STRUCTURE AND BONDING Organic chemistry
Chapter 1 - STRUCTURE AND BONDING
Organic chemistry - the chemistry of CARBON Historically the "chemistry of compounds obtained from living organisms". "Vital force" Early nineteenth century: Michel Chevreul found soap could be broken down into several organic compounds he called "fatty acids" Animal fat Soap
NaOH HOH
Soap + Glycerine
H3O+
"Fatty acids"
Later Friedrich Wöhler found he could convert "inorganic salt" ammonium cyanate into "organic" urea simply with heat. O +
-
NH4 OCN
H2 N C NH2
Today many "organic" compounds still isolated from natural products but many synthesized. Why study to chemistry of one element? Number of compounds that contain carbon is much greater than of any other element. Carbon atoms can attach themselves to one another in chains, rings and branches off the chains and rings. Can be attached to other atoms also - mainly hydrogen but also O, S, N, HALOGENS, P, etc. Over 11 million ORGANIC COMPOUNDS. But not as overwhelming as it appears at first. The compounds can be classified into families by the reactive portion of the molecule - called "functional groups".
We will study these functional groups which are much fewer in number than 10 million. STRUCTURE: Review: Atomic Structure Atoms consist a dense nucleus containing protons (positively charges) and neutrons (neutral - no charge) and moving around the nucleus are the negatively charged electrons. The electrons have negligible mass compared to protons and neutrons. Most of the mass of the atom is contained in the nucleus. Early on the atom was described by Niels Bohr (Bohr atom). 8e 2e
Nuc
Nuc = nucleus with protons and neutrons surrounded by orbitals 1st holds two electrons maximum 2nd holds eight electrons maximum
Atom described by atomic number, Z, which gives the number of protons (as well as the number of electrons since the atom must be electronically neutral) and the mass number, A, which is the number of protons plus neutrons in the nucleus. The Bohr atom was a rather simple description since it does not really explain how the electrons are distributed around the nucleus except in terms of non distinct energy levels that could "hold" 2, 8, 18....... electrons. The proposed theory of QUANTUM MECHANICS (Schrödinger, Weisenberg and Dirac) helped explain electron distribution and bonding more precisely. This is a very complicated mathematical theory and will be explained in simple, nonmathematical terms. Motion of electrons around the nucleus is described using mathematical wave equations. A particular wave equation has solutions known as wave functions corresponding to different energy levels. The wave functions (REMEMBER MATHEMATICAL EXPRESSIONS) express NOT the exact position or speed of electron, but describe an area in space in which the electron is most likely to be found. Describes the PROBABILITY of finding an electron in a certain region. Since these are mathematical equations, they can be plotted to show a 3-D space. This area is called an ORBITAL.
ORBITAL: a region in space where an electron is most likely to be found. Review of ELECTRONIC CONFIGURATION: Primary energy levels or shells Similar idea to original idea of atomic theory (Bohr).
Shell 1 2 electrons Shell 2 8 electrons Shell 3 18 electrons
But developed further to state that within each of these shells there are discrete energy levels or orbitals KINDS OF ORBITALS SHAPE AND POSITION OF ORBITALS IN RELATION TO ATOMS NUCLEUS AND OTHER ORBITALS.
DEPENDS ON THE ENERGY OF THE ELECTRON. SHAPE DENOTED s, p , d and f. CONCERNED PRIMARILY WITH s and p ORBITALS IN ORGANIC CHEMISTRY.
Electrons exist in discrete energy levels: Lowest energy level is the 1s orbital (Remember that the numeral indicates the shell number in which the orbital appears. the lower case letter following indicates the shape of the orbital and should also give you an indication of how many of these orbitals are present). 1s orbital is spherical as are all s orbitals. Holds a maximum of 2 electrons as do all orbitals. Very small probability of finding electron very far out from the nucleus. The high probability of finding an electron in a region nearer the nucleus is referred to as ELECTRON DENSITY. Next energy level is the 2s level Again this orbital is spherical as are all s orbitals but is somewhat larger than the 1s and is a little further out from the nucleus.
This orbital also holds a maximum of 2 electrons. But the second shell can have 8 electrons. Each shell can have only one s orbital. Therefore there must be three other orbitals, different from an s orbital in this shell. There are 3 other orbitals, they are of equal energy to each other but of slightly higher energy than the 2s orbitals. They also have a different shape than an s orbital. These are known as the p orbitals. These p orbitals are dumbbell-shaped, and oriented on three mutually perpendicular axes: px, py, pz.
y x
y x
px z
y x
py z
pz
The third shell can have a maximum of 18 electrons; so in addition to the 3s and the three 3p's, there are five d orbitals called the 3d orbitals. Generally not concerned with these in organic chemistry but understand that they exist from row three and above in the Periodic Table and these are spaces (orbitals) available for more than an octet of electrons around these atoms.
_____ _______ ______
3d
_____ _____ _____ _____ _____
4s
______
3p
_____ _____ _____
3s
_____
2p
_____ _____ _____
2s
______
1s
______
px
py
pz
five d's Energy
4p
three p's
ELECTRONIC CONFIGURATION 1. Aufbau Principle - electrons enter orbitals of the lowest energy available first. 2. Pauli Exclusion Principle - only two electrons in an orbital - must have opposite spins (must be spin "paired"). Electrons of like spin tend to get as far away from each other as possible. 3. Hund's Rule - if orbitals of the same energy are available, one electron goes in each (of parallel spin) before pairing occurs. Samples of electronic configurations: Ground State Confirgurations: C = 1s2 , 2s2 , 2p2 (4 valence electrons) F = 1s2 , 2s2 , 5p2 (7 valence electrons) N = 1s2 , 2s2 , 3p3 (5 valence electrons) Na = 1s2 , 2s2 , 6p2 , 3s1 (1 valence electrons) Development of Chemical Bonding 1858 - Kekule and Cooper - Carbon has four (4) affinity units, i.e. forms four bonds. Carbon is "tetravalent". Can form chains (rows) of carbon connected to each other. Erlenmeyer proposed that carbon could form a triple bond C=C.
Brown proposed that carbon could form a double bond C=C. Kekule later proposed that carbon could form rings also (in addition to chains). 1874 - van'tHoff and LeBel said that carbon was three-dimensional - that is the four bonds of carbon did not lie in a plane but were directed toward the corners of a tetrahedron. CHEMICAL BONDING IONIC - For elements in row two and above, a complete octet in outer shell is stable - configuration of the inert gases (noble gases). Atoms tend to achieve this stability. Alkali metals tend to lose one electron - low ionization energy - easy to lose one electron. They are ELECTROPOSITIVE, Alkaline earths tend to lose two electron. Again this is easier to do than to gain 6 electrons to form a complete octet in the out shell. Halogens, on the other hand, easily gain one electron to complete their outer shell. They are ELECTRONEGATIVE. IONIC BONDS are formed between elements of very different electronegativities. Elements from groups of the left of the periodic table tend to form ionic bonds with elements on the right of the table. COVALENT BONDS are formed by sharing of electrons - that is electrons are neither fully gained nor lost by the atoms involved in the bonds. Elements in the center of the table tend to share electrons and form covalent bonds. Bonds between atoms actually occur along a contimuum between a pure ionic bond (complete transfer) and a pure covalent bond (equal sharing). Ionic Bond
Na+F-
Polar Bond
H-F
Covalent Bond
H-H F-F
We have a simple shorthand for representing structures with covalent bonds on paper. These are LEWIS DOT STRUCTURES. It is a way of bookkeeping valence electrons. Another simpler method of representation are KEKULE or LINE FORMULAS. Lewis dot structures show all valence electrons including nonbonding or lone pair electrons. Kekule formulas do not necessarily show nonbonding electrons unless for emphasis of some kind. LEWIS DOT STRUCTURES
