© Copyright 1999 Lawrence Chasin and Deborah Mowshowitz Department of Biological Sciences Columbia University New York, NY

Last updated Tuesday, July 18, 2000 12:47 AM

Bio C2005/F1402x     Lec.2         L. Chasin     September 9, 1999

1 handout of 14 pages. Please bring this handout to class for the next two weeks.

The problem book (Mowshowitz, 14th edition) is available at the Village Copier (115th&Bway starting Thursday afternoon.  The 14th edition is required this year.

Recitations start next week. Look for lists, times on the Web site by Monday at noon..  No quiz the first week.

Outline:

The 5 categories of E. coli's molecules:

Polysaccharides, Lipids, Nucleic Acids, Proteins, and Small molecules

Water

Chemical bond types:

strong (covalent);

weak: hydrogen, ionic, hydrophobic, van der Waals

Macromolecules vs. small molecules

Polysaccharides

Lipids

General links of interest for this lecture:

See E. coli grow

A quick basic review of molecular structure and chemical bonding

A simple view of water

More about water

A simple view of chemical bonds

A more in-depth discussion of functional groups in organic chemistry

We  are starting on the problem of how the bacterium E. coli reproduces, how it grows; how we get two E. coli cells from one. But first we need to know what are the chemicals that need to be made if we are to create one net E. coli cell.

Let's start with the most abundant and most important molecule in the cell, not an organic molecule, but water, H2O.  We will use our discussion of the water molecule as a springboard for introducing different types of chemical bonds that are important in biology.

A more accurate representation than H20 would be: HOH, showing that each hydrogen atom is bonded to the oxygen atom.

Or H-O-H, where the dashes indicate covalent bonds, the sharing of electrons.

Or, showing the actual angle formed by the 2 bonds to H:

,

or, showing the sharing of electrons more explicitly:

 .

Covalent bonds are very strong bonds. We can compare the strength of various chemical bonds by noting the amount of energy it would take to break the bond. For most covalent bonds this is about 100 kilocalories per mole (kcal/mole), where calories are a unit of energy (1 cal = energy needed to raise 1 gram water 1 degree C; Calorie in dieting (with an uppercase C) = kilocalorie (kcal)). 

A yet more accurate portrayal would be:

This configuration [Purves 2.18] has important consequences, because although the electrons are indeed shared between the H's and the O, they are not shared equally. The oxygen nucleus is more electronegative than the hydrogen nucleus, that is, it attracts the shared electrons more strongly than the hydrogen nuclei. As a result, the O is slightly negatively charged and the H's are slightly positively charged.

 

The 's indicate a partial charge, as opposed to a full charge You'd get a full charge if the electron were to be completely captured by one of the partners, resulting in the formation of charged ions, as chlorine atom does in table salt (NaCl --> Na+ and Cl-).

So water is a polar molecule (one with a charge separation), and this property has profound consequences for biological molecules. [Purves 2.8]

As a result of this polarity, each water molecule can be attracted to another water molecule, depending on the orientation. This attraction is very sensitive to orientation, being sharply maximal when the O - H - O atoms are lined up:

waterhb.gif (5131 bytes)

These connections between two molecules are called hydrogen bonds. Their strength is about 3 kcal/mole, thus they are weak bonds compared to the strong covalent bonds of ~100 kcal/mole.  To chemically break covalent bonds by the thermal motion induced by heat, you would typically need hundreds of degrees (e.g., breaking oxygen-oxygen bonds when burning coal). In contrast, hydrogen bonds are readily disrupted at temperatures between freezing and boiling (0o - 100oC). In fact, freezing and boiling of water is a reflection of the hydrogen bonding: Gas = so much thermal motion that no hydrogen bonding is possible; liquid = H-bonds are constantly forming, breaking, and re-forming; solid = hydrogen bonds are locked in a stable crystal structure, which is ice [Purves 2.19]  (see another ice) .

Are there H-bonds in compounds other than water? Sure. Consider ethanol (alcohol), which has an hydroxyl group (-OH, see functional groups handout and polarity I handout and optionally other links on functional groups; we will be discussing most of these functional groups at one time or other). Compare CH3-CH2-OH, vs. ethane (CH3-CH3) which does not have this polar hydroxyl group. The hydroxyl group is polar, for the same reason as in water. So it can H-bond to water when it is in an aqueous solution (as most biological molecules are).  It is for this reason that most compounds with polar groups are very soluble in water. That is, they are constantly forming thse weak bonds to the water molecules.

Note that carbon always forms 4 bonds.

And the H-bonds are not limited to oxygen in O-H groups: nitrogen is also more electronegative than hydrogen, as in an amide (-CO-NH2), and oxygen is more electronegative than carbon (as in the carbonyl in the same amide):

("R" is shorthand for any general organic group, one that is not necessarily relevant for the discussion at hand.)

 

How about H-bonds between organic molecules? Sure, if they can find each other: e.g., ethanol-acetamide, and the orientation is important here, as with water. (If the amide in the diagram were acetamide, the R would be CH3.)

