Chemical Foundations for Cells

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Chapter 2

• Chemical Foundations for Cells

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Chapter Outline

• Review of elements and atomic structure
• Radioactive elements and health/medicine
• Chemical bonding
• Ionic
• Covalent nonpolar and polar
• Hydrogen bonding
• Properties of water
• Acids, bases, and buffers
• Chemical change

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Elements (2.1, 2.3)

• Living organisms are composed of matter
• Matter is composed of atoms of elements
• Element
• substance that cannot be broken down into other
  substances by chemical means

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Elements (2.1)

• 92 naturally occurring elements
• Life requires 25 of these
• 96 of human body is made up of
• Carbon (C)
• Hydrogen (H)
• Oxygen (O)
• Nitrogen (N)

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Compounds (2.1)

• Atoms of one element can join with atoms of other
  elements to form compounds.
• A given compound is always made of the same
  elements combined in the same ways.
• NaCl table salt
• H2O - water
• C6H12O6 - glucose

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Compounds of Life

• Only living organisms have the ability to make
  the compounds of life
• Carbohydrates C, H, O
• Lipids C, H, O
• Proteins C, H, O, N, S
• Nucleic acids C, H, O, N, P

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Atoms (2.3)

• An atom is the smallest unit of an element
• Atoms are composed of 3 subatomic particles
• Protons
• Neutrons
• Electrons

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Subatomic Particles
Subatomic Particle Charge Mass, amu Location in atom
Electron (e-) -1 0 amu Outside of nucleus
Proton (p) 1 1 amu Nucleus
Neutron (n) 0 1 amu Nucleus
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Subatomic Particles and the Elements

• Each element has a unique number of protons.
• Number of protons defines the element.
• Atomic number - number of protons in an atom
• Also indicates the number of electrons in the
  atomWhy???

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Terms

• Mass number sum of the protons and neutrons in
  an atoms nucleus
• Atoms of a given element may differ in the number
  of neutrons in the nucleus called isotopes

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Isotopes

• Isotopes of carbon
• 12C carbon-12 6 neutrons
• 13C carbon-13 7 neutrons
• 14C carbon-14 8 neutrons
• All contain ____ protons and electrons.

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Isotopes and Radioactivity

• Radioactive isotopes have an unstable nucleus
• Emit energy and particles from the nucleus in an
  effort to become more stable
• In doing so they may change the number of protons
  in the nucleus and become a different element.

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Radioactive Isotopes

• Possible to target the energy and detect the
  radioactivity
• Radioactive (RA) isotopes are used in medicine to
  treat cancer and diagnose disease
• Radiation therapy - treatment
• PET - diagnosis

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Fig. 2-4a, p.21
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Fig. 2-4bcd, p.21
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Radioactive Isotopes

• Overexposure to RA isotopes is HARMFUL.
• Energy emitted damages cells radiation therapy
  takes advantage of this
• Can cause mutations that lead to cancers

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Electron Arrangement (2.5)

• When compounds form, the electrons of the bonding
  atoms interact in attempt to obtain a more stable
  state.
• Some electron arrangements are more stable than
  others.see board

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Chemical Bonding (2.6-2.7)

• Chemical bonding atoms gain, lose, or share
  electron(s) to obtain a stable number of electrons

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Chemical Bonding

• Types of bonds
• Ionic bonds transfer of electrons to form ions,
  ions are attracted to each other by ionic bonds
• Covalent bonds bonding atoms share electrons
• Nonpolar covalent bonds equal sharing of
  eleactrons
• Polar covalent bonds unequal sharing of
  electrons

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Chemical Bonding

• Ionic Bond strong attractive force between
  oppositely charged ions
• Atoms form ions by losing or gaining enough
  electron(s) to obtain a stable of electrons in
  their outer shell

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Ionic Bonding
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Chemical Bonding

• Covalent Bond bonded atoms share pair(s) of
  electrons and form molecules.
• Occurs between nonmetals such as C, O, H, N,
  P, S
• Covalent bonding occurs in
• H2
• O2
• H2O

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Two Classes of Covalent Bonds

• Nonpolar Covalent Bond bonded atoms share
  electrons equally
• Occurs between like atoms or between atoms with a
  similar ability to attract shared electrons
• Polar Covalent Bond unequal sharing of
  electrons by the bonded atoms
• Occurs between atoms with very different ability
  to attract shared electrons