Add up valence electrons for all atoms in the formula (determined from the group on the periodic chart). Connect atoms in most symmetrical way or as described. In organic chem carboncarbon bonds are important. Each bond has two electrons. Place "bond" between all atoms. Put rest of electrons around outer elements so that each has an octet. If central atom does not have an octet, make multiple (double or triple) bond. Make sure there are the correct number of electrons total. If a cation, less one electron; if an anion, plus one electron. Remember number of bonds that some common elements have in neutral state C - 4, H - 1, X - 1, O - 2 (plus two lone pairs), same S, N - 3 (plus lone pair), same P (but remember d orbitals, Easier to draw Kekule or line structures. CH4 C - four valence electrons each H - one valence electron
C
+
BH3 B - three valence electrons each H - one valence electron
H H C H H
4H H BH H
H H C NH H H These, however, do not tell us about the shape of the molecule or where electrons may be in relation to the two atoms involved in the bond. Have to go to molecular orbital theory for a more complete picture. CH3NH 2
MOLECULAR ORBITALS - Describes covalent bonding in organic compounds in electronic terms. In molecules, electrons occupy orbitals. In reality the molecular orbitals that hold electrons are centered around many nuclei - perhaps the whole molecule. Make assumptions to better understand the picture. A pair of electrons in an m.o. are localized around two nuclei.
Shape and position of m.o. related to shape and position of the atomic orbitals from which the m.o. was formed. "Bond orbitals" hold a maximum of 2 electrons. M.O. formed from the OVERLAP of atomic orbitals. Bonds are "DIRECTIONAL" to obtain maximum overlap. New orbital called a MOLECULAR ORBITAL - associated with molecule rather than individual atoms.
Formation of hydrogen molecule:
H
+
H
H
H
H H
This is called a sigma (s) orbital - symmetrical about the bond axis.
New arrangement of electrons is considerably more stable than the individual atoms. (Paired electrons - complete outer shell). When bond forms, energy given off (BOND STRENGTH). In H2, = 104 kcal/mol (436 kJ/mol). It would take this same amount of energy to break this bond = BOND DISSOCIATION ENERGY. BOND LENGTH - distance between nuclei - is optimum. In H2 is 0.74 Å. 1Å = 10-8 cm = 10-10 m FOR EVERY ORBITAL WE START WITH, MUST GET THE SAME NUMBER OF ORBITALS AT THE END. True no matter what we do to orbitals. In formation of hydrogen bond, started out with two atomic orbitals when we formed the bond. Since energy given off in bond formation, the molecular orbital in which we now find the electrons is of lower energy than the original atomic orbitals. Also form a molecular orbital that is equally higher in energy (conservation of
energy) but is not occupied = no electrons in it. Bonding and antibonding orbitals are formed. H-H antibonding mol. orb.
Hydrogen 1s at. orb.
Hydrogen 1s at. orb.
Both electrons are in the bonding molecular orbital and are spin-paired. The antibonding orbital is empty. Total energy in m.o.s is the same as energy in a.o.s
Linus Pauling and the Principle of Maximum Overlap s + s = sigma bond s bond
p + p (head-on overlap) = sigma bond s p + p (sideways) = pi bond π bond
+
Nodal plane at bond axis
BONDING IN ORGANIC MOLECULES IS MADE UP ALMOST EXCLUSIVELY BY A COMBINATION OF SIGMA AND PI BONDS
HYBRIDIZATION: Can use valence shell electron pair repulsion (VSEPR) theory to explain the shapes of orbitals around individual atoms and positions of electrons. Since attached atoms bond to the central atom via these orbitals, the shapes of molecules can usually be predicted by VSEPR. But this and orbital overlap does not fully explain some other properties about bonding that we know as fact. C + 2H (ground state)
H C H + ~200kcal/mol energy released ~ 100 kcal/mol each C-H bond formed
But we know that carbon forms four bonds with hydrogen. Carbon can also exist in excited state (excite one 2s electron to the empty 2pz orbital). Now have four unpaired electrons in the second shell available for bonding. H H C H H
96kcal/mol
C
C + 4H
+ ~400 kcal/mole
Net gain from this explanation ≈ 200kcal/mol.
