Chapter 2 Chemical Foundations

Lecture 3, Slides The Chemicals of Life

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

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

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

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

Covalent Bonds ECB, Fig. 2-6

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

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

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

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

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

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

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

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

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

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

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

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

Ionic bonds

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

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

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

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]

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.

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

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

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)

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

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.

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

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

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

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

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

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