Chapter 2: Chemical Foundations for Cells

Like it or not, chemistry is extremely useful and AP biology involves a great deal of chemistry. This chapter should be largely review – if it is basic review, it will be easy for you. If you don’t remember much of the stuff, be sure to study the chapter in detail, as it will be used throughout the year.

I. Basic (Very Basic) Organization in Chemistry

In terms of size and basic organization, the chemical world works like this (from small/simple to large/complex):

1. Subatomic particles: These particles are smaller than atoms and lack the properties of a specific atom (e.g. 8 protons together do not act like oxygen until they are linked with neutrons and protons in an atom of oxygen). We will work with the basic three subatomic particles (proton, neutron and electron). Protons are positively charged, reside in the nucleus of an atom, and have an atomic mass of about 1 amu. Neutrons are neutrally charged, reside in the nucleus of an atom, and have an atomic mass of about 1 amu. Electrons are negatively charged, orbit around the nucleus in the electron clouds, and have so little mass that we can neglect the number of electrons when calculating mass.

2. Atoms: The smallest unit of mass that still retains the properties of an element. In nature we find 92 elements, and only about 20% are commonly found in living things (this is not to say other elements don’t affect life – read the case study at the beginning of the chapter to see how the element selenium can be detrimental to life). To remember the common atoms of life, remember SPONCH (sulfur, phosphorus, oxygen, nitrogen, carbon and hyrdrogen).

3. Molecules: Molecules are compounds made of two or more atoms chemically combined. If you dissolve salt in water, the salt (NaCl) dissociates in water but is not chemically combined (you can separate them out by boiling the water) so this is a mixture, not a compound. If the chemical formula is written out with no spaces between the atomic symbols, it is a compound (e.g. H2O, NO2, NH4).

II. Some Key Terms

Of course, science wouldn’t be science without some confusing (yet decipherable) exceptions to the rules:

1. Ions: In a balanced, neutral molecule the number of electrons and protons are equal, giving a net charge of 0. When the number of electrons is greater or less than the number of protons, the atoms or molecule becomes an ion, meaning it has a net charge. Diatomic oxygen is neutral; NO3 (-) is a negatively charged polyatomic (more than one atom) ion.

2. Isotopes: The concept of an isotope is perhaps the single most challenging concept to physical science students. If you haven’t learned it in PS, Bio and Chem, study tonight until you get it – you need to understand isotopes!!! If an atom of hydrogen has one proton and one neutron (ignore electrons for all isotopes as they have negligible atomic mass), it has a mass number of 2 amu. If an atom of hydrogen has one proton and no neutrons, it has a mass number of 1 amu. In dating ancient (and not-so ancient) materials, carbon is a famous isotope. In nature, cosmic rays convert some C-12 (how many P and N?) to C-14 (two extra neutrons). (Note: The number of protons NEVER varies from atom to atom of the same element, only the number of neutrons. So even if there were a carbon-312, there would still be only 6 protons with 306 neutrons. If there were a different number of protons, it wouldn’t be carbon!)

3. Radioactivity: So how can C-14 be used in dating materials? Some isotopes are radioactive, meaning they have an unstable nucleus and will tend to break apart over time, emitting neutrons and energy. C-14 is radioactive and unstable so over time (at a very predictable rate of decay), the amount of C-14 will decline. Radioisotopes are important is many, many aspects of biology. Currently, PET scans use isotopes to judge the rate of cell activity inside a living organism. Other radioisotopes can be used a tracers, meaning the location of the isotope is used to judge where a certain physiological activity is occurring.

III. Electrons

If they have nearly no mass, why worry about electrons? They are vital in determining which chemical bonds will occur between atoms. They are the outer portion of an atom so when two or more atoms collide, the electrons will be the subatomic particles interacting with each other. The text uses the term "shell," much to my chagrin. Don’t use this term – electrons move around the nucleus, not within defined shells but rather within clouds of less-than-definite size and shape. Remember, two electrons can fit within the first cloud, eight in the second and eight in the third (don’t worry about the fourth and beyond or the ‘s’ and ‘p’ orbitals for this chapter).