2p
2p 96 kcal/mol
2s 1s
Valence shell
2s 1s
ground state electronic configuration
excited state electronic configuration
But we also know that all four bonds in methane are equal in bond strength and bond length. Cannot be explained if the bonding is through three p orbitals and one s orbital. There must be some explanation that makes the four orbitals on the carbon the same. We explain this with the theory of HYBRIDIZATION. - a mathematical mixing of the wave functions of the orbitals. Explain sp3 hybridization in METHANE.................
One s plus three p = four orbitals; so must get four hybrid orbitals: 4 sp3 orbitals. The tetrahedral caarbon. Shape of sp3: Overlap of s and p orbital -------> changed shape. Remember is a mathematical mix of the wave functions. +
+
-
These are mathematical signs –NOT charges.
-
+
Positive plus positive end gets big. Positive (s) plus negative (p) gets small.
There are FOUR of these orbitals and they are directed toward the corners of a tetrahedron.
MODELS Molecular orbital diagrams of formation of methane from one hybridized carbon and four hydrogens. s-sp3 antibonding m.o.
sp3
sp3
sp3
sp3
carbon atom
1s 1s 1s four hydrogn atoms
1s
s-sp3 bonding m.o.
Show pictures of ETHANE. (To compare later to ethene and ethyne). Explain sp2 hybridization for BCl3 and ethene. Explain sp hybridization for BeCl2 and ethyne. SAME KINDS OF HYBRIDIZATION CAN BE SHOWN FOR OTHER ATOMS NITROGEN, OXYGEN, SUFLUR, ETC.
Organic chemistry - the chemistry of CARBON Historically the "chemistry of compounds obtained from living organisms". "Vital force" Early nineteenth century: Michel Chevreul found soap could be broken down into several organic compounds he called "fatty acids" Animal fat Soap
NaOH HOH
Soap + Glycerine
H3O+
"Fatty acids"
Later Friedrich Wöhler found he could convert "inorganic salt" ammonium cyanate into "organic" urea simply with heat. O +
-
NH4 OCN
H2 N C NH2
Today many "organic" compounds still isolated from natural products but many synthesized. Why study to chemistry of one element? Number of compounds that contain carbon is much greater than of any other element. Carbon atoms can attach themselves to one another in chains, rings and branches off the chains and rings. Can be attached to other atoms also - mainly hydrogen but also O, S, N, HALOGENS, P, etc. Over 11 million ORGANIC COMPOUNDS. But not as overwhelming as it appears at first. The compounds can be classified into families by the reactive portion of the molecule - called "functional groups".
We will study these functional groups which are much fewer in number than 10 million. STRUCTURE: Review: Atomic Structure Atoms consist a dense nucleus containing protons (positively charges) and neutrons (neutral - no charge) and moving around the nucleus are the negatively charged electrons. The electrons have negligible mass compared to protons and neutrons. Most of the mass of the atom is contained in the nucleus. Early on the atom was described by Niels Bohr (Bohr atom). 8e 2e
Nuc
Nuc = nucleus with protons and neutrons surrounded by orbitals 1st holds two electrons maximum 2nd holds eight electrons maximum
Atom described by atomic number, Z, which gives the number of protons (as well as the number of electrons since the atom must be electronically neutral) and the mass number, A, which is the number of protons plus neutrons in the nucleus. The Bohr atom was a rather simple description since it does not really explain how the electrons are distributed around the nucleus except in terms of non distinct energy levels that could "hold" 2, 8, 18....... electrons. The proposed theory of QUANTUM MECHANICS (Schrödinger, Weisenberg and Dirac) helped explain electron distribution and bonding more precisely. This is a very complicated mathematical theory and will be explained in simple, nonmathematical terms. Motion of electrons around the nucleus is described using mathematical wave equations. A particular wave equation has solutions known as wave functions corresponding to different energy levels. The wave functions (REMEMBER MATHEMATICAL EXPRESSIONS) express NOT the exact position or speed of electron, but describe an area in space in which the electron is most likely to be found. Describes the PROBABILITY of finding an electron in a certain region. Since these are mathematical equations, they can be plotted to show a 3-D space. This area is called an ORBITAL.
