SOME FUNCTIONAL GROUPS
WHAT IS ORGANIC CHEMISTRY?:
Organic chemistry is the study of the chemistry of carbon (C) compounds. Exceptions include CO, CO2 , carbonates, bicarbonates, cyanates (CNO-), thiocyanates (CNS-), and elemental C compounds like diamond and graphite. Although these compounds contain carbon, their chemistry is typical of inorganic chemicals and they are generally classed as inorganic.
- Most organic compounds contain H.
- Many also contain N, O, P, S, Cl, Br and I.
- Considering only C, H, O, and N, there are over 18 106 known C-containing compounds and this number increases by about 500,000 each year.
- C is unique among the elements in its ability to bond with itself forming long chains of compounds from simple to immense, i.e., from methane (CH4 ) to DNA containing 10’s of billions of C atoms.
- Synthetic organic compounds include medicines, dyes, paints, polymers, food additives, pesticides, fibers, etc.
- Natural organic compounds include the matter contained in all living and once-living organisms, e.g., hair, skin, muscles, genes, food, etc. Aside from water, living organisms are made up primarily of organic compounds.
HISTORY:
- In 1807, J. J. Berzelius coined the name ‘organic’ chemistry for materials derived from living organisms. (‘Inorganic’ refers to materials derived from minerals).
- Up until ~ 1800, the only source of organic chemicals was from living organisms (hence the name ‘organic’). In 1826, Friedrich Wohler converted an inorganic C-containing chemical, ammonium cyanate into urea, a previously known ‘organic’ substance isolated from urine .
Wohler was intrigued that chemicals of the same elemental composition could be different and invented the term ‘isomerism’ to describe this. - Organic chemistry began with chemists synthesizing natural organic compounds. For example, aspirin, which was originally obtained as an extract from the bark of the willow tree, is now produced synthetically in millions of tons per year.
- The availability of large, inexpensive sources of raw materials, i.e., petroleum, coal, and natural gas, has caused synthetic organic chemistry to flourish.
- Unfortunately, < 10% of the fossil fuels consumed are used to make chemicals; > 90% is burned to supply energy.
Some Basics:
Carbon atoms have 4 bonds or less, never 5 bonds!
The 4 bonds on carbon may exist in any of 4 arrangements.
Note that each line (---) in the structures represents a covalent bond, i.e., a pair of electrons that is shared between two atoms.
When carbon has 4 bonds it is neutral (as above), but when carbon has only three bonds it exists in one of 3 possible forms, a cation (‘carbocation’ or ‘carbonium ion’), an anion (‘carbanion’) or a neutral radical.
Electron Configuration of the Elements:
In order to understand organic chemistry we must learn the electron configuration of the first 20 elements (H to Ca) plus Br and I.
Electrons are continuously buzzing around the nucleus at mind-boggling speeds (ca. 1/10 the speed of light). We don’t know the exact position of electrons from one moment to the next (Heizenburg uncertainty principle) but we do know that their movement is not entirely random.
Electrons fly within well-defined flight paths (orbitals) around the nucleus. Each orbital can hold a maximum of 2 electrons. Think about the heavier elements on the periodic table, with 100+ electrons flying around the nucleus in 50+ different flight paths. Inevitably, some of the orbitals overlap. Just imagine how busy their flight controllers must be while trying to prevent all those flying electrons from colliding.
The orbitals lowest to ground zero (the nucleus) are lowest in energy and are occupied by electrons before the outer, high-energy orbitals. The 50+ orbitals around all atoms are grouped into 7 different ‘energy levels’ (also called layers or ‘shells’) with n = 1 being the shell closest to the nucleus (lowest energy) and n = 7 being the farthest from the nucleus (highest energy). Study the order of orbitals and shells in the planetary model of the atom.
PLANETARY MODEL OF THE ATOM SHOWING ENERGY LEVELS n = 1 to 6
This model is not spatially correct. There is some overlap of orbitals in the 3rd shell and higher.
Within each shell, there exist subshells or types of orbitals. The types of orbitals are named s, p, d, f, g, h, etc.
The 1st shell has only an s orbital, named 1s.
The 2nd shell has both s and p-type orbitals, named 2s and 2p.
The 3rd shell has s, p and d-type orbitals, named 3s, 3p and 3d.
