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Summary Essentials Of Organic Chemistry (NWI-MOL101)
Summary of the book Organic Chemistry written by Francis Carey et al. isbn:4248. The chapters 1-6 are summarized with images in this document
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Voorbeeld 4 van de 59 pagina's
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Voorbeeld van de inhoud
Chapter 1 Structure DeterminesProperties
1.1 Atoms, Electrons, and Orbitals
Each element is characterized by a unique atomic number Z which is equal to the number
of protons in its nucleus.
A neutral atom has equal numbers of protons (+ charge) and electrons (- charge).
According to the Heisenberg uncertainty principle, we can’t tell exactly where an electron is,
but we can tell where it is most likely to be. The probability of finding an electron at a
particular spot relative to an atom’s nucleus is given by the square of the wave function Ψ2
at that point.
The probability of finding an electron at a particular point is greatest near the nucleus and
decreases with increasing distance from the nucleus but never becomes zero.
Orbitals (wave functions) are described by specifying their size, shape, and directional
properties.
s orbital
- Letter s is preceded by the principal quantum number, which specifies the shell
and is related to the energy of the orbital.
- An electron in a 1s orbital is likely to be found closer to the nucleus
- Lower in energy than 2s
- More strongly held than 2s
Electron configuration of Hydrogen: 1s1 , Helium: 1s2 , Lithium: 1s22s1
Electrons possess the property of spin. The spin quantum number
of an electron can have a value of either +½ or -1/2. According to the
Pauli exclusion principle, two electrons may occupy the same
orbital only when they have opposite, or “paired”, spins. For this
reason, no orbital can contain more than two electrons.
The period or row of the periodic table in which an element appears
corresponds to the principal quantum number of the highest numbered occupied orbital
(n=1 in the case of hydrogen and helium). Hydrogen and helium are first-row elements;
lithium (n=2) is a second-row element.
Valence electrons are the outermost electrons, the ones most likely to be involved in
chemical bonding and reactions.
For main-group elements, the number of valence electrons is equal to its group number in
the periodic table.
Noble gases are characterized by an extremely stable “closed-shell” electron configuration
and are very unreactive. They have a complete octet of electrons.
1.2 Ionic Bonds
Atoms combine with one another to give compounds having properties different from the
atoms they contain.
The attractive force between atoms in a compound is a chemical bond. One type of
chemical bond, called an ionic bond, is the force of attraction between oppositely charged
species (ions).
Cations: + charged ions
Anions: - charged ions
,In forming ionic compounds, elements at the left of the periodic table typically lose
electrons, giving a cation that has the same
electron configuration as the preceding noble
gas.
Loss of an electron from sodium, yield Na+,
which has the same electron configuration as
neon.
A large amount of energy, called the ionization energy, must be transferred to any atom to
dislodge an electron.
Processes that absorb energy are said to be endothermic. Compared with other elements,
sodium and its relatives in group 1A have relatively low ionization energies. In general,
ionization energy increases across a row in the periodic table.
Elements at the right of the periodic table tend to gain
electrons to reach the electron configuration of the next
higher noble gas.
Adding an electron to chlorine gives the anion Cl-, which
has the same closed-shell electron configuration as the noble gas argon.
Energy is released when a chlorine atom captures an electron. Energy-releasing reactions
are described as exothermic, and the energy change for an exothermic process has a
negative sign. The energy change for addition of an electron to an atom is referred to as its
electron affinity.
Forces between charged particles are called electrostatic, or Coulombic, and constitute
an ionic bond when they are attractive.
Ionic bonds are common in inorganic compounds, but rare in organic ones. Instead of losing
or gaining electrons, carbon shares electrons with other elements to give what are called
covalent bonds.
1.3 Covalent Bonds, Lewis Formulas, and the Octet Rule
The covalent, or shared electron pair, model of chemical bonding was suggested by G.N.
Lewis. He proposed that a sharing of two electrons by two hydrogen atoms permits each one
to have a stable closed-shell electron configuration analogous to that of helium.
The amount of energy required to dissociate a hydrogen molecule H 2 to two separate
hydrogen atoms is its bond dissociation enthalpy.
Only the electrons in an atom’s valence shell are involved in covalent bonding. Fluorine has
nine electrons, but only seven are in its valence shell. Pairing a valence electron of one
fluorine atom with one of a second fluorine gives a fluorine molecule (F 2) in which each
fluorine has eight valence electrons and an electron configuration equivalent to that of the
noble gas neon.