ORBITAL: a region in space where an electron is most likely to be found. Review of ELECTRONIC CONFIGURATION: Primary energy levels or shells Similar idea to original idea of atomic theory (Bohr).
Shell 1 2 electrons Shell 2 8 electrons Shell 3 18 electrons
But developed further to state that within each of these shells there are discrete energy levels or orbitals KINDS OF ORBITALS SHAPE AND POSITION OF ORBITALS IN RELATION TO ATOMS NUCLEUS AND OTHER ORBITALS.
DEPENDS ON THE ENERGY OF THE ELECTRON. SHAPE DENOTED s, p , d and f. CONCERNED PRIMARILY WITH s and p ORBITALS IN ORGANIC CHEMISTRY.
Electrons exist in discrete energy levels: Lowest energy level is the 1s orbital (Remember that the numeral indicates the shell number in which the orbital appears. the lower case letter following indicates the shape of the orbital and should also give you an indication of how many of these orbitals are present). 1s orbital is spherical as are all s orbitals. Holds a maximum of 2 electrons as do all orbitals. Very small probability of finding electron very far out from the nucleus. The high probability of finding an electron in a region nearer the nucleus is referred to as ELECTRON DENSITY. Next energy level is the 2s level Again this orbital is spherical as are all s orbitals but is somewhat larger than the 1s and is a little further out from the nucleus.
This orbital also holds a maximum of 2 electrons. But the second shell can have 8 electrons. Each shell can have only one s orbital. Therefore there must be three other orbitals, different from an s orbital in this shell. There are 3 other orbitals, they are of equal energy to each other but of slightly higher energy than the 2s orbitals. They also have a different shape than an s orbital. These are known as the p orbitals. These p orbitals are dumbbell-shaped, and oriented on three mutually perpendicular axes: px, py, pz.
y x
y x
px z
y x
py z
pz
The third shell can have a maximum of 18 electrons; so in addition to the 3s and the three 3p's, there are five d orbitals called the 3d orbitals. Generally not concerned with these in organic chemistry but understand that they exist from row three and above in the Periodic Table and these are spaces (orbitals) available for more than an octet of electrons around these atoms.
_____ _______ ______
3d
_____ _____ _____ _____ _____
4s
______
3p
_____ _____ _____
3s
_____
2p
_____ _____ _____
2s
______
1s
______
px
py
pz
five d's Energy
4p
three p's
ELECTRONIC CONFIGURATION 1. Aufbau Principle - electrons enter orbitals of the lowest energy available first. 2. Pauli Exclusion Principle - only two electrons in an orbital - must have opposite spins (must be spin "paired"). Electrons of like spin tend to get as far away from each other as possible. 3. Hund's Rule - if orbitals of the same energy are available, one electron goes in each (of parallel spin) before pairing occurs. Samples of electronic configurations: Ground State Confirgurations: C = 1s2 , 2s2 , 2p2 (4 valence electrons) F = 1s2 , 2s2 , 5p2 (7 valence electrons) N = 1s2 , 2s2 , 3p3 (5 valence electrons) Na = 1s2 , 2s2 , 6p2 , 3s1 (1 valence electrons) Development of Chemical Bonding 1858 - Kekule and Cooper - Carbon has four (4) affinity units, i.e. forms four bonds. Carbon is "tetravalent". Can form chains (rows) of carbon connected to each other. Erlenmeyer proposed that carbon could form a triple bond C=C.
