Chemical Bonding and Non-Covalent Interactions for Biochemistry Student Edition 8/27/13 Pharm. 304 Biochemistry Fall 2014 Dr. Brad Chazotte 213 Maddox Hall Web Site: Original material only © B. Chazotte

GOALS Review and be able to identify different chemical functional groups. Understand the significance of and remember the electronegativities of atoms that are biologically important. Review the nature of covalent bonds and the importance of bond, strengths, angles and lengths. To review the nature and importance of dipole moments. Review the types of noncovalent interactions and there biological importance. Review the classes of chemical reactions in cells. Develop an understanding of the chemical logic of biochemistry from the molecular forces and chemical principles that determine chemical properties.

Lehninger 2000 Fig 31 Elements Essential for Animal Life and Health

Lehninger 2000 Fig 3.2 Covalent Bonding and Carbon Atoms H, O, N, & C are the most abundant biological elements. Above are the lightest elements capable of forming 1, 2, 3, & 4 covalent bonds, respectively. Generally true that lightest elements form the strongest bonds.

An Element’s Chemical Reactivity and It’s Outermost Electron Shell Alberts et al, 2004 Figure 2.5

Bonding & Stability Alberts et al, 2004 Figure 2.6 Valance: the number of electrons an atom must gain or lose to have a filled outer shell.

Biomolecules are Carbon Compounds Chemistry of living organisms is organized around the element carbon (“carbon-based life forms”) Molecules with covalently bonded carbon backbones are termed organic compounds. Carbon accounts for 50% of the dry weight of cells Carbon can form: single bonds with hydrogen double bonds with oxygen bonds with other carbon atoms (important!)

Carbon Versatility in Bonding Lehninger 2000 Fig 3.3 Red dot=unpaired electron Carbon can form single, double or triple bonds. Covalently linked carbon atoms in biomolecules can form linear chains, branched chains, or cyclic structures. These chains or skeletons can have other groups of atoms, called functional groups, added to them which give rise specific chemical properties.

Carbon-Bonding Geometry Lehninger 2000 Fig 3.4 The geometry of the bonds and the extent of free or hindered rotation affects the properties of the molecules. tetrahedral arrangement freedom of rotation about the single bond Double bonds are shorter and DO NOT allow for free rotation

Van der Waals Radii and Bond length Lehninger 2000 Table 3.1Matthews et al, 1999 Table 2.2 Van der Waals Radii Note: OTHER, ADJACENT ATOMS EFFECT BOND LENGTHS.

Van der Waals Radii, Bond Lengths & Angles, and the Water Molecule: Diagram Voet. Voet & Pratt, 2013 Fig 2.1 1Å = 1 x cm = 0.1 nm

Biomolecule Chemical Functional Groups Most biomolecules can be considered derivatives of hydrocarbons. Compounds with covalently linked carbon backbone to which only hydrogens are bonded. Backbones of hydrocarbons are very stable. Substitution of the hydrogens by various chemical functional groups a)determines molecule’s chemical properties b)yield different families of organic compounds

Hydroxyl, Carboxyl, & Carbonyl Functional Groups Lehninger 2000 Fig 3.5 R R is used to represent a substituent group (R can be a simple hydrogen atom to complex carbon-containing moieties.)

Methyl, Ethyl & Phenyl Functional Groups Lehninger 2000 Fig 3.5a

Amino, Amido, Imidazole, and Guanidino Functional Groups Lehninger 2000 Fig 3.5b (In Histidine)

Sulfhydral & Disulfide Functional Groups Lehninger 2000 Fig 3.5c

Phosphoryl Functional Group Lehninger 2000 Fig 3.5d

Ester & Ether Functional Groups Lehninger 2000 Fig 3.5e

Anhydride Functional Groups Lehninger 2000 Fig 3.5f

Lehninger 2000 Fig 3.6 Common Functional Groups in a Biomolecule

Chemical Reactivity Matthews et al, 1999 Figure 2.X Biochemical Reactions may be understood and predicted from the nature of the reactant’s functional groups similar to regular chemical reactions. Key concept: Functional groups alter the electron distribution and the geometry of neighboring atoms, thus affecting the chemical reactivity of the entire molecule

Covalent Bond Strength Matthews et al, 1999 Figure 2.X Strength of a chemical bond depends on: Relative electronegativities of the bonding atoms The distance of the bonding electrons from each nucleus Nuclear charge on each atom The number of electrons shared triple > double > single (Bond strength is expressed in terms of bond energy in joules)

Lehninger 2000 Table 3.2 Selected Elements’ Electronegativities The higher the number the higher the affinity for electron and the more strongly electrons will be pulled toward the nucleus. Increasing electronegativity

Lehninger 2000 Table 3.3 Common Biomolecule Bond Strengths Bond energy (one definition): the amount of energy to break a bond Bonds with lower energies can be made to break before other stronger bonds. When bonds are broken and formed in a chemical reaction the energy can be approximated as the enthalpy change,  H = the difference between the energy extracted from the surroundings to break the bonds and the energy released to the surroundings by the formation of a new bond.

Lehninger 2000 Fig 3.11 Conformation and Bond Energy The conformation of a molecule can also affect its bond strength.

Dipole Moments of Some Molecules Matthews et al, 1999 table 2.1 Molecules that have NO net charge may have an asymmetric distribution of the internal charge. that is, a polar molecule, or it is said to have a permanent dipole moment. red arrow =dipole vectorsblue arrow = vector sum dipole moment Note: a dipole moment affects the properties of a molecule.

Noncovalent Interactions of Molecules Matthews et al, 1999 Chapter 2 Much weaker than covalent bond. (10 to 100-fold weaker than, e.g. C-C and C-H bonds) Weakness actually makes them essential - allows them to be continually be broken and reformed –essential to dynamic biological processes. Depends on the rapid interchange of partners. (would be impossible with strong interactions.) Fundamentally electrostatically-based.

Dipole Moments can Interact Voet. Voet & Pratt, 2013 Fig 2.5 Even though very weak the dispersion forces can have a very significant effect on the structure of biological molecules that have many closely packed chemical groups in their interior. Permanent dipole-dipole interaction are weaker than ionic bonds. Dipole-induced dipole interaction are weaker than permanent dipole interactions.

Types of Noncovalent Interactions Matthews et al, 1999 Figure 2.2 a. Charge-Charge b. Charge-Dipole c. Dipole-Dipole d. Charge-induced dipole e. Dipole-induced Dipole f. (London) Dispersion g. van der Waals h. Hydrogen bond

Noncovalent Interaction Energy of Two Approaching Particles Matthews et al, 1999 Figure 2.6

Hydrogen Bonding Matthews et al, 1999 Figure 2.7

H-Bonds in Biologically Important Molecules Matthews et al, 1999 Table 2.3

The Classes of Reactions in Cells Oxidation-Reduction (all involve electron transfer) Cleavage & Formation of Carbon-Carbon Bonds Internal Rearrangements Group Transfers Condensation Reactions (monomeric subunits joined & water eliminated)

Oxidation-Reductions & Oxidation State of Carbon in Biomolecules Lehninger 2000 Figure 3.15 Reduction gain of electrons. Oxidation loss of electrons.

C-C Bond Cleavage Homolytic vs Heterolytic

C-C Bond Heterolytic Cleavage: Nucleophilic Substitution Reactions: S N 1 & S N 2 Lehninger 2000 Figure 3.19a,b Occurs when a second electron- rich group replace the departing anion Nucleophile: functional groups rich in electrons that can donate them.

Internal Rearrangements Lehninger 2000 Figure 3.20

Group Transfer General metabolic theme: attachment of a good leaving group, e.g. phosphoryl group, to a metabolic intermediate to “activate” the intermediate for subsequent reactions

Condensation and Hydrolysis CondensationHydrolysis Important concept: monomeric subunits that make up proteins, nucleic acids, and polysaccharides are joined by nucleophilic displacement reactions that replace a good leaving group.

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