Chapter 1 Structure and Bonding

John E. McMurry

www.cengage.com/chemistry/mcmurry

Paul D. Adams • University of Arkansas

Chapter 1Structure and Bonding

Living things are made of organic chemicals (carbon-based compounds)

Proteins that make up hair

DNA, controls genetic make-up

Foods, medicines

What is Organic Chemistry?

Foundations of organic chemistry from mid-1700’s.

Compounds obtained from plants, animals hard to isolate, and purify.

Compounds also decomposed more easily.

Torben Bergman (1770) first to make distinction between organic and inorganic chemistry.

It was thought that organic compounds must contain some “vital force” because they were from living sources.

Origins of Organic Chemistry

Because of “vital force”, it was thought that organic compounds could not be synthesized in laboratory like inorganic compounds.

1816, Chevreul showed that not to be the case, he could prepare soap from animal fat and an alkali and glycerol is a product

1828, Woehler showed that it was possible to convert inorganic salt ammonium cyanate into organic substance “urea”

Origins of Organic Chemistry

Organic chemistry is study of carbon compounds. Why is it so special? 90% of more than 30 million chemical compounds contain carbon. Examination of carbon in periodic chart answers some of these questions. Carbon is group 4A element, it can share 4 valence electrons and form 4

covalent bonds.

Origins of Organic Chemistry

Abundance of Organic Compounds

• Why are there so many more organic compounds than inorganic?

• Carbon has unique bonding characteristics – Strong, covalent bonds with C and H

• Isomerism– Groups of carbon atoms can form more than one unique compound

C C O

H

H

H

H

H

H

C O C

H

H

H

H

H

H

Structural Isomers of C2H6O

Structure of an atom Positively charged nucleus (very dense, protons and neutrons) and small (10-15

m)

Negatively charged electrons are in a cloud (10-10 m) around nucleus

Diameter is about 2 10-10 m (200 picometers (pm)) [the unit ångström (Å) is 10-10 m = 100 pm]

1.1 Atomic Structure

The atomic number (Z) is the number of protons in the atom's nucleus

The mass number (A) is the number of protons plus neutrons

All the atoms of a given element have the same atomic number

Isotopes are atoms of the same element that have different numbers of neutrons and therefore different mass numbers

The atomic mass (atomic weight) of an element is the weighted average mass in atomic mass units (amu) of an element’s naturally occurring isotopes

Atomic Number and Atomic Mass

Shells, Subshells, OrbitalsShell

(Corresponds to period)

# of subshells

subshells # orbitals # electrons

n = 1 1 1s 1 2n = 2 2 2s 1 2

2p 3 6n = 3 3 3s 1 2

3p 3 63d 5 10

n = 4 4 4s 1 24p 3 64d 5 104f 7 14

The number of subshells in a shell = shell number The first subshell s has 1 orbital. Each successive subshell adds 2 more orbitals (1, 3, 5,

7, etc). Each orbital can hold only 2 electrons of opposite spin. An atom with n = 3 also includes all subshells and orbitals for n < 3:

1s, 2s, 2p, 3s, 3p, 3d

Quantum mechanics: describes electron energies and locations by a wave equation Wave function solution of wave equation Each wave function is an orbital, ψ

A plot of ψ describes where electron most likely to be Electron cloud has no specific boundary so we show

most probable area, i.e., this is a probability function.

1.2 Atomic Structure: Orbitals

Four different kinds of orbitals for electrons based on those derived for a hydrogen atom

Denoted s, p, d, and f s and p orbitals most important in organic and biological

chemistry s orbitals: spherical, nucleus at center p orbitals: dumbbell-shaped, nucleus at middle d orbitals: elongated dumbbell-shaped, nucleus at center

Shapes of Atomic Orbitals for Electrons

Orbitals are grouped in shells of increasing size and energy

Different shells contain different numbers and kinds of orbitals

Each orbital can be occupied by two electrons

Orbitals and Shells (Continued)

First shell contains one s orbital, denoted 1s, holds only two electrons

Second shell contains one s orbital (2s) and three p orbitals (2p), eight electrons

Third shell contains an s orbital (3s), three p orbitals (3p), and five d orbitals (3d), 18 electrons

Orbitals and Shells (Continued)

In each shell there are three perpendicular p orbitals, px, py, and pz, of equal energy

Lobes of a p orbital are separated by region of zero electron density, a node

P-Orbitals

Bonding Characteristics of Carbon

2p2s1s

Valence Valence shell shell

electronselectrons

C Q: If 2s electrons are already paired, with only 2 2p electrons unpaired, how does carbon form 4 covalent bonds?

Ground-state electron configuration (lowest energy arrangement) of an atom lists orbitals occupied by its electrons. Rules:

1. Lowest-energy orbitals fill first: 1s 2s 2p 3s 3p 4s 3d (Aufbau (“build-up”) principle)

2. Electrons act as if they were spinning around an axis. Electron spin can have only two orientations, up and down . Only two electrons can occupy an orbital, and they must be of opposite spin (Pauli exclusion principle) to have unique wave equations

3. If two or more empty orbitals of equal energy are available, electrons occupy each with spins parallel until all orbitals have one electron (Hund's rule).

1.3 Atomic Structure: Electron Configurations

How are electrons arranged?Aufbau PrincipleElectrons fill orbitals starting at the lowest available (possible) energy states before filling higher states (e.g. 1s before 2s).

Sometimes a low energy subshell has lower energy than upper subshell of preceding shell (e.g., 4s fills before 3d).

Pauli exclusion principle QM principle: no two identical fermions (particles with half-integer spin) may occupy the same quantum state simultaneously (why paired electrons have different spin).

Hund's ruleEvery orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin.

2p2s1s

Energy

Kekulé and Couper independently observed that carbon always has four bonds

van't Hoff and Le Bel proposed that the four bonds of carbon have specific spatial directions Atoms surround carbon as corners of a tetrahedron

1.4 Development of Chemical Bonding Theory

Atoms form bonds because the compound that results is more stable than the separate atoms

Ionic bonds in salts form as a result of electron transfers

Organic compounds have covalent bonds from sharing electrons (G. N. Lewis, 1916)

Development of Chemical Bonding Theory

Lewis structures (electron dot) show valence electrons of an atom as dots Hydrogen has one

dot, representing its 1s electron

Carbon has four dots (2s2 2p2)

Kekulé structures (line-bond structures) have a line drawn between two atoms indicating a 2 electron covalent bond.

Stable molecule results at completed shell, octet (eight dots) for main-group atoms (two for hydrogen)

Development of Chemical Bonding Theory

Atoms with one, two, or three valence electrons form one, two, or three bonds.

Atoms with four or more valence electrons form as many bonds as they need electrons to fill the s and p levels of their valence shells to reach a stable octet.

Carbon has four valence electrons (2s2 2p2), forming four bonds (CH4).

Development of Chemical Bonding Theory

Nitrogen has five valence electrons (2s2 2p3) but forms only three bonds (NH3).

Oxygen has six valence electrons (2s2 2p4) but forms two bonds (H2O)

Development of Chemical Bonding Theory

Development of Chemical Bonding Theory

Valence electrons not used in bonding are called nonbonding electrons, or lone-pair electrons Nitrogen atom in ammonia (NH3)

Shares six valence electrons in three covalent bonds and remaining two valence electrons are nonbonding lone pair

Non-Bonding Electrons

Covalent bond forms when two atoms approach each other closely so that a singly occupied orbital on one atom overlaps a singly occupied orbital on the other atom

Two models to describe covalent bonding.

Valence bond theory, Molecular orbital theory

Valence Bond Theory: Electrons are paired in the overlapping orbitals and are attracted to

nuclei of both atoms H–H bond results from the overlap of two singly occupied

hydrogen 1s orbitals H-H bond is cylindrically symmetrical, sigma () bond

1.5 Describing Chemical Bonds: Valence Bond Theory

Reaction 2 H· H2 releases 436 kJ/mol Product has 436 kJ/mol less energy than two atoms:

H–H has bond strength of 436 kJ/mol. (1 kJ = 0.2390 kcal; 1 kcal = 4.184 kJ)

Bond Energy

Distance between nuclei that leads to maximum stability

If too close, they repel because both are positively charged

If too far apart, bonding is weak

Bond Energy

Covalent bonds can have ionic character These are polar covalent bonds

Bonding electrons attracted more strongly by one atom than by the other

Electron distribution between atoms is not symmetrical

2.1 Polar Covalent Bonds: Electronegativity

4.0EN 1.24.0 EN 1.2EN

Electronegativity (EN): intrinsic ability of an atom to attract the shared electrons in a covalent bond

Differences in EN produce bond polarity Arbitrary scale. As shown in Figure 2.2,

electronegativities are based on an arbitrary scale F is most electronegative (EN = 4.0), Cs is least (EN =

0.7) Metals on left side of periodic table attract electrons

weakly, lower EN Halogens and other reactive nonmetals on right side of

periodic table attract electrons strongly, higher electronegativities

EN of C = 2.5

Bond Polarity and Electronegativity

The Periodic Table and Electronegativity

Nonpolar Covalent Bonds: atoms with similar EN Polar Covalent Bonds: Difference in EN of atoms < 2 Ionic Bonds: Difference in EN > 2

C–H bonds, relatively nonpolar C-O, C-X bonds (more electronegative elements) are polar

When C bonds with more EN atom C acquires partial positive charge, + Electronegative atom acquires partial negative

charge, - Inductive effect: shifting of electrons in a bond in

response to EN of nearby atoms

Bond Polarity and Inductive Effect

Electrostatic potential maps show calculated charge distributions

Colors indicate electron-rich (red) and electron-poor (blue) regions

Arrows indicate direction of bond polarity

Electrostatic Potential Maps

Molecules as a whole are often polar from vector summation of individual bond polarities and lone-pair contributions

Strongly polar substances are soluble in polar solvents like water; nonpolar substances are insoluble in water.

Dipole moment () - Net molecular polarity, due to difference in summed charges - magnitude of charge Q at end of molecular dipole times distance r

between charges = Q r, in debyes (D), 1 D = 3.336 1030 coulomb meter length of an average covalent bond, the dipole moment would be 1.60

1029 Cm, or 4.80 D.

2.2 Polar Covalent Bonds: Dipole Moments

Large dipole moments EN of O and N > H Both O and N have lone-pair electrons oriented away from all nuclei

Dipole Moments in Water and Ammonia

In symmetrical molecules, the dipole moments of each bond have one in the opposite direction

The effects of the local dipoles cancel each other

Absence of Dipole Moments

Partial Charge vs. Formal Charge

Partial charge is a real valueFormal charge may or may not correspond to a real charge

Atoms with FC usually bear at least partial charge ( positive or negative) FC helps us determine overall charge distribution and is useful for

understanding reaction mechanismsNeutral molecules with both a “+” and a “-” are dipolar

2.3 Formal Charges

How to Determine FC

FC = [# of valence e-] – [non-bonding e-] – [shared e-/2]

How to Determine FC

FC = [# of valence e-] – [non-bonding e-] – [shared e-/2]

How to Determine FC

FC = [# of valence e-] – [non-bonding e-] – [shared e-/2]

Atomic sulfur has 6 valence electrons.

Dimethyl sulfoxide sulfur has only 5.

It has lost an electron and has positive charge.

Oxygen atom in DMSO has gained electron and has negative charge.

Formal Charge for Dimethyl Sulfoxide

Formal Charges (Continued)

The terms “acid” and “base” can have different meanings in different contexts

For that reason, we specify the usage with more complete terminology

The idea that acids are solutions containing a lot of “H+” and bases are solutions containing a lot of “OH-” is not very useful in organic chemistry

Instead, Brønsted–Lowry theory defines acids and bases by their role in reactions that transfer protons (H+) between donors and acceptors

2.7 Acids and Bases: The Brønsted–Lowry Definition

“Brønsted-Lowry” is usually shortened to “Brønsted” A Brønsted acid is a substance that donates a hydrogen cation (H+) A Brønsted base is a substance that accepts the H+

“proton” is a synonym for H+ - loss of an electron from H leaving the bare nucleus—a proton

Brønsted Acids and Bases

Hydronium ion, product when base H2O gains a proton

HCl donates a proton to water molecule, yielding hydronium ion (H3O+) [conjugate acid] and Cl [conjugate base]

The reverse is also a Brønsted acid–base reaction of the conjugate acid and conjugate base

The Reaction of Acid with Base

Acidity/Basicity

0 7 14

pH

basicneutralacidic

pH = -log[H+]

H+ OH-

The pH of solution determines form of carboxylic acid Ex. Carboxylate ion predominates at pH 7.4 (physiological pH)

Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

2.8 Acid and Base Strength Strong acids/bases: Dissociate completely in water Weak acids/bases: Dissociate incompletely in water

Strength of acid can be related to acid dissociation constant (Ka) Stronger acids have larger Ka, lower p Ka values. Ka ranges from 1015 for the strongest acids to very small values (10-60) for the weakest

[HB] undissociated acid

[H3O+], [B-] dissociated acid components

pKa’s of Some Common Acids

pKa = –log Ka

The free energy in an equilibrium is related to –log of Keq (G = –RT log Keq)

A smaller value of pKa indicates a stronger acid and is proportional to the energy difference between products and reactants

The pKa of water is 15.74

H2O H

2O OH H

3O

Keq

H

3O OH H

2O

2 and K

aK

eq H

2O

H3O OH

H2O

pKa – the Acid Strength Scale

pKa values are related as logarithms to equilibrium constants Useful for predicting whether a given acid-base reaction will take

place The difference in two pKa values is the log of the ratio of equilibrium

constants, and can be used to calculate the extent of transfer The stronger base holds the proton more tightly

2.9 Predicting Acid–Base Reactions from pKa Values

Organic chemistry is 3-D space Molecular shape is critical in determining the chemistry

a compound undergoes in the lab, and in living organisms

2.12 Molecular Models

2.12 Molecular Models

Build the following compounds with your molecular modeling kit and look at the geometry:

Hexane 2-methylhexane Benzene ethyne