Brown proposed that carbon could form a double bond C=C. Kekule later proposed that carbon could form rings also (in addition to chains). 1874 - van'tHoff and LeBel said that carbon was three-dimensional - that is the four bonds of carbon did not lie in a plane but were directed toward the corners of a tetrahedron. CHEMICAL BONDING IONIC - For elements in row two and above, a complete octet in outer shell is stable - configuration of the inert gases (noble gases). Atoms tend to achieve this stability. Alkali metals tend to lose one electron - low ionization energy - easy to lose one electron. They are ELECTROPOSITIVE, Alkaline earths tend to lose two electron. Again this is easier to do than to gain 6 electrons to form a complete octet in the out shell. Halogens, on the other hand, easily gain one electron to complete their outer shell. They are ELECTRONEGATIVE. IONIC BONDS are formed between elements of very different electronegativities. Elements from groups of the left of the periodic table tend to form ionic bonds with elements on the right of the table. COVALENT BONDS are formed by sharing of electrons - that is electrons are neither fully gained nor lost by the atoms involved in the bonds. Elements in the center of the table tend to share electrons and form covalent bonds. Bonds between atoms actually occur along a contimuum between a pure ionic bond (complete transfer) and a pure covalent bond (equal sharing). Ionic Bond
Na+F-
Polar Bond
H-F
Covalent Bond
H-H F-F
We have a simple shorthand for representing structures with covalent bonds on paper. These are LEWIS DOT STRUCTURES. It is a way of bookkeeping valence electrons. Another simpler method of representation are KEKULE or LINE FORMULAS. Lewis dot structures show all valence electrons including nonbonding or lone pair electrons. Kekule formulas do not necessarily show nonbonding electrons unless for emphasis of some kind. LEWIS DOT STRUCTURES
Add up valence electrons for all atoms in the formula (determined from the group on the periodic chart). Connect atoms in most symmetrical way or as described. In organic chem carboncarbon bonds are important. Each bond has two electrons. Place "bond" between all atoms. Put rest of electrons around outer elements so that each has an octet. If central atom does not have an octet, make multiple (double or triple) bond. Make sure there are the correct number of electrons total. If a cation, less one electron; if an anion, plus one electron. Remember number of bonds that some common elements have in neutral state C - 4, H - 1, X - 1, O - 2 (plus two lone pairs), same S, N - 3 (plus lone pair), same P (but remember d orbitals, Easier to draw Kekule or line structures. CH4 C - four valence electrons each H - one valence electron
C
+
BH3 B - three valence electrons each H - one valence electron
H H C H H
4H H BH H
H H C NH H H These, however, do not tell us about the shape of the molecule or where electrons may be in relation to the two atoms involved in the bond. Have to go to molecular orbital theory for a more complete picture. CH3NH 2
MOLECULAR ORBITALS - Describes covalent bonding in organic compounds in electronic terms. In molecules, electrons occupy orbitals. In reality the molecular orbitals that hold electrons are centered around many nuclei - perhaps the whole molecule. Make assumptions to better understand the picture. A pair of electrons in an m.o. are localized around two nuclei.
Shape and position of m.o. related to shape and position of the atomic orbitals from which the m.o. was formed. "Bond orbitals" hold a maximum of 2 electrons. M.O. formed from the OVERLAP of atomic orbitals. Bonds are "DIRECTIONAL" to obtain maximum overlap. New orbital called a MOLECULAR ORBITAL - associated with molecule rather than individual atoms.
Formation of hydrogen molecule:
H
+
H
H
H
H H
This is called a sigma (s) orbital - symmetrical about the bond axis.
New arrangement of electrons is considerably more stable than the individual atoms. (Paired electrons - complete outer shell). When bond forms, energy given off (BOND STRENGTH). In H2, = 104 kcal/mol (436 kJ/mol). It would take this same amount of energy to break this bond = BOND DISSOCIATION ENERGY. BOND LENGTH - distance between nuclei - is optimum. In H2 is 0.74 Å. 1Å = 10-8 cm = 10-10 m FOR EVERY ORBITAL WE START WITH, MUST GET THE SAME NUMBER OF ORBITALS AT THE END. True no matter what we do to orbitals. In formation of hydrogen bond, started out with two atomic orbitals when we formed the bond. Since energy given off in bond formation, the molecular orbital in which we now find the electrons is of lower energy than the original atomic orbitals. Also form a molecular orbital that is equally higher in energy (conservation of
energy) but is not occupied = no electrons in it. Bonding and antibonding orbitals are formed. H-H antibonding mol. orb.
Hydrogen 1s at. orb.
Hydrogen 1s at. orb.
Both electrons are in the bonding molecular orbital and are spin-paired. The antibonding orbital is empty. Total energy in m.o.s is the same as energy in a.o.s
Linus Pauling and the Principle of Maximum Overlap s + s = sigma bond s bond
p + p (head-on overlap) = sigma bond s p + p (sideways) = pi bond π bond
+
Nodal plane at bond axis
BONDING IN ORGANIC MOLECULES IS MADE UP ALMOST EXCLUSIVELY BY A COMBINATION OF SIGMA AND PI BONDS
HYBRIDIZATION: Can use valence shell electron pair repulsion (VSEPR) theory to explain the shapes of orbitals around individual atoms and positions of electrons. Since attached atoms bond to the central atom via these orbitals, the shapes of molecules can usually be predicted by VSEPR. But this and orbital overlap does not fully explain some other properties about bonding that we know as fact. C + 2H (ground state)
H C H + ~200kcal/mol energy released ~ 100 kcal/mol each C-H bond formed
But we know that carbon forms four bonds with hydrogen. Carbon can also exist in excited state (excite one 2s electron to the empty 2pz orbital). Now have four unpaired electrons in the second shell available for bonding. H H C H H
96kcal/mol
C
C + 4H
+ ~400 kcal/mole
Net gain from this explanation ≈ 200kcal/mol.
2p
2p 96 kcal/mol
2s 1s
Valence shell
2s 1s
ground state electronic configuration
excited state electronic configuration
But we also know that all four bonds in methane are equal in bond strength and bond length. Cannot be explained if the bonding is through three p orbitals and one s orbital. There must be some explanation that makes the four orbitals on the carbon the same. We explain this with the theory of HYBRIDIZATION. - a mathematical mixing of the wave functions of the orbitals. Explain sp3 hybridization in METHANE.................
One s plus three p = four orbitals; so must get four hybrid orbitals: 4 sp3 orbitals. The tetrahedral caarbon. Shape of sp3: Overlap of s and p orbital -------> changed shape. Remember is a mathematical mix of the wave functions. +
+
-
These are mathematical signs –NOT charges.
-
+
Positive plus positive end gets big. Positive (s) plus negative (p) gets small.
There are FOUR of these orbitals and they are directed toward the corners of a tetrahedron.
MODELS Molecular orbital diagrams of formation of methane from one hybridized carbon and four hydrogens. s-sp3 antibonding m.o.
sp3
sp3
sp3
sp3
carbon atom
1s 1s 1s four hydrogn atoms
1s
s-sp3 bonding m.o.
Show pictures of ETHANE. (To compare later to ethene and ethyne). Explain sp2 hybridization for BCl3 and ethene. Explain sp hybridization for BeCl2 and ethyne. SAME KINDS OF HYBRIDIZATION CAN BE SHOWN FOR OTHER ATOMS NITROGEN, OXYGEN, SUFLUR, ETC.
