Chapter 2 THE CHEMICAL CONTEXT OF LIFE
Summary of Chapter 2, BIOLOGY, 10TH ED Campbell, by J.B. Reece et al. 2014.
The physical and chemical principles that govern the universe also govern the composition and metabolic processes of living organisms.
Chemistry is the study of matter. It is fundamental to an understanding of life.
Resource: http://www.chem1.com/acad/webtext/virtualtextbook.html
CHEMICAL ELEMENTS AND COMPOUNDS
Organisms are made of matter.
Matter is anything that takes up space (volume) and has mass.
Mass is different from weight although it is often used interchangeably.
· Mass is the amount of matter an object has.
· Weight is the pull of gravity on the mass of an object.
Matter consists of chemical elements in pure form and in combinations called compounds.
Elements are the simplest substances. They cannot be broken down into simpler substances by chemical reactions.
· All matter is made of elements.
· There are 92 naturally occurring elements.
· Elements are designated by a symbol of one or two letters, e.g. C for carbon; Na for sodium.
A compound is a substance consisting of two or more elements combined in a fixed ratio.
Humans need 25 elements that are essential to life. Plants need only 17. These called essential elements of life.
· C, O, H and N make about 96% of living matter.
· Ca, P, K, S, Na, Cl and Mg make about 4% of the organism's weight.
· Trace elements are required in extremely small amounts, e.g. B, Cu, Mn, Mo, Se, etc.
THE STRUCTURE OF ATOMS AND MOLECULES
An element’s properties depend on the structure of its atoms
An atom is the simplest portion of an element that retains its chemical properties.
The properties of an element depend on the atomic structure of its atoms.
Each element has its own characteristic atom represented by a chemical symbol.
Subatomic particles: protons, neutrons and electrons.
Protons and neutrons are packed together in the nucleus of the atom.
Atomic number
The number of protons, atomic number, identifies the atom.
Protons carry a positive electrical charge; neutrons are electrically neutral, and electrons are negative.
The number of protons and neutrons determines the mass of the atom, atomic mass.
The masses of a proton and a neutron are almost identical, about 1.7 x 10-24 grams.
For practical reasons the mass of the atom is expressed in daltons also known as atomic mass units, amu. One dalton equals 1.7 x 10-24 grams.
· Example: the atomic mass of Na is commonly said to be 23 daltons, but more precisely is 22.9898 daltons.
The symbol 2He means that the atom of helium has two protons in its nucleus: its atomic number.
The superscript 4 means the mass number (atomic mass) of the atom, e.g. 42He
The subscript is the atomic number, e.g. the 2 in 42He.
The mass of the electron is about 1/1800 of the mass of a proton or neutron or about 9.109 x 10-31 grams, and it is disregarded in calculating the atomic mass (mass number) of an atom.
Protons plus neutron make the mass number (atomic mass, atomic weight) of an atom.
Isotopes
Isotopes of an element are atoms that have the same number of protons and different number of neutrons, e.g. 12C and 13C are isotopes of carbon; both have 6 protons in the nucleus.
A radioactive isotope is one in which the nucleus decays spontaneously, giving off particles and energy. 14C is a radioactive isotope of carbon.
If the isotope gives off a proton, its atomic number changes and the atom becomes that of another element. E. g. radioactive carbon becomes nitrogen.
Energy levels of electrons
Energy is the ability to do work and cause change.
Potential energy is the energy matter has due to its location or structure, e.g. water in a tank on top of a hill.
Matter has the tendency to move to the lowest possible state of potential energy.
The electrons of an atom have potential energy because of how they are arranged in relation to the nucleus.
Electrons are attracted by the positive nucleus. It takes energy to move electrons farther away from the nucleus.
Electrons have fixed amounts of potential energy that correspond to a position at a distance from the nucleus.
Electrons move around the nucleus of the atom in areas of space called orbitals.
For an electron to occupy an orbital it must have a specific amount of energy.
Orbitals, therefore, correspond to certain amount of energy in the electron, which is called energy level. Energy levels are also referred to as electron shells.
The distribution of electrons in energy levels or shells around the nucleus is called the electron configuration of the atom.
Electrons that are in an energy level distant from the nucleus have more energy than those closer to the nucleus.
The chemical behavior of an atom depends mostly on the number of electrons in its outermost shell.
Electrons in the outermost shell have the greatest amount of energy and are called valence electrons. They occupy the valence shell. See Fig. 2.9.
Elements with the same number of electrons in their valence shell have similar chemical properties, e.g. K and Na; Cl and F.
Electrons can change from one energy level to another unoccupied level by gaining or loosing energy.
The reactivity of atoms arises from the presence of unpaired electrons in one or more orbitals of their valence shells.
An orbital is a three-dimensional space where the electron is found 90% of the time.
An orbital contains a maximum of two electrons.
Electrons in orbitals with similar energies occupy the same principal energy level.
The orbitals of an energy level are designated by the letters s, p, d and f. See fig. 2.10.
Periodic table of elements: http://www.ktf-split.hr/periodni/en/
ATOMS COMBINE TO FORM MOLECULES
Atoms with incomplete valence shells interact with certain other atoms in such a way that each partner completes its valence shell.
Atoms do this by either sharing or transferring valence electrons.
Atoms may combine chemically, bond, to form molecules.
· Molecules of an element have atoms of the same kind, e.g. H2, N2.
