BSC 1010C
General Biology I
Dr. Graeme Lindbeck
glindbeck@valenciacollege.edu

The Chemical Context of Life

Outline

1. Matter consists of chemical elements in pure form and in combinations called compounds
2. Life requires about 25 chemical elements
3. Atomic structure determines the behavior of an element
   1. Subatomic Particles
   2. Atomic Number and Atomic Weight
   3. Isotopes
   4. Energy Levels
   5. Electron Configuration and Chemical Properties
4. Atoms combine by chemical bonding to form molecules
   1. Covalent Bonds
   2. Ionic Bonds
5. Weak chemical bonds play important roles in the chemistry of life
   1. Hydrogen Bonds
6. A molecule's biological function is related to its shape
7. Chemical reactions change the composition of matter

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I. Matter consists of chemical elements in pure form and in combinations called compounds

Chemistry is fundamental to an understanding of life, because living organisms are made of matter.

Matter = Anything that takes up space and has mass.

Mass = A measure of the amount of matter an object contains.

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II. Life requires about 25 chemical elements

Element = A substance that cannot be broken down into other substances by chemical reactions.

• All matter made of elements.
• 92 naturally occurring.
• Designated by a symbol of one or two letters.

About 25 of the 92 naturally occurring elements are essential to life. Biologically important elements include:

Trace element = Element required by an organism in extremely minute quantities.

• Though in small quantity, are indispensable for life.
• For example: B, Cr, Co, Cu, F, I, Fe, Mn, Mo, Se, Si, Sn, V and Zn.

Elements can exist in combinations called compounds.

Compound = A pure substance composed of two or more elements combined in a fixed ratio.

• For example: NaCl (sodium chloride).
• Has unique emergent properties beyond those of its combined elements. (Na and Cl have very different properties from NaCl.)

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III. Atomic structure determines the behavior of an element

Atom = Smallest possible unit of matter that retains the physical and chemical properties of its element.

• Atoms of the same element share similar chemical properties.
• Made up of subatomic particles.

A. Subatomic Particles

The three most stable subatomic particles are:

1. Neutrons [no charge (neutral)].
2. Protons [+1 electrostatic charge].
3. Electrons [-1 electrostatic charge].

NEUTRON	PROTON	ELECTRON
no charge	+1 charge	-1 charge
found together in a dense core called the nucleus (positively charged because of protons)		orbit around nucleus (held by electrostatic attraction to charged nucleus)
1.009 dalton	1.007 dalton	1/2000 dalton
masses of both are about the same (about 1 dalton)		mass is so small, usually not used to calculate atomic mass

If an atom is electrically neutral, the number of protons equals the number of electrons, which yields an electrostatically balanced charge.

B. Atomic Number and Atomic Weight

Atomic number = Number of protons in an atom of a particular element.

• All atoms of an element have the same atomic number.
• Written as a subscript to the left of the element's symbol (e.g. ₁₁Na).
• In a neutral atom, # protons = # electrons.

Mass number = Number of protons and neutrons in an atom.

• Written as a superscript to left of an element's symbol (e.g. ²³Na).
• Is approximate mass of the whole atom, since the mass of a proton and the mass of a neutron are both about 1 dalton.
• Can deduce the number of neutrons by subtracting atomic number from mass number.
• Number of neutrons can vary in an element, but number of protons is constant.
• Is not the same as an element's atomic weight, which is the weighted mean of the masses of an element's constituent isotopes.

C. Isotopes

Isotopes = Atoms of an element that have the same atomic number but different mass number.

• Have the same number of protons, but different number of neutrons.
• Under natural conditions, elements occur as mixtures of isotopes.
• Different isotopes of the same element react chemically in same way.
• Some isotopes are radioactive.

Radioactive isotope = Unstable isotope in which the nucleus spontaneously decays emitting sub-atomic particles and/or energy as radioactivity.

• Loss of nuclear particles may transform one element to another
• Has a fixed half life,

Half life = Time for 50% of radioactive atoms in a sample to decay.

Biological applications of radioactive isotopes include:

1. Dating Geological strata and fossils.
   • Radioactive decay is at a fixed rate.
   • By comparing the ratio of radioactive and stable isotopes in a fossil with the ratio of isotopes in living organisms, one can estimate the age of a fossil.
   • The ratio of ¹⁴C to ¹²C is frequently used to date fossils less than 50,000 years old.

2. Radioactive tracers
   • Chemicals labeled with radioactive isotopes are used to trace the steps of a biochemicalreaction or to determine the location of a particular substance within an organism.
   • Radioactive isotopes are useful as biochemical tracers because they chemically react likethe stable isotopes and are easily detected at low concentrations.
   • Isotopes of P, N and H were used to determine DNA structure.
   • Used to diagnose disease (e.g. PET scanner).
   • Because radioactivity can damage cell molecules, radioactive isotopes can also be hazardous.

3. Treatment of cancer
   • e.g. radioactive cobalt.

D. Energy Levels

Electrons = Light negatively charged particles that orbit around nucleus.

• Equal in mass and charge.
• Are the only stable subatomic particles directly involved in chemical reactions.
• Have potential energy because of their position relative to the positively charged nucleus.

Energy = Ability to do work.

Potential energy = Energy that matter stores because of its position or location.

• There is a natural tendency for matter to move to the lowest state of potential energy.
• Potential energy of electrons is not infinitely divisible, but exists only in discrete amounts called quanta.
• Different fixed potential energy states for electrons are called energy levels or electron shells.
• Electrons with lowest potential energy are in energy levels closest to the nucleus.
• Electrons with greater energy are in energy levels further from nucleus.

Electrons may move from one energy level to another. In the process, they gain or lose energy equal to the difference in potential energy between the old and new energy level.

E. Electron Configuration and Chemical Properties

An atom's electron configuration determines its chemical behavior.

Electron configuration = Distribution of electrons in an atom's electron shells.

The first 18 elements of a periodic chart are arranged sequentially by atomic number into 3 rows (periods). In reference to these representative elements, note the following:

• Outermost shell of these atoms never have more than 4 orbitals (one s and three p) or 8 electrons.
• Electrons must first occupy lower electron shells before the higher shells can be occupied. (A reflection of the natural tendency for matter to move to the lowest possible state of potential energy - the most stable state.)
• Electrons are added to each of the p orbitals singly, before they can be paired.
• If an atom does not have enough electrons to fill all shells, the outer shell will be the only one partially filled.

Chemical properties of an atom depend upon the number of valence electrons.

Valence electrons = Electrons in the outermost energy shell (valence shell).

Octet rule = Rule that a valence shell is complete when it contains 8 electrons (except H and He).

• An atom with a complete valence shell is unreactive or inert.
• Noble elements (e.g. helium, argon and neon) have filled outer shells in their elemental state and are thus inert.
• An atom with an incomplete valence shell is chemically reactive (tends to form chemical bonds until it has 8 electrons to fill the valence shell).
• Atoms with the same number of valence electrons show similar chemical behavior.

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IV. Atoms combine by chemical bonding to form molecules

Atoms with incomplete valence shells tend to fill those shells by interacting with other atoms. These interactions of electrons among atoms may allow atoms to form chemical bonds.

Chemical bonds = Attractions that hold molecules together.

Molecules = Two or more atoms held together by chemical bonds.

A. Covalent Bonds

Covalent bond = Chemical bond between atoms formed by sharing a pair of valence electrons.

• Strong chemical bond.
• For example: molecular hydrogen (H₂) - when 2 hydrogen atoms come close enough for their electron shells to overlap, they share electrons - thus completing the valence shell of each atom.

Structural formula = Formula which represents the atoms and bonding within a molecule (e.g. H-H). The line represents a shared pair of electrons.

Molecular formula = Formula which indicates the number and type of atoms (e.g. H₂)

Single covalent bond = Bond between atoms formed by sharing a single pair of valence electrons.

• Atoms may freely rotate around the axis of the bond.

Double covalent bond = Bond formed when atoms share two pairs of valence electrons (e.g. O₂)

Triple covalent bond = Bond formed when atoms share three pairs of valence electrons (e.g. N₂).

NOTE: Double and triple covalent bonds are rigid and do not allow rotation.

Valence = Bonding capacitoy of an atom which is the number of covalent bonds that must be formed to complete the outer electron shell.

• Valences of some common elements: hydrogen = 1, oxygen = 2, nitrogen carbon = 4, phosphorus = 3, sulfur = 2.

Compound = A pure substance composed of two or more elements combined in a fixed ratio.

• For example: water (H₂O), methane (CH₄)
• Note that two hydrogens are necessary to complete the valence shell of oxygen in water, and four hydrogens are necessary for carbon to complete the valence shell in methane.

B. Nonpolar and Polar Covalent Bonds

Electronegativity = Atom's ability to attract and hold electrons.

• The more electronegative an atom, the more strongly it attracts shared electrons.
• Scale determined by Linus Pauling:
  • O = 3.5
  • N = 3.0
  • S and C = 2.5
  • P and H = 2.1

Nonpolar covalent bond = Covalent bond formed by an equal sharing of electrons between atoms.

• Occurs when electronegativity of both atoms is about the same (e.g. CH₄)
• Molecules made of one element usually have nonpolar covalent bonds (e.g. H₂, O₂, Cl₂, N₂)

Polar covalent bond = Covalent bond formed by an unequal sharing of electrons between atoms.

• Occurs when the atoms involved have different electronegativities.
• Shared electrons spend more time around the more electronegative atom.
• In H₂O, for example, the oxygen is strongly electronegative, so negatively charged electrons spend more time around the oxygen than the hydrogens. This causes the oxygen atom to have a slight negative charge and the hydrogens to have a slight positive charge.

Ion = Charged atom or molecule.

Anion = An atom that has gained one or more electrons from another atom and has become negatively charged; a negatively charged ion.

Cation = An atom that has lost one or more electrons and has become positively charged; a positively charged ion.

Ionic bond = Bond formed by the electrostatic attraction after the complete transfer of an electron from a donor atom to an acceptor.

• The acceptor atom attracts the electrons because it is much more electronegative than the donor atom.
• Are strong bonds in crystals, but are fragile bonds in water; salt crystals will readily dissolve in water and dissociate into ions.
• Ionic compounds are called salts (e.g. NaCl or table salt).

NOTE: The difference in electronegativity between interacting atoms determines if electrons are shared equally (non-polar covalent), shared unequally (polar covalent), gained or lost (ionic bond). Non-polar covalent bonds and ionic bonds are two extremes of a continuum from interacting atoms with similar electronegativities to interacting atoms with very different electronegativities.

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V. Weak chemical bonds play important roles in the chemistry of life

Biologically important weak bonds:

• Include hydrogen bonds, ionic bonds in aqueous solutions and other weak forces such as Van der Waals and hydrophobic interactions.
• Make chemical signaling possible in living organisms, because they are only temporary in associations. Signal molecules can briefly and reversibly bind to receptor molecules on a cell, causing a short-lived response.
• Can form between molecules or between different parts of a single large molecule.
• Help stabilize the three-dimensional shape of large molecules (e.g. DNA and proteins).

Hydrogen bond = Bond formed by the charge attraction when a hydrogen atom covalently bonded to one electronegative atom is attracted to another electronegative atom.
• Weak attractive force that is about 20 times easier to break than a covalent bond.
• Is a charge attraction between oppositely charged portions of polar molecules.
• Can occur between a hydrogen that has slight positive charge when covalently bonded to an atom with high electronegativity

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VI. A molecule's biological function is related to its shape

The function of many molecules depends upon their shape.

• Molecules with only two atoms are linear.
• Molecules with more than two atoms have more complex shapes.
• When an atom forms covalent bonds, electrons in the valence shells rearrange into the most stable configuration

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VII. Chemical reactions change the composition of matter

Chemical reactions = process of making and breaking chemical bonds leading, to changes in the composition of matter.

• Process where reactants undergo changes into products.
• Matter is conserved, so all reactant atoms are only rearranged to form products.
• Some reactions go to completion (all reactants converted to products), but most reactions are reversible
• The relative concentration of reactants and products affects reaction rate (tile higher the concentration, the greater probability of reaction).

Chemical equilibrium = Equilibrium established when the rate of forward reaction equals the rate of the reverse reaction.

• Is a dynamic equilibrium with reactions continuing in both directions.
• Relative concentrations of reactants and products stay the same.

Course Pages maintained by
Dr. Graeme Lindbeck.