When electrons interact to form bonds, we observe two types of compounds:

1. Ionic: Opposites attract, likes repel. If two (or more) charged ions interact, they sometimes form ionic bonds (electrons not ‘shared’). Table salt (NaCl) is such an example. The Cl will have a negative charge and the Na will have a positive charge. When in contact, there is an ionic bond between the two. Ionic compounds don’t form molecules as such because all ions will form a crystal matrix (given the appropriate conditions and time) so we express ionic compounds with a ionic formula (NaCl) meaning the two atoms occur in a one-to-one ratio, not meaning that every Na is connected to one-and-only-one Cl. You commonly find millions of ions all ionically bound together. Ionic bonds are usually very strong, display a crystal matrix and have high melting points.

2. Covalent bonds: We often say atoms want a full outer cloud (2, 10, 18). The noble gases have a full outer cloud and are therefore very inactive. When atoms lack a full outer cloud, they often share electrons with other atoms, forming covalent bonds. Think of oxygen (6 outer electrons): if it shared four electrons with another oxygen, they both have a full outer cloud. Likewise, if it shares electrons with two hydrogens, each atom will attain a full outer cloud.

2.a. Special Case of Covalent Bond:

Polar vs. Nonpolar: In the case of water, oxygen covalently binds with two hydrogens. Oxygen is more electronegative, meaning it ‘pulls’ more strongly on the shared electrons because it has eight protons (+) versus one for each of the hydrogens. We find that one end of the molecule has more of the electrons than the other end. In this case, the molecule is polar (have a positive and negative end, like a magnet). In biology, we will look often at a special case, called hydrogen bonding. When hydrogen atoms bond, we often see an uneven sharing of electrons (with hydrogen displaying a slightly positive charge). When a covalently-bonded hydrogen comes near a positively charged particle, it will form a weak bond known as a hydrogen bond. This bonding accounts for the water tension of an overfilled water glass and the bonds holding together two strands of DNA.

IV.Water: The Elixir of Life

Water is a polar molecule and there is a general rule that ‘like-dissolves-like.’ Consequently, atoms with even slight charges will freely dissolve in water. In the salt water tank in my room, there are over 80 minerals dissolved in the water. In fact, water is such an effective solvent that it is often called the ‘universal solvent.’ Of course, not everything is water-soluable. The term used to describe molecules that readily dissolve in water is hydrophilic (water-loving). The term used to describe molecules that do not readily dissolve in water is hydrophobic (water-fearing). A substance like grease from an oily pan is an example of a hydrophobic substance.

Because water forms many hydrogen bonds, it takes a great deal of heat energy to change the temperature of water. Thus water remains liquid even with a large input of heat. This property helps stabilize the temperature of organisms as most living things are made of lots of water. (If living things contained a substance that didn’t have hydrogen bonding, they would experience great temperature swings, making many metabolic processes impossible.) Water also forms a crystal lattice when it gets near freezing, resulting in the unusual situation in which solid water is less dense than liquid water. As a result, ice floats and things living under the ice are insulated from temperature swings.

Finally, water has a strong cohesive property. Thus water will adhere to itself, allowing trees to pull water up trunks.

V. pH and Buffers

The presence of H(+), hydronium ions, makes a solution acidic. If a substance accepts hydronium ions, the solution is said to be basic (in other words, contains -OH, hydroxide, ions). We measure pH on a scale from 0 to 14. Below 7, a solution is acidic, and above 7, a solution is basic. As a basic rule, living organisms prefer pH levels close to 7.

The presence of buffers in cells helps guard against wild swings in pH levels. Buffers are weak acid and their associate base that can accept hydroxide or hydronium ions. Look at the equation on page 31 for an example of the use of carbonic acid as a buffer. Buffers are found in many living organism to prevent deadly changes in pH levels.