The six valence electrons of each fluorine that are not involved in bonding comprise three
unshared pairs (lone pairs)
Most elements obey the octet rule: in forming compounds they gain, lose, or share
electrons to achieve a stable electron configuration characterized by eight valence electrons.
Ethylene (C2H4) has 12 valence electrons, which can be distributed as follows:
,1.4 Polar Covalent Bonds, Electronegativity, and Bond Dipoles
Electrons in covalent bonds are not necessarily shared equally by the two atoms that they
connect.
If one atom has a greater tendency to attract electrons toward itself that the other, the
electron distribution is polarized, and the bond is
described as polar covalent. The tendency of an atom
to attract the electrons in a covalent bond toward itself
defines its electronegativity. An electronegative
element attracts electrons; an electropositive one
donates them.
Two ways of representing the polarization in HF are:
A third way of illustrating the electron polarization in HF is graphically, by way of an
electrostatic potential map, which uses the colors of the rainbow to show the charge
distribution. Blue through red tracks regions of greater positive charge to greater negative
charge.
The H-H bond is nonpolar, and a neutral yellow-green color dominates the electrostatic
potential map. Likewise, the F-F bond in F2 is nonpolar.
The covalent bond in HF unites two atoms of different electronegativity, and the electron
distribution is very polarized. Blue is the dominant color near the positively polarized
hydrogen, and red the dominant color near the negatively polarized fluorine.
Electronegativity increases from left to right across a row in the periodic table.
Of the second-row elements, the most electronegative is fluorine, the least electronegative is
lithium.
Electronegativity decreases going down a column.
Of the halogens, fluorine is the most electronegative.
In general, the greater the electronegativity difference between two elements, the more
polar the bond between them.
A dipole exists whenever opposite charges are separated from each other, and a dipole
moment μ is the product of the amount of the charge e multiplied by the distance d
between the centers of charge.
μ=e×d
The bond dipoles depend on the difference in electronegativity of the bonded atoms and on
the bond distance.
1.5 Formal Charge
Formal charges correspond to the difference between the number of valence electrons in
the neutral free atom and the number of valence electrons in its bonded state. The number
of electrons in the neutral free atom is the same as the atom’s group number in the periodic
table.
To determine the electron count of an atom in a Lewis formula, we add the total number of
electrons in unshared pairs to one-half the number of electrons in bonded pairs.
, Formal charge = Group number in periodic table – Electron count
Electron count =1/2(number of shared electrons) + Number of unshared electrons
With ions, the sums of the positive and negative formal charges will not be equal.
1.6 Structural Formulas of Organic Molecules: Isomers
Systematic Approach to writing Lewis Formulas
1. The molecular formula is determined experimentally
2. Based on the molecular formula, count the number of valence electrons
3. Given the connectivity, connect bonded atoms by a shared electron-par bond (:)
represented by a dash (-)
4. Count the number of electrons in the bonds (twice the number of bonds), and subtract
this from the total number of valence electrons to give the number of electrons that
remain to be added
5. Add electrons in pairs so that as many atoms as possible have eight electrons. It is
usually best to begin with the most electronegative atom. (No second-row element
have more than eight valence electrons)
6. If one or more atoms (excluding hydrogens) have fewer than eight electrons, use an
unshared pair from an adjacent atom to form a double or triple bond to complete the
octet. Use one double bond for each deficiency of two electrons to complete the octet
for each atom
7. Calculate formal charges.
In step 3 we set out a partial structure that shows the order in which the atoms are
connected. This is called the connectivity of the molecule. It frequently happens in organic
chemistry that two or more different compounds have the same molecular formula, but
different connectivities.
Different compounds that have the same molecular formula are classified as isomers.
Isomers can be either constitutional isomers (differ in connectivity) or
stereoisomers/structural isomers (differ in arrangement of atoms in space).
In a condensed formula, we omit the bonds altogether. Atoms and their attached
hydrogens are grouped and written in sequence; subscripts indicate the number of identical
groups attached to a particular atom. Bond-line formulas are formulas in which labels for
individual carbons are omitted and hydrogens attached to carbon are shown only when
necessary for clarity.
Heteroatoms are atoms other that carbon or hydrogen and are shown explicitly as are
hydrogens attached to them.
1.7 Resonance and Curved Arrows
Sometimes more than one Lewis formula can be written for a molecule, especially if the
molecule contains a double or triple bond. According to the resonance concept, when two
or more Lewis formulas that differ only in the distribution of electrons can be written for a
molecule, no single Lewis formula is sufficient to describe the true electron distribution. The
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