chem 220 notes page 1 structure and bonding of organic ... · patterns for the formal charges of...

Chem 220 Notes

Lecture Notes © 2003-2015 Dr. Thomas Mucciaro. All rights reserved.

Structure and Bonding of Organic Molecules

I. Types of Chemical Bonds

A. Why do atoms forms bonds?

� Atoms want to have the same number of electrons as the nearest noble gas atom (noble gas configuration). This requires having a completely full or completely empty valence shell of electrons.

� Most main groups atoms will try to have eight valence electrons to completely fill their valence shell; this is the octet rule. There are some exceptions: hydrogen only needs two electrons (duet rule), and boron and aluminum are stable with only six valence electrons (incomplete octet).

� Atoms form bonds in order to gain or lose electrons and satisfy the octet rule.

B. Two Types of Bonds

� Atoms have two major strategies for gaining (or losing) electrons to fulfill the octet rule. These two different strategies lead to the two main types of chemical bonds.

1. Ionic Bonds

� One strategy is to completely transfer valence electrons between atoms:

� This forms oppositely charged ions, which are then attracted and form ionic bonds.

2. Covalent Bonds

� The other strategy is to share valence electrons between atoms:

� This strategy forces the atoms to remain in close proximity to continue sharing electrons.

� This forms covalent bonds (co = together and valent = valence).

Na Cl ClNa Oppositely charged ