1. Chapter 2
Chemical Foundations

2. Lecture 3, Slides
The Chemicals of Life

3. Lecture 3, Slides
The Chemicals of Life
(b) Macromolecules (23%)

4. Atoms Carbon atom neutron electron proton Hydrogen atom
ECB Fig. 2-2
Carbon atom
atomic number (protons) = 6
atomic mass (protons + neutrons) = 12
Hydrogen atom
atomic number = 1
atomic mass = 1

5. Energy levels, Energy Shells, Orbitals
ECB, Fig. 2-5

6. Covalent bonds
Formed when two different atoms share electrons in the outer atomic orbitals
Each atom can make a characteristic number of bonds (e.g., carbon is able to form 4 covalent bonds)
Covalent bonds in biological systems are typically single (one shared electron pair) or double (two shared electron pairs) bonds

7. Covalent Bonds
ECB, Fig. 2-6

8. The making or breaking of covalent bonds involves large energy changes
Lecture 3, Slides
The making or breaking of covalent bonds involves large energy changes
In comparison, thermal energy at 25ºC is < 1 kcal/mol

9. Covalent bonds have characteristic geometries
Lecture 3, Slides
Covalent bonds have characteristic geometries
Figure 2-2

10. Covalent double bonds cause all atoms to lie in the same plane
Lecture 3, Slides
Covalent double bonds cause all atoms to lie in the same plane

11. Lecture 3, Slides
A water molecule has a net dipole moment caused by unequal sharing of electrons
Figure 2-3

12. Asymmetric carbon atoms are present in most biological molecules
Lecture 3, Slides
Asymmetric carbon atoms are present in most biological molecules
Carbon atoms that are bound to four different atoms or groups are said to be asymmetric
The bonds formed by an asymmetric carbon can be arranged in two different mirror images (stereoisomers) of each other
Stereoisomers are either right-handed or left-handed and typically have completely different biological activities
Asymmetric carbons are key features of amino acids and carbohydrates

13. Stereoisomers of the amino acid alanine
Lecture 3, Slides
Stereoisomers of the amino acid alanine
Figure 2-12

14. Lecture 3, Slides
Different monosaccharides have different arrangements around asymmetric carbons
Figure 2-8

15.  and  glycosidic bonds link monosaccharides
Lecture 3, Slides
 and  glycosidic bonds link monosaccharides
Figure 2-17

16. Noncovalent bonds
Several types: hydrogen bonds, ionic bonds, van der Waals interactions, hydrophobic bonds
Noncovalent bonds require less energy to break than covalent bonds
The energy required to break noncovalent bonds is only slightly greater than the average kinetic energy of molecules at room temperature
Noncovalent bonds are required for maintaining the three-dimensional structure of many macromolecules and for stabilizing specific associations between macromolecules

17. The hydrogen bond underlies water’s chemical and biological properties
Lecture 3, Slides
The hydrogen bond underlies water’s chemical and biological properties
Molecules with polar bonds that form
hydrogen bonds with water can dissolve in water and are termed hydrophilic
Figure 2-6

18. Hydrogen bonds within proteins
Lecture 3, Slides
Hydrogen bonds within proteins

19. Ionic bonds
Ionic bonds result from the attraction of a positively charged ion (cation) for a negatively charged ion (anion)
In ionic bonds, electrons are not shared. The electron is completely transferred from one atom to another atom.
Ions in aqueous solutions are surrounded by water molecules, which interact via the end of the water dipole carrying the opposite charge of the ion

20. Ionic bonds

21. Ions in aqueous solutions are surrounded by water molecules
Lecture 3, Slides
Ions in aqueous solutions are surrounded by water molecules
Figure 2-5

22. van der Waals interactions are caused by transient dipoles
Lecture 3, Slides
van der Waals interactions are caused by transient dipoles
When any two atoms approach each other closely, a weak nonspecific attractive force (the van der Waals force) is created due to momentary random fluctuations that produce a transient electric dipole
Figure 2-8

