chapter 1 structure and bonding
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Chapter 1 Structure and Bonding. What is Organic Chemistry?. Living things are made of organic chemicals (carbon-based compounds) Proteins that make up hair DNA, controls genetic make-up Foods, medicines. Origins of Organic Chemistry. Foundations of organic chemistry from mid-1700’s. - PowerPoint PPT PresentationTRANSCRIPT
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