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Two hydrogen atoms, each with one proton, share
two electrons in a single nonpolar covalent bond.
molecular hydrogen (H2) HH
Fig. 2-8b(1), p.25
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water (H2O) HOH
Two oxygen atoms share four electrons in a
nonpolar double covalent bond.
molecular oxygen (O2) OO
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Types of Covalent Bonds

• Nonpolar covalent bonded atoms share the
  electrons equally
• Examples of nonpolar bonds
• H2

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• Atoms with different electronegativity values
  form polar covalent bonds.
• Electronegativity (EN) measure of an atoms
  ability to attract shared electrons in a covalent
  bond
• Oxygen and nitrogen have fairly large EN values
  often d -
• Carbon and hydrogen have low EN values often d

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• H2O
• Oxygen needs 2 electrons in its outer shell.
• Oxygen forms 2 polar covalent bonds with
  hydrogen atoms to obtain 8 electrons in the outer
  shell

water (H2O) HOH
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• Polar Covalent unequal pull on shared electrons
  by the bonded atoms
• Results in partial charges on the bonded atoms

d -
O
H
H
d
d
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Common Polar Covalent Bonds

• O-H N-H
• C-O CO
• Label the polarity in each bond.

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Forces between Molecules

• Molecules are weakly attracted to each other by
  intermolecular (IM) forces,
• The most important IM force in biology is the
  hydrogen bond (2.8)
• Attractive force between d H and d O, N or F

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• Hydrogen bond is a weak attractive force
  between a d hydrogen and a d- O, N, or F in a
  second polar bond

Water is a polar molecule capable of hydrogen
bonding.
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Properties of Water

• Water is cohesive and has high surface tension.
• Cohesion ability of molecules to stick together
• Surface tension - ability to resist rupturing
  when under tension

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Properties of Water

• Water resists changes in temperature.
• When heat is applied to an aqueous solution much
  of the heat (energy) is used to break hydrogen
  bonds, not to increase the movement of the
  molecules.

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Properties of Water

• Solid water (ice) is less dense than liquid water
• Ice floats
• Therefore, ice forms on the top of lakes and
  insulates the liquid water below.

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• Water is a good solvent for ionic compounds and
  small polar molecules.
• Water hydrates ions
• Water H bonds to polar
• molecules

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Related Terms

• Hydrophilic
• Water loving
• Capable of hydrogen bonding to water (polar)
• Hydrophobic
• Water fearing
• Cannot hydrogen bond to water (nonpolar)

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Acids, Base, and Buffers (2.14)

• Many ions are dissolved in the fluids in/outside
  of cells called electrolytes
• Na, Ca2, K
• H
• Level of each ion is critical
• Our focus is on H (hydrogen ions)

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Acids, Base, and Buffers

• Acid Substance that produces H when dissolved
  in water.
• Examples
• Hydrochloric acid stomach acid
• Lactic acid made when cells run out of oxygen
• Amino acids building blocks of proteins

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Acids, Base, and Buffers

• Base substance that accepts H1 (hydrogen ions)
  in water
• Examples
• Sodium hydroxide - NaOH
• Most nitrogen containing compounds
• Ammonia NH3
• Urea in urine
• Amino acids building blocks for proteins

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Acids, Base, and Buffers

• Classify substances as acid, base or neutral by
  their pH
• Acids pH lt 7
• Base pH gt 7
• Neutral pH 7
• Pure water has a pH of 7
• See page 28

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Acids, Base, and Buffers

• How the pH scale works
• The lower the pH the more acidic
• The higher the pH the more basic (alkaline)
• A difference of 1 pH unit is a 10-fold difference
  in acidity or alkalinity

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Why is pH important?

• Most cells require a pH near 7. Above or below
  this pH for too long and they die.
• Proteins function only at specific pHs.
• In lab you will determine the optimal pH for a
  protein that is needed to breakdown hydrogen
  peroxide in cells

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Acids, Base, and Buffers

• Buffers solution that resists changes in pH
  even when acid or base is added
• Buffers can both produce H and neutralize H
• Buffers are key to maintaining pH homeostasis
• Most body solutions are buffered

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Why is pH important?

• Blood has a pH of 7.3 7.4
• If the pH is above or below this range for more
  than a couple of days death occurs.
• The blood buffer system helps keep blood pH in a
  range that supports life.