The Chemical Basis of Life

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The Chemical Basis of Life

• Biological function starts
• at the chemical level...

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Objectives Section 1

• Define element, atom, compound, and molecule.
• Draw a model of the structure of an atom.
• Explain what determines an atoms stability.
• Contrast ionic, covalent, and hydrogen bonding.

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Lets review

• Matter - anything that occupies space and has
  mass - makes up everything in the universe.
• Mass - the measurement of the amount of matter in
  an object.
• Mass and weight are NOT the same!
• Weight is the pull of gravity on an objects
  mass.

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4 states of matter solid, liquid, gas and plasma

• Solid anything that has both a definite volume
  and a definite shape.
• Liquid anything that has definite volume but no
  definite shape.(Liquids can flow and can be
  poured.)
• Gas has no definite volume and no definite
  shape, it takes the shape and volume of the
  container into which it is placed.

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• Plasma has no definite volume or shape and is
  composed of electrical charged particles.

• Plasmas are conductive assemblies of charged
  particles, neutrals and fields that exhibit
  collective effects. Further, plasmas carry
  electrical currents and generate magnetic fields.
  Plasmas are the most common form of matter,
  comprising more than 99 of the visible universe

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Color is determined by the gas that emits the
photon.
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Northern Lights a.k.a. Aurora Borealis

• What causes them?
• Northern lights originate from our sun. During
  large explosions and flares, huge quantities of
  solar particles are thrown out of the sun and
  into deep space. These plasma clouds travel
  through space with speeds varying from 300 to
  1000 kilometers per second.
• Takes 2-3 days to reach Earth.

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Element

• Chemical element - a substance that cannot be
  broken down into other substances by ordinary
  chemical means (92 naturally occurring elements
  and about 12 man made) ex. Carbon

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• Carbon, hydrogen, oxygen, and nitrogen make up
  the bulk of living matter, but there are other
  elements necessary for life

Table 2.2
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• Goiters are caused by iodine deficiency

Figure 2.2
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The Periodic Table

• A chart with all of the known elements arranged
  providing information for each. Ex. Atomic
  number, atomic mass number and chemical symbol

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Atoms

• The simplest particle of an element that retains
  all of the properties of the element.
• Composed of 3 kinds of particles, protons (p)
  and neutrons (n0) found in the atoms nucleus and
  electrons (e-) orbiting the nucleus.
• Most of the mass of an atom is
  concentrated in its
  nucleus.

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The Nucleus Protons Neutrons

• Protons ALL atoms of a given element have the
  same number of protons.
• This number is called the Atomic Number.
• Protons have a positive charge.
• The number of protons defines the element.

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• Neutrons Determined by subtracting the atomic
  number from the atomic mass number.
• Zero electrical charge.
• The number can vary between atoms of the same
  element, creating isotopes

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• ISOTOPES

• The number of neutrons may vary
• Variant forms of an element (due to a different
  number of neutrons) are called isotopes
• Some isotopes are radioactive

Table 2.4
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Radioactive isotopes can help or harm us

• Radioactive isotopes can be useful tracers for
  studying biological processes
• PET scanners use radioactive isotopes to create
  anatomical images

Figure 2.5B
Figure 2.5A
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Outside the nucleus...

• Electrons high energy particles with very
  little mass.
• Have a negative charge.
• Move about the nucleus at very high speeds in one
  of seven different energy (E) levels.
• Electrons in outer E levels have more E than
  those is inner levels.
• Each E level (shell) can hold only a certain
  number of electrons.
• 1st level 2 electrons
• 2-7 level up to 8 electron
• A stable atom is an atom that has a full outer
  level (The Octet Rule)

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• Atoms whose shells are not full tend to interact
  with other atoms and gain, lose, or share
  electrons

Outermost electron shell (can hold 8 electrons)
Electron
First electron shell (can hold 2 electrons)
HYDROGEN (H) Atomic number 1
CARBON (C) Atomic number 6
NITROGEN (N) Atomic number 7
OXYGEN (O) Atomic number 8
Figure 2.6
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Important concepts

• The identity of an element is determined by its
  number of protons.
• Atoms are electrically neutral. This means that
  the number of protons an element has is also the
  number of electrons the element has.
• Most atoms Do Not have a full outer level, this
  drives chemical reactions.

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• Compound - a substance containing two or more
  elements in a fixed ratio.
• Chemical formulas shows the kind and proportion
  of atoms of each element that forms a particular
  compound.
• Ex. Table salt NaCl, Water H2O

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• Since most atoms are not stable in their natural
  state and tend to undergo chemical reactions,
  combining in ways that cause their atoms to
  become stable.
• In chemical reactions, chemical bonds are broken,
  atoms are rearranged, and new chemical bonds,
  or attachments,
  are formed.

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Covalent Bond

• Forms when two or more atoms share one or more
  pairs of electrons.
• Results in the formation of molecules, simplest
  part of a substance that retains all the
  properties of the substance.
• The strongest of the bonds.

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• Molecules can be represented in many ways

Table 2.8
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Ionic Bond

• Forms when there is a transfer of electrons
  creating charged atoms (ions) which are then
  attracted to each other due to their opposite
  charges.
• 2nd strongest bond
• Ex. Table salt (NaCl)

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IONS

• When atoms gain or lose electrons, charged atoms
  called ions are created
• cation loses e- and has an overall positive
  charge
• anion gains e- and has an overall negative
  charge

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• Sodium and chloride ions bond to form sodium
  chloride, common table salt

Na
Cl
Figure 2.7B
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Energy (The big E!)