23. Multiple weak bonds stabilize large molecule interactions
Lecture 3, Slides
Multiple weak bonds stabilize large molecule interactions
Figure 2-10

24. Chemical equilibrium Keq = [X][Y] A + B X + Y [A][B]
The extent to which a reaction can proceed and the rate at which the reaction takes place determines which reactions occur in a cell
Reactions in which the rates of the forward and backward reactions are equal, so that the concentrations of reactants and products stop changing, are said to be in chemical equilibrium
At equilibrium, the ratio of products to reactants is a fixed value termed the equilibrium constant (Keq) and is independent of reaction rate
A + B X + Y
Keq = [X][Y]
[A][B]

25. Equilibrium constants reflect the extent of a chemical reaction
The Keq is always the same for a reaction, whether a catalyst is present or not.
Many reactions involve non-covalent binding of one molecule to another. For these reactions we usually refer to KD, dissociation constant, which is the inverse of the Keq.
For example, KD is the term we use to describe the affinity of a ligand for a receptor.
The lower the KD, the higher the affinity for the receptor.

26. Biological fluids have characteristic pH values
All aqueous solutions, including those in and around cells, contain some concentration of H+ and OH- ions, the dissociation products of water
In pure water, [H+] = [OH-] = 10-7 M
The concentration of H+ in a solution is expressed as pH
pH = -log [H+]
So for pure water, pH = 7.0
On the pH scale, 7.0 is neutral, pH < 7.0 is acidic, and pH > 7.0 is basic
The cytosol of most cells has a pH of 7.2

27. Hydrogen ions are released by acids and taken up by bases
When acid is added to a solution, [H+] increases and [OH-] decreases
When base is added to a solution, [H+] decreases and [OH-] increases
The degree to which an acid releases H+ or a base takes up H+ depends on the pH

28. Biochemical energetics
Living systems use a variety of interconvertible energy forms
Energy may be kinetic (the energy of movement) or potential (energy stored in chemical bonds or ion gradients)

29. G = Gproducts - Greactants
The change in free energy determines the direction of a chemical reaction
Living systems are usually held at constant temperature and pressure, so one may predict the direction of a chemical reaction by using a measure of potential energy termed free energy (G)
The free-energy change (G) of a reaction is given by
G = Gproducts - Greactants
If G < 0, the forward reaction will tend to occur spontaneously
If G > 0, the reverse reaction will tend to occur
If G = 0, both reactions will occur at equal rates

30. Many cellular processes involve oxidation-reduction reactions
The loss of electrons from an atom or molecule is termed oxidation and the gain of electrons is termed reduction
If one atom or molecule is oxidized during a chemical reaction then another molecule must be reduced
The readiness with which an atom or molecule gains electrons is its redox potential E. Molecules with -E make good electron donors. Molecules with +E make good electron acceptors.

31. The oxidation of succinate to fumarate
Lecture 3, Slides
The oxidation of succinate to fumarate
Figure 2-25

32. An unfavorable chemical reaction can proceed if it is coupled to an energetically favorable reaction
Many chemical reactions are energetically unfavorable (G > 0) and will not proceed spontaneously
Cells can carry out such a reaction by coupling it to a reaction that has a negative G of larger magnitude
Energetically unfavorable reactions in cells are often coupled to the hydrolysis of adenosine triphosphate (ATP), which has a Gº = -7.3 kcal/mol
The useful free energy in an ATP molecule is contained is phosphoanhydride bonds

33. The phosphoanhydride bonds of ATP
Lecture 3, Slides
The phosphoanhydride bonds of ATP
Figure 2-24

34. ATP is used to fuel many cell processes
The ATP cycle
Figure 1-14

35. Activation energy and reaction rate
Lecture 3, Slides
Activation energy and reaction rate
Many chemical reactions that exhibit a negative G°´ do not proceed unaided at a measurable rate
Chemical reactions proceed through high energy transition states. The free energy of these intermediates is greater than either the reactants or products

36. Lecture 3, Slides
Example changes in the conversion of a reactant to a product in the presence and absence of a catalyst
Enzymes accelerate biochemical reactions by reducing transition-state free energy