Introduction & Review: Orbitals, Bonds, Structures & Acidity Highland Hall Biochem Block: Handout #1 Introduction & Review: Orbitals, Bonds, Structures & Acidity Adapted from Organic Chemistry, L. G. Wade, Jr.

Particle Location Mass (amu) Atomic Structure 3 components Particle Location Mass (amu) proton nucleus ~ 1 neutron nucleus ~ 1 electron surrounds ~ 1/1800 the nucleus Electrons determine an atom’s reactivity Recall that opposite charges attract each other, same charges repel…

Atomic Structure atomic number = number of protons atomic weight ≈ # of protons + # of neutrons typically #(protons) = #(neutrons) Isotope: atoms with different number of neutrons than protons

Electronic Structure Atomic orbital : space occupied by an electron associated with a nucleus Electron density : statistical probability of finding the electron at a given location (density maps) Orbital shells described by principal quantum number (n) As n ↑, energy ↑, and (# of e-) ↑ Atomic orbitals are defined by three quantum numbers: n, ℓ, mℓ

Quantum Numbers n : orbital shell, determines size + energy of an orbital aka : principle quantum number ℓ : orbital subshell, determines shape of an orbital ranges from 0 to n-1, for a given n. mℓ : magnetic number, determines spatial orientation of an orbital, an integer ranging from - ℓ, … , 0 , …, +ℓ ms : spin of an electron (-½, +½) ℓ 1 2 3 4 orbital type

1s Atomic Orbital e- density highest at nucleus Spherically symmetrical (Same in all three dimensions)

In-class exercise Which is a better model for a 1s orbital – a basketball or an avocado? Can you think of another object/food that might be an even better model than the two above?

2s Atomic Orbital e- density highest at nucleus spherically symmetrical Has a node: region of zero e- density Chapter 1

2p Atomic Orbitals node (plane) at nucleus directional degenerate: 2px, 2pY, 2pz degenerate: Chapter 1

The Aufbau “building up” principle Distribute the e- so that the energy is as low as possible # of e- in an atom = its atomic number Each added electron will enter orbitals in order of increasing energy Any given orbital cannot take more than 2 e- (__________ _________ Principle)

Electronic Configurations  =>    Hund’s rule:

Partial Periodic Table Valence electrons:

Table: electronic configurations =>

Octet Rule Filled shells of electrons are particularly stable Example : noble gases definition 1 : atoms transfer or share e- in such a way as to attain a filled shell definition 2 : atoms undergo bonding by gaining, losing, or sharing e- so as to acquire the electron configuration arrangement of a noble gas. For atoms in 2nd row of the periodic table, a filled shell = 8 valence e- “octet rule”

Ionic Bond Formation Ionic bonding: electrons are transferred from one atom to another to attain noble gas configuration. The ionic bond is the result of attraction between oppositely charged ions.

Covalent Bond Formation Covalent bonding: electron pair is shared between 2 nuclei. e- not transferred Polar covalent bond: electron pair is unequally shared

Lewis Structures A way to symbolize covalent bonds in molecules Bonding electrons : symbolized by a line Nonbonding electrons/ lone pairs: symbolized by dots

Valence :

Multiple Bonding Single bond : sharing of one pair of e- between two atoms Double bond : sharing of two pairs of e- between two atoms Triple bond : sharing of three pairs of e- between two atoms Carbon has a valence of 4 =>

Multiple Bonding Double bond example : ethylene and formaldehyde Triple bond example : acetylene => Chapter 1

In class exercise: Drawing Lewis Structures Chapter 1

Covalent Bond Types Non-polar : bond between atoms in which electrons are shared equally (C―C, H―H, C―H) Polar : bond between atoms in which electrons are shared unequally (O―H, C―Cl) Ionic bonds are simply the most extreme case of a polar bond. To predict polarity of a bond, one exploits electronegativities. =>

Electronegativity Electronegativity : ability of a bonded atom in a molecule to attract electrons. Atoms with higher electronegativities pull the shared electrons more strongly. Electronegativity increases from left to right and from bottom to top.

Electronegativity and Bond Polarity The larger the difference in electronegativities between two atoms the more polar the covalent bond. If DE.N. > 2, bond is ionic To quantify how polar a bond is, calculate the dipole moment. Electronegativity values:

Dipole Moment (μ) Dipole moment : amount of charge separation. Charge separation shown by electrostatic potential map Red = δ+ region Blue = δ- region Chapter 1

Structural Formulas CH3CHOHCH2CHFCH3 Full structural formula (no lone pairs shown) Line-angle formula stick figure ends of lines = carbon assume enough H’s to give each carbon 4 bonds show all heteroatoms Condensed structural formula not shown in actual bonding order CH3CHOHCH2CHFCH3

Chapter 1

Drawing molecules in 3-D using the line-wedge formula

In class exercise: Make these molecules What is the relationship between molecules A and B? Between B and C?

BrØnsted-Lowry Acids and Bases Acids can donate a proton (H+) in water Bases can accept a proton (H+) in water These reactions form conjugate acid-base pairs H2O acid base H2O acid base

Conjugate Acid-base Pairs Conjugate acid : acid formed upon protonation of a base Conjugate base : base formed upon the deprotonation of an acid H2O acid base conjugate base conjugate acid H2O acid base conjugate base conjugate acid

Acid and Base Strength Acid dissociation constant, Ka Ka is the approximate strength of the acid. Base dissociation constant, Kb For conjugate pairs, (Ka)(Kb) = Kw Spontaneous acid-base reactions proceed from stronger to weaker. pKa 4.74 pKb 3.36 pKb 9.26 pKa 10.64

Acid Strength Example : Water

Typical pKa values pKa α 1/ acid strength

In class exercise: predicting acid base reactions which reactions are favorable?