temperature

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Water Is a Good Solvent

• Ions and polar molecules dissolve easily in water
  
• When solute dissolves, water molecules cluster
  around its ions or molecules and keep them
  separated

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Water Cohesion

• Hydrogen bonding holds molecules in liquid water
  together
• Creates surface tension
• Allows water to move as continuous column upward
  through stems of plants

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Example of Waters Cohesion
Fig. 2-11a, p.27
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The pH Scale

• Ions dissolved in fluids inside and outside the
  each living cell influence its structure and
  function.
• Among the most influential are hydrogen ions.
• They have far reaching effects largely because
  they are chemically active and because there are
  so many of them.

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The pH Scale

• Measures H concentration of fluid
• Change of 1 on scale means 10X change in H
  concentration
• Highest H Lowest H
• 0---------------------7-------------------14
• Acidic Neutral Basic

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pH Scale

• Water splits into ions
• H is an acid
• OH- is a base
• The greater the number of H concentration the
  lower the pH.
• The greater the number of OH- concentration the
  higher the pH.

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Acid/Base

• Acids
• Donate H
• Below 7

• Bases
• Accept OH-
• Above 7
• Also called alkaline

Both can cause severe damage
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The pH Scale
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Acid Rain

• A coal-burning power plant emits sulfur dioxide,
  which dissolves in water vapor to form acid rain

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Salt Water

• A salt is any substance that dissolves in water
  and releases ions (not H)
• Forms when and acid and a base are mixed.
• HCL NaOH? H2O NaCl
• Hydrochloric acid reacting with sodium hydroxide
  produces water and sodium chloride (salt)
• NaCl (solute) dissolves in water (solvent) forms
  Na and Cl- (ions)

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Buffers Against Shifts in pH

• Cells must maintain homeostasis
• Cells must respond to shifts in pH quickly
• Respiratory acidosis high carbon dioxide levels
  causing pH drop in blood , muscles tetany
  (continued muscle contraction)
• Respiratory alkaosis pH rises in blood
• Blood pH drops can cause a coma if not corrected
• Buffer systems reponds to readjust pH
• Limited in ability

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Carbonic Acid-Bicarbonate Buffer System

• When blood pH rises, carbonic acid dissociates to
  form bicarbonate and H
• H2C03 -----gt HC03- H
• When blood pH drops, bicarbonate binds H to form
  carbonic acid
• HC03- H -----gt H2C03

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Summary

• Ions dissolved in fluids on the inside and
  outside of cells have key roles in cell function.
• Acidic substances release hydrogen ions, and
  basic substances accept them.
• Salts are compounds that release ions other than
  H and OH-.
• Acid-base interactions help maintain pH, which is
  the H concentration in a fluid.
• Buffer systems help maintain the bodys acid-base
  balance at levels suitable for life.

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