In aqueous solutions such interactions will always be competing with water molecules, which are usually more abundant....

Having introduced the subject of weak bonds, I want to now complete the discussion of bonds by introducing all of the bonds that play important roles in the behavior of biological molecules. There are five:

 

1. covalent bonds: electrons shared, strong [bond energy of ~ 100 kcal/mole] = energy needed to pull the 2 bonding atoms apart]

2. hydrogen bonds:

water-water [~2-3 kcal/mole]

organic molecule - water

organic - organic molecule

orientation-dependent

 3. Ionic bonds: Full charge transfer; NaCl = strong in the dry crystal (need a hammer to break it)

But ionic bonds are weak in water. Why? Water can H-bond to the charged ions: Na+ and Cl-.  This process is called solvation [Purves 2.12].

So is this bond between water and the ion an H-bond or ionic bond? Half and half. Maybe 5 kcal/mole.

 --- 3A. Organic molecules can form ions too (acids and bases):

ACIDS: molecules that are able to lose a proton (hydrogen ion) easily, such as a carboxyl group (a carboxylic acid):

R-CO-OH  ---> R-COO- (net charge = -1), + H+.

BASES: molecules that are able to take up a proton easily (protons being always around to some extent in water [ e.g., at 10-7M at pH7]), such as amines:

R-NH2 + H+ --> R-NH3+

Carboxylic acid will be the only organic acid and amines will be the only organic base we will discuss this semester

Acidity and basicity are measured by pH  (= -log[H+]) [Purves 2.21]

Under the right conditions, ionic bonds can form between two organic ions, with a bond strength of about 5 kcal/mole (in water):

Where are we going with all this chemistry, and these weak bonds? We started describing the molecules of E. coli, with the idea that we have to know what we have, in order to know what we have to make, to replicate an E. coli cell. The weak bonds I am cataloging for you now show how these molecules can interact - but as we proceed to consider larger and larger molecules, they will help us to understand the structure of the individual large molecules, such as proteins and DNA. So this is more than just a listing, the weak bonds will be very important, as we will see in the next few lectures.

 4. Van der Waals bonds: Exist between any 2 molecules

Only effective at very close range (e.g., 0.1 nm, or 1A).

Fluctuating induced dipole.

~1 kcal/mole.

These are the weakest of the bonds we'll discuss, about 1 kcal/mole, but they are able to form between any two molecules. Van der Waals interactions form between fluctuating induced dipoles. Take for example two methyl groups, where the C and H have about the same electronegativity, so there is no intrinsic charge separation. A momentary negative charge can develop in the electron distribution around one of these atoms, and this charge will induce the opposite charge in a nearby atom's electron cloud. These bonds are only effective at extremely short range (~ "touching"). Indeed, the size of an atom in space is often estimated by its "van der Waals radius." (Closest approach before repulsion between nuclei sets in). 

5. Hydrophobic interactions ("bonds")

Not really bonds, but often referred to as such.

Caused by the effects of water on the association of other molecules.

Non-polar (apolar) molecules are unable to form H-bonds with water.

E.g., octane, CH3-(CH2)6-CH3.  (~= gasoline)

Water molecules surrounding an apolar molecule take on a relatively ordered structure compared to the constantly switching H-bonding patterns made with other water molecules.

This ordered cage structure is minimized by interfacing the apolar molecules with each other:

Systems will change so as to maximize entropy (disorder).   Even though the octane molecules are more ordered when aggregated, the increase in disorder of the water molecules that become freed from the cage structure is so great that the entropy of the system is greater with the octanes coalesced.  This increase in entropy provides a hydrophobic force equivalent to about 2-3 kcal/mole (mole of octane, in this case).

The actual bonds between the octane molecules in a coalesced glob in water are just the van der Waals bonds.

This state of affairs is not intuitively obvious.

The bottom line is that apolar groups will tend to associate with other apolar groups in aqueous solutions. Click here for an alternate view of hydrophobic interactions. [Purves 2.14]

This finishes our introduction to the chemical bonds that will be important in our consideration of biological molecules.

Let's get back to the chemical make-up of an E. coli cell: ……………. Water was molecule #1.

We have about 5000 more different types to consider. Before proceeding to #2, let's place all molecules into 2 classes:

1) small; and 2) large.

Small ~< 500 daltons, or molecular weight units), corresponding to about 50 atoms; Large ~> 5000 daltons (~500 atoms). 
These size distinctions are not sharp boundaries, just a rough gauge.

The large molecules are usually called macromolecules. The small molecules are just called small molecules.

Small = e.g., a molecule like propylene, a synthetic organic chemical: CH3-CH=CH2. What is a large version of this small molecule? It is not like the picture on the left (where each circle represents propylene):  

 ( Incorrect picture)

 

 

Rather, polypropylene, a familiar plastic, is a linear polymer of propylene: O-O-O-O-O-O-O-O-O-O-O-, where -O- = the bracketed repeating group in the picture below:

 

 

Virtually all of the biological macromolecules are built this same way: they are linear polymers of small molecules. This simplifies matters greatly.