· A chemical compound is made of different type atoms, e.g. H2O, Ca (OH)2.
The formation and function of molecules depend on chemical bonding between atoms.
Covalent bonds
Covalent bonds involve the sharing of electrons between two atoms in such a way that each valence shell is filled.
· Covalent compounds.
· Single, double, triple and more covalent bonds.
The atoms of a compound are held together by electrical forces of attraction between the nuclei and electrons of the atoms involved.
Each bond contains certain amount of energy called bond energy.
Valence electrons determine how many bonds the atom can form.
Molecules can be represented
1. by their molecular formula (O2),
2. by the Lewis dot structure of O.
3. by their structural formula: O=O
Check: http://www.dataworks-ed.com/sites/default/files/temp/1/CHEM_SCI_2e_LEWIS_DOT_STRUCTURE_DW.pdf
Each molecule has its own characteristic shape, which determines its function.
· Structural formulas show the atoms and their bonds: H-O-H
· Molecular formulas show only the numbers and kinds of atoms in one molecule: H2O.
The attraction of an atom for the electrons of a covalent bond is called its electronegativity.
Covalent bonds can be polar or nonpolar depending on the affinity for electrons of the atoms forming the bond.
· Nonpolar bonds: electrons are shared equally between two atoms.
· Polar bonds: one atom is more electronegative and pulls the electrons closer to its nucleus.
When the electronegativity is very uneven, the more electronegative atom strips one or several electrons away from its partner and the electron is transferred to the more electronegative atom.
When this happens, both atoms end with a complete outer orbit, either by the loss or gain of electrons.
The acceptor of electrons becomes negative because it has increased its number electrons. This number is now greater than the number of protons in the nucleus. It becomes a negative ion.
The loser of electrons becomes positive because the number of electrons is now less than its number of protons. It becomes a positive ion.
Ions are electrically charged atoms or groups of atoms.
· Cations are positive ions.
· Anions are negative ions.
Ionic bonds result from the attraction between anions and cations.
Compounds formed by ionic bonds are called ionic compounds or salts, e.g. NaCl.
The solid structure or crystal is not made of molecules but of ions in a certain ratio.
Ionic compounds dissolve readily in water and form free ions.
The formula for an ionic compound indicates only the ration of atoms in the crystal structure.
Weak bonds
Molecules interact with each other and attract each other temporarily. This temporary and weak intermolecular attractions are very important biologically.
- Hydrogen bonds are relatively weak bonds formed when a hydrogen atom with a partial positive charge is attracted to an atom, usually a nitrogen or oxygen, with a partial negative charge already bonded to another molecule or another part of the same molecule.
- Van der Waals interactions are weak temporary attractions between atoms and molecules.
Electrons are not evenly distributed around the nucleus and are not static.
Van der Waals interactions occur when transient positive and negative regions of molecules attract each other.
Summary:
· Polar covalent bonds
· Non-polar covalent bonds
· Ionic bonds
· Hydrogen bonds
· Van der Waals forces
Check: http://bio.phys.unm.edu/500-09/Panel_02-3.pdf
Note: there is another kind of bond called “metallic bond” that is found in solid pure metals. Metallic bond is out of the scope of our biology course but be informed that it exists.
Molecular shape and biological function of molecules
The biological function of a molecule depends on its shape.
The shape of a molecule depends on the position of its valence orbitals.
When covalent bonds form, the s and p orbitals of reacting atoms will form a hybrid orbital that will change the outer arrangement of the electron orbitals and therefore the outline of the atom.
The molecular shape is important because it determines how biological molecules recognize and respond to one another with specificity.
Only molecules with complementary shapes are able to bind to each other by weak bonds.
Check:
http://intro.chem.okstate.edu/1314F00/Lecture/Chapter10/VSEPR.html
http://www.sparknotes.com/testprep/books/sat2/chemistry/chapter4section8.rhtml
http://www.chemmybear.com/shapes.html
http://library.thinkquest.org/C006669/data/Chem/bonding/shapes.html
CHEMICAL REACTIONS
Chemical reactions make and break chemical bonds.
Chemical reactions change reactants into products while conserving matter.
Reactions cannot create or destroy matter, only rearrange it.
Some chemical reactions go to completion: all reactants are converted to products.
Many chemical reactions are reversible. The products of the forward reaction become the reactants of the reverse reaction: 3H2 + N2 ↔ 2NH3
Chemical equilibrium occurs when the forward and reverse reaction rates are equal.
Summary of key concepts – MUST KNOW:
- Elements and compounds – what is the difference?
- Essential elements of life – why are they called essential? Which elements make 96% of living matter?
- Subatomic particles – name them; characteristics; location in the atom.
- Atomic number and atomic mass – definition
- Isotopes; radioactive isotopes – what are they?
- Energy levels – arrangement in the atom.
- Electron configuration; valence electrons; chemical properties.
- Electron orbitals – shape; hybridization
- Molecules
- Ions
- Bonds
- Covalent bonds
- Ionic bonds
- Weak bonds
· Hydrogen bonds
· Van der Waals forces
- Molecular shape and function – specificity
- Chemical reactions – reactants, products; reversibility; chemical equilibrium
ANOTHER SUMMARY OF CHAPTER 2:
http://www.mansfield.ohio-state.edu/~sabedon/campbl02.htm
July 13, 2017