• Energy (E) is the ability to do work or cause
  change.
• Types of E chemical, thermal, electrical, and
  mechanical
• As E flows through an organism, it is often
  converted from one form to another. Ex.
  Chemical E found in food converts into thermal E
  (heat) and mechanical E (move bone).
• All chemical reactions involve E.

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Exergonic Reactions

• Chemical reactions which involve a net release of
  free E.
• Ex. Cellular respiration - the chemical reaction
  in which sugars are broken down to CO2 and H2O,
  releasing E.
• Requires an initial input of E to get the
  reaction started, called activation E.

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Endergonic Reactions

• Reactions that involve a net absorption of free
  energy.
• Ex. photosynthesis - the chemical reaction in
  which CO2 and H2O are combined to make sugars,
  storing E E.
• Requires an initial input of E to get the
  reaction started, called activation E.

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• In many chemical reactions, the amount of energy
  needed to start the reaction, called ACTIVATION
  ENERGY, is high.
• Certain chemical substances, known as CATALYSTS,
  reduce the amount of activation energy needed. A
  reaction in the presence of a catalyst will
  proceed spontaneously or with the addition of a
  small amount of E.
• ENZYMES are an important class of catalysts
  in living things. A single organism may
  have thousands of different enzymes.

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Reduction-Oxidation Reactions

• Many of the chemical reactions that help transfer
  E in living things involve the transfer of
  electrons.
• These reactions in which e- are transferred
  between atoms are known as REDUCTION-OXIDATION
  REACTIONS, OR REDOX REACTIONS. Ex. The formation
  of table salt
• LEO the lion goes GER
• LEO losing electrons is oxidation
• GER gaining electrons is reduction

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• In an OXIDATION REACTION, a reactant LOSES One or
  More e-, becoming more POSITIVE in charge. The
  Sodium atom undergoes oxidation to form Na ion.
• 4. In a REDUCTION REACTION, a reactant GAINS One
  or more E-, Becoming more NEGATIVE in charge. The
  Chlorine atom undergoes Reduction to form Cl-
  ion.
• 5. REDOX REACTIONS ALWAYS OCCUR TOGETHER. An
  Oxidation Reaction occurs, and the e- given up by
  one substance is then accepted by another
  substance in a Reduction Reaction.

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Solutions

• Types of mixtures in which one or more substances
  are uniformly distributed in another substance.
• Can be solids (brass), liquids (sugar water), and
  gases (air).

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• A saturated solution is one in which no more
  solute can dissolve.
• An aqueous (aq) solution is one in which water is
  the solvent. Ex. Nutrients in moist soil, fluid
  in and around body cells

• Solute - the substance dissolved in the solution.
• Solvent - the substance in which the solute is
  dissolved.
• Water is the universal solvent (Why?)
• Concentration - the measurement of the amount of
  solute dissolved in a fixed amount of the
  solution.

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Acids and Bases

• A solution that has an excess of H ions. It
  comes from the Latin word acidus that means
  "sharp".
• Base A solution that has an excess of OH- ions.
  Another word for base is alkali.

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Acids and Bases

• The pH scale measures the relative concentration
  of H ions to OH- ions.
• Acids pH from 0-6.9
• The lower the pH the stronger the acid
• Neutral pH of 7
• Base pH of 7.1-14
• The high the pH the stronger the base

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• The control of pH is important for living things.
  
• Enzymes can function only within a very narrow pH
  range.
• The control of pH in organisms is often
  accomplished with buffers.
• Buffers are chemical substances that neutralize
  small amounts of either an acid or a base added
  to a solution, preventing rapid/large swings in
  pH.

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• Some of your body's fluids, such as stomach acid
  and urine, are acidic. Others, such as intestinal
  fluid and blood are basic.
• Complex buffering systems maintain the pH values
  of your body's many fluids at normal or safe
  levels.

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pH of cytoplasm 7.2 (resting cell)
But the pH of cell sap may be alkaline (9 in
Begonia) or acidic (2.4 in Lemon).

• Many biochemical pathways are affected by the
  degree of acidity in the nucleus or cytoplasm in
  which they occur. In order for enzymatic
  reactions to proceed with the highest efficiency,
  the pH of the medium must be controlled
  (homeostasis!)

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This diagram shows the diffusion directions for
H, CO2, and O2 between the blood and the muscle
cells during exercise. The resulting
concentration changes affect the buffer
equilibria, shown in the upper right-hand corner
of the diagram (yellow).
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The following steps outline the processes that
affect the buffers in the blood during exercise.

• Hemoglobin carries O2 from the lungs to the
  muscles through the blood.
• The muscles need more O2 than normal, because
  their metabolic activity is increased during
  exercise. The amount of oxygen in the muscle is
  therefore depleted in the muscles, setting up a
  concentration gradient between the muscle cells
  and the blood in the capillaries. Oxygen diffuses
  from the blood to the muscles, via this
  concentration gradient.
• The muscles produce CO2 and H as a result of
  increased metabolism, setting up concentration
  gradients in the opposite direction from the O2
  gradient.
• The CO2 and H flow from the
  muscles to the blood, via these
  concentration gradients.

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Blood Buffers Cont

• The buffering action of hemoglobin picks up the
  extra H and CO2.
• If the amounts of H and CO2 exceed the capacity
  of hemoglobin, they affect the carbonic acid
  equilibrium. As a result, the pH of the blood is
  lowered, causing acidosis.

• Hence, the body has developed finely-tuned
  chemical processes (based on buffering and
  acid-base equilibria) that work in combination to
  handle the changes that exercise produces.
• The lungs and kidneys respond to pH changes by
  removing CO2, HCO3 -, and H from the blood.

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