Monomers --> di-mers (two small molecules linked together covalently) --> tri-mers  --> tetramers, etc. ….--> oligo-mers (moderate numbers of repeating units),   --> poly-mers (lots of repeating units). 

The monomers could be all the same, as in certain polysaccharides like cellulose (glucose (n)) or they could be different, the extreme example being proteins, where there is a mixture of 20 different monomers present.

While this greatly simplifies our consideration of these large molecules, there is plenty of complexity left.

Many of the important small molecules in the cell are these monomers, the basic building blocks of the biological polymers. The polymers, or macromolecules, comprise 4 classes:

polysaccharides,

lipids,

nucleic acids, and

proteins.

 The total number of such monomers is about 40..... Pretty simple...

Now there are about a dozen or so other small molecules that serve other functions (they are not monomers that will end up in polymers). These are co-factors that are important in the catalysis of chemical reactions in the cells. So this brings us to ~50 different small molecules so far.

Then, necessary but less generally important, are the "intermediates".  All the carbon in E. coli can flow from glucose via biochemical pathways (see flow handout (handout p.13) for overview):

glucose --> A --> B --> C --> D --> E --> monomer -->  polymer

[A, B, C, D, and E here are the intermediates.]

These pathways are of various lengths. If we take 10 as a generous estimate of the intermediate steps in an average pathway, then we get another 450 (i.e., 9 x 50) different small molecules to add to our total in the E. coli cells. So our final number of small molecules is about 500. Not too great a number to master. Most are known. We will get to know the majority of the end-products, the monomers, as well as a few of the intermediates.

We will continue our discussion of the molecules of E.coli by focusing on the polymers, the monomers being considered in the context of the macromolecules of which they are a part.

A simple overview of the kinds of molecules in the cell, then, is:

 

 

 

 

 

 

 

 

So, first, POLYSACCHARIDES:

The MONOMERS here are sugars, which are poly-hydroxy-aldehydes or ketones, as we see below.  (They are also carbohydrates [compounds with the general formula (CH2O)n].This terminology is a bit confusing and in fact not too important.) Most common is glucose (also called dextrose). Let's look at the structure of this hexose (sugar with 6 carbon atoms).  We will use this opportunity to describe the conventions of organic chemistry in depicting molecules on paper

 

 

 

 

 

 

Note the functional groups here: a carbonyl (aldehyde or ketone) and a bunch of hydroxyls.

Got this far in live lecture. Lecture 3 starts from this point. ***********************************

A ring tends to form from this straight chain in water; this ring closing reaction occurs by itself, and produces eitherof two distinct kinds of rings, rings with different conformations in space:

Alpha-glucose: (C1 and C2 OH's on the same side) or Beta-glucose: (C1's OH on the opposite side from C2's): see glucose handout, handout p.11, right side. 

Note one shorthand convention of organic chemistry: the intersection of lines is assumed to have a carbon atom, and it is understood that carbon always makes 4 bonds. If a bond line ends with nothing attached, then it is assumed that a H atom resides there. For numbering, the carbon at the end where most bonds to oxygen occur is number one.

When carbon has bonds to four different atoms, the bonds extend to form a tetrahedron; that is, they all extend at equivalent angles away from each other. As a result, the ring cannot be flat, but is is actually puckered, into a reclining chair-like shape, which is hard to draw: (see glucose handout, left side). In this chair-view the hydrogens and the hydroxyls end up either axial (vertical, up OR down) or equatorial (horizontal, sticking out) relative to the rough plane of the reclining chair.  A direct comparison of the "flat" ring depiction and the more realistic puckered chair can be seen on the handout, p.12.

Note that in glucose all the hydroxyls are equatorial except that of the #1 carbon in the alpha conformation, which is axial.

Two glucose monomers can be connected to form a DIMER. This connection, WHICH DOES NOT HAPPEN BY ITSELF (i.e., without some help, to be defined later), involves a dehydration, the removal of one molecule of water, from the 2 monomers:

                         dehydration
2 monomers -------------------> dimer.

Dehydration (sometime called condensation) to produce a connection between 2 monomers is a common theme in biochemical reactions,i.e., reactions that take place in living cells.

Conversely, the breakdown of polymers back to their constituent monomers involves the reversal of this chemistry, the addition of water, or hydrolysis.  

            hydrolysis
dimer ------------------> 2 monomers

Both of these reactions require different catalysts in the cell in order to occur.  Specific catalysts are generally required to effect all the biochemical reactions we will discuss.

A dimer of 2 sugars is called a disaccharide. Note the 1-4 linkages on the disaccharide handout.

There are several different hexoses in most cells. Fructose, galactose, and mannose are some common ones.

And there are several common disaccharides (Becker: p.65):

glucose-glucose via a 1-4 alpha-link is maltose, where alpha refers to the state of the OH in the monomer joined at its C1 carbon [Purves 3.12a]

glucose + galactose via a 1-4 beta-link is lactose (in milk),

glucose + fructose = sucrose (table sugar). [Purves 3.12c]

But these are not yet polymers, or macromolecules.   Next time: POLY-saccharides.