The 4th shell has s, p, d and f-type orbitals, named 4s, 4p, 4d and 4f.
etc.
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ORGANIC CHEMISTRY INTRO
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ORGANIC CHEMISTRY INTRO
The number of each type of orbital, their shape and orientation are listed below.
There is only one s-orbital in each shell and it is spherical.
There are three p-orbitals in the 2nd and all higher shells. The three p-orbitals are propeller shaped and are oriented along an x, y or z axis in space. They are named px, py and pz, respectively.
There are five d-orbitals in the 3rd and all higher shells. Four of the five
d-orbitals look like four-leaf clovers each oriented differently around the nucleus. The fifth d-orbital looks like a propeller inside a donut. They are named dxy, dxz, dyz, dx2-y2 and dz2.
There are seven f-orbitals in the 4th and all higher shells. All but one have six lobes. Each one is oriented differently around the nucleus.
In writing the electron configuration of the elements we fill lowest energy orbitals first (Aufbau principle), with a maximum of 2 electrons per orbital –with opposite spins (Pauli Exclusion principle). Orbitals of the same energy level (‘degenerate orbitals’) are all singly filled (half-filled) before electrons pair up. This occurs, for example in the 2px, 2py, and 2pz orbitals.
The filling order (increasing energy level) of the various orbitals is shown in the following chart.
Reading the table left to right and top to bottom, the orbital filling order is as follows:
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p
For the purpose of learning organic chemistry, we need only study the electron configuration of the first 20 elements (H to Ca) plus Br and I, i.e., in the following orbitals:
1s 2s 2p 3s 3p 4s
Br and I have electrons in the 4p and 5p orbitals respectively, but their electron arrangement is analogous to F and Cl.
The electron configuration of atoms is shown using a notation in which the number of electrons in each orbital is written as a superscript. The orbital is shown as a line, _ or as a circle, O. Each electron in the orbital is written as an arrow, . The direction of the arrow is either up, , (indicating clockwise rotation) or down, , (indicating counterclockwise rotation). Complete the following table.
Full Orbital Notation / Simplified Notation1s / 2s / 2px / 2py / 2pz
1H / __ / 1s1
2He / __ / 1s2
3Li / __ / __ / 1s2 / 2s1
4Be / __ / __ / 1s2 / 2s2
5B / __ / __ / __ / __ / __ / 1s2 / 2s2 / 2p1
1s / 2s / 2px / 2py / 2pz / 3s
6C / __ / __ / __ / __ / __ / 1s2 / 2s2 / 2p2
7N / __ / __ / __ / __ / __ / 1s2 / 2s2 / 2p3
8O / __ / __ / __ / __ / __ / 1s2 / 2s2 / 2p4
9F / __ / __ / __ / __ / __ / 1s2 / 2s2 / 2p5
10Ne / __ / __ / __ / __ / __ / 1s2 / 2s2 / 2p6
11Na / __ / __ / __ / __ / __ / __ / 1s2 / 2s2 / 2p6
Problem: Write out the electron configuration for Mg through Ca in both 'Full Orbital Notation' and 'Simplified Orbital Notation'. Recall that the 4s orbital is filled before the 3d orbital.
Full Orbital Notation / Simplified Orbital Notation3s / 3px / 3py / 3pz / 4s
12Mg / [Ne] / / [Ne] / 3s2
13Al
14Si
15P
16S
17Cl
18Ar
19K
[Ca
The outermost occupied shell is referred to as the ‘valence’ shell. Orbitals of the valence shell are thus ‘valence orbitals’ and electrons in the valence orbitals are ‘valence electrons’.
The outer (valence) electrons are transferred or shared in chemical reactions. Chemistry is understood in terms of valence electron arrangement.
The number of valence electrons determines the ability of an atom to combine with other atoms
- The number of covalent bonds an atom forms to become isoelectronic with its nearest noble gas is called its ‘covalence’. (Isoelectronic means ‘having the same valence electronic configuration’.)
- The number of valence electrons in an atom is shown with a Lewis Symbol. One dot is drawn for each valence electron. The dots are placed into four positions (one for each of the one s plus three p orbitals) around the symbol of the element, i.e., north, south, east or west. Once all four positions are singly filled, electrons (dots) are paired up until a maximum of 8 valence electrons (dots) have been drawn.
Although Lewis symbols do not always show the lowest energy electron arrangement of an unbonded atom, they are a good depiction of the electron arrangement just prior to bonding.
When dots (electrons) are drawn, each orbital is first half-filled before electrons are paired up in orbitals.
The arrangement of elements in the periodic table is based on the number of valence electrons. For example, elements in Group IVA have 4 valence electrons.
For all representative elements (A-group elements), the number of valence electrons equals the group number.
- Complete the following table. Note that He is an exception. Although it has 2 valence electrons like Be and Mg, it is unreactive (like other noble gases) and therefore has a valence of 0.
Group # / 1A / 2A / 3A / 4A / 5A / 6A / 7A / 8A
# valence electrons / 1 / 2 / 3 / 4 / 5 / 6 / 7 / 8
covalence
(# bonds)
Period
1 / H / He
Period
2 / Li / Be / B / C / N / O / F / Ne
Period
3 / Na / Mg / Al / Si / P / S / Cl / Ar
Period
4 / K / Ca / Br / Kr
electron
config. / s1 / s2 / s2 p1 / s2 p2 / s2 p3 / s2 p4 / I
s2 p5 / Xe
s2 p6
show lone pairs of electrons after bonding / H / Be / B / C / N / O / F
- Note that covalence is the same as the group number for groups 1A to 4A, but covalence is equal to [8 - (group number)] for groups 5A to 8A.
- It is very important to appreciate the relationship between the Lewis symbols and the number of covalent bonds formed by an atom (its covalence). Usually, in most stable organic compounds, the atoms form a covalent bond for each unpaired electron in the Lewis symbol of the atom.
Study the bonding arrangements of the neutral atoms shown below. Note that all single (unpaired) electrons in a Lewis structure will bond (as shown in the bonded structure). The non bonded electron pairs (‘lone pairs’) may either remain unbonded or form two additional bonds per electron pair. The two additional bonds may be two single bonds or one double bond
Octet Rule:
Note however that 2nd period elements (B, C, N, O and F) will never have more than 4 bonds (8 electrons) around themselves as they can only use four orbitals - their 2s and 2p orbitals for bonding (‘octet rule’).
Hypervalent Atoms:
3rd period elements and higher (Si, P, S, Cl and Br) can form more than 4 bonds (more than 8 electrons) by using their d-orbitals. Examples include PCl5, SF6, ClF7, and BrF7. Such elements that exceed the octet rule are called ‘hypervalent’.
Group # / 3A / 4A / 5A / 6A / 7ALewis
Symbol / / / / /
Bonded
Structure / / / / /
Lewis
Symbol / / / / /
Bonded
Structure / / / /
Carbon and nitrogen are the only two elements that can form a triple bond.
ELECTRONEGATIVITY:
One way to estimate the degree of ionic or covalent character in a chemical bond is to compare electronegativities of atoms involved. Electronegativity is a measure of the force of an atom’s attraction for electrons that it shares in a chemical bond with other atoms.
- In the 1930’s, Linus Pauling assigned electronegativity values to all elements relative to F (the most electronegative element), which he gave a value of 4.0 .
Linus Pauling's Table of Electronegativities
H2.1
Li
1.0 / Be
1.5 / B
2.0 / C
2.5 / N
3.0 / O
3.5 / F
4.0
Na
1.0 / Mg
1.2 / Al
1.5 / Si
1.8 / P
2.1 / S
2.5 / Cl
3.0
K
0.9 / Ca1.0 / Sc
1.3 / Ti
1.4 / V
1.5 / Cr
1.6 / Mn
1.6 / Fe
1.7 / Co
1.7 / Ni
1.8 / Cu
1.8 / Zn
1.6 / Ga
1.7 / Ge
1.9 / As
2.1 / Se
2.4 / Br
2.8
Rb
0.9 / Sr
1.0 / Y
1.2 / Zr
1.3 / Nb
1.5 / Mo
1.6 / Tc
1.7 / Ru
1.8 / Rh
1.8 / Pd
1.8 / Ag
1.6 / Cd
1.6 / In
1.6 / Sn
1.8 / Sb
1.9 / Te
2.1 / I
2.5
Cs
0.8 / Ba
1.0 / La
1.1 / Hf
1.3 / Ta
1.4 / W
1.5 / Re
1.7 / Os
1.9 / Ir
1.9 / Pt
1.8 / Au
1.9 / Hg
1.7 / Tl
1.6 / Pb
1.7 / Bi
1.8 / Po
1.9 / At
2.1
Fr
0.8 / Ra
1.0 / Ac
1.1
Bear in mind that these values can vary slightly depending upon the chemical environment and so the values are average values.
Note that EN increases across any period and decreases down any group (in most cases).
Pure covalent bonds involve equal sharing of the bonding electron pairs. Pure covalent bonds occur when both atoms involved have equal EN (i.e., EN = 0) ...
For example, H2, N2, O2, F2, Cl2, Br2, I2, Cx, S8, PH3, and CS2 are all pure covalent.
Nonpolar covalent bonds are those in which EN 0.4. Examples include all the pure covalent compounds listed above as well as compounds, e.g., CH4 (EN = 0.4) and BH3 (calculate EN)
Polar covalent bonds are those in which the bonding pair of electrons is unequally shared (0.5 EN 1.7). Examples include HBr, HCl, etc. In polar covalent compounds, the more electronegative atom has a partial negative charge (-) and the less electronegative atom has a partial positive charge (+).
Ionic bonds are those in which EN 1.8 and are generally considered to have complete charge separation, i.e., considered to be made up of cations and anions. Examples include NaCl, Li3N, and CaO. Exceptions are HF and alkali metal iodides. EN in HF is 1.9 but this compound behaves as a polar covalent compound. EN in LiI is only 1.5 but LiI behaves ionically. BF3 is also anomalous. It is covalent although EN = 2.0.
- A scale of bond type versus EN follows...
In polar covalent bonds, the electron distribution is said to be polarized, i.e., not equally distributed but closer to the more electronegative atom. For example, in HCl....
EN = (3.0 – 2.1) = 0.9 i.e., polar covalent
A separation of '+' and '-' charge is called a dipole. Dipoles are sometimes illustrated with an arrow pointing toward the more electronegative atom.
+ -The tail of the arrow is crossed to look like a + sign.
The head of the arrow points in the direction of electron shift.
H Cl
The shifting of electron density through sigma bonds due to EN differences between atoms is called an 'inductive effect'. Electropositive elements (metals) such as Zn and Hg, inductively donate electrons through sigma (single) bonds with carbon. Electronegative elements (nonmetals) such as oxygen and chlorine inductively withdraw electrons from carbon through their sigma bonds with carbon.
Problem: Calculate EN values and show bond dipoles and dipole arrows.
BONDING (INTRAMOLECULAR FORCES):
Atoms bond because the resulting compound is more stable (lower energy). For the representative elements, filled valence orbitals (isoelectronic with the nearest noble gas) is a stable arrangement.
Note the # of electrons in each PEL for the noble gases and the # of valence electrons..
period / noble gas / PEL = 1 / PEL = 2 / PEL = 3 / PEL = 4 / PEL = 51 / He / 2
2 / Ne / 2 / 8
3 / Ar / 2 / 8 / 8
4 / Kr / 2 / 8 / 18 / 8
5 / Xe / 2 / 8 / 18 / 18 / 8
- For main group elements, bonds form in which the combining atoms obtain a noble gas electron configuration by either transferring electrons (ionic) or sharing electrons (covalent). A-Group element bond to become isoelectronic with their nearest noble gas.
- Ionization energy (Ei) is the amount of energy added to remove an electron form an isolated atom (endothermic process). Metals have low Ei whereas metalloids and non metals have high Ei. Ei decreases down all groups (due to increased shielding and distance from the nucleus) and from right to left across all periods (due to increased distance from the nucleus).
- Write the full electron configuration of the reactants and products in the reaction ...
Na + Cl [Na+ + Cl-]
Na (1s2 2s2 2p6 3s1) + Cl (1s2 2s2 2p6 3s2 3p5) Na+ (1s2 2s2 2p6) + Cl- (1s2 2s2 2p6 3s2 3p6)
isoelectronic with Neon isoelectronic with Argon
- Ionic bonds result when metals lose electrons to nonmetals forming cations (+) and anions (-). Electrostatic attraction holds the solid together, e.g. ...
Na (g) Na+ (g) + 1e-Ei = + 119 kcal/mol
Cl (g) + 1e- Cl- (g)Eea = - 83 kcal/mol
Na+ (g) + Cl- (g) NaCl (g)Elattice -121 kcal/mol
The lattice energy (Elattice) is the energy released due to coulombic (electrostatic) attraction as cations and anions combine in a crystallatice. - Ionic bonds are common in inorganic compounds, but are less common in organic compounds. The Ei of C is too high (+261 kcal/mol). C shows little tendency to act as a source of a cation or anion in an ionic bond so most of its compounds are characterized by covalent bonds.
Covalent bonding is readily shown using Lewis diagrams or Lewis structures.
Using Lewis structures show the bonding in CH4, CH3OH, and H3O+ (hydronium ion)
Lewis structures show a bonding electron pair as a line ( ) and a nonbonded pair as two dots (.. or :).
Draw Lewis structures for H2, and N2 and include lone pair electrons.
Most of the 2nd and 3rd period representative elements we encounter in organic chemistry (C, N, O, F, Na, Mg, Al, Si, P, S, Cl, etc.) obey the ‘octet rule’, i.e., in forming compounds, atoms will gain, lose or share electrons to obtain 8 valence electrons. This produces a stable valence electron configuration (stable octet) like that of Ne and Ar.
H and Li also react to obtain a noble gas electron configuration, i.e., that of Helium with 2 valence electrons.
Reaction Mechanisms: We show the transfer of electrons with curved arrows. A one-electron transfer is shown with a half arrow head, while a two-electron transfer is shown with a full (double) arrowhead. Arrows point away from the donor atom and toward the acceptor atom. Be careful to draw covalent bonds as a line (a shared pair) but recall that an ionic bond is shown by a non bonded pair and electric charges on the anion and cation. .
EXCEPTIONS TO THE OCTET RULE:
The octet rule indicates that atoms bond to share enough electrons so that each atom has 8 valence electrons after bonding. Most main group elements (Group 1A to 7A) obey this but there are a few exceptions, divided into 2 kinds, i.e., molecules in which atoms contain fewer than 8 valence electrons and molecules in which atoms contain more than 8 valence electrons.
- Molecules in which atoms contain less than 8 valence electrons include BeCl2, AlCl3 and BF3. Be has only 4 valence electrons around it while both Al and B have only 6 valence electrons around them. These compounds thus behave as Lewis acids or electrophiles (electron acceptors). Learn the Lewis structures for these compounds.
- Molecules in which atoms contain more than 8 valence electrons.....
Atoms of the 2nd period use 2s and 2p orbitals for bonding and these orbitals can contain only 8 valence electrons, hence, the octet rule.
Atoms of the 3rd, 4th, and 5th period have ns, np, nd, etc. orbitals and can accommodate more than 8 electrons in their valence shells.
Phosphorus and sulfur have 3s, 3p, and 3d orbitals and in some compounds have > 8 valence electrons. Study the Lewis structures for the following .....
PI3
PI5
SF6
DRAWING LEWIS STRUCTURES OF COMPOUNDS:
- The least electronegative atom is central, e.g., in CO2 O=C=Onot O=O=C
- H is never central (max 1 bond), e.g., in H2OH-O-Hnot H-H-O
- In oxyacids (‘ternary’ acids), e.g., HNO3, H2SO4, HClO3, etc., acidic H’s are bonded to O and O’s are bonded to the central atom,
e.g., in HClO (hypochlorous acid)H-O-Cl notH-Cl=O - Join all atoms with single bonds. Add additional bonds and non-bonded electron pairs consistent with the Lewis symbols and normal covalence of atoms. Non-bonded e pairs may remain unbonded but single electrons always form bonds. Make each atom isoelectronic with its nearest noble gas but ‘3rd period plus’ atoms may be hypervalent. Check for normal bonding as shown on page 10.
- Count (check) the number of valence electrons in the structure. Add 1 e- for each negative charge on an ion and subtract 1 e- for each positive charge on an ion.
Note: Oxygen does not normally bond to itself, except in peroxides,e.g., hydrogen peroxide,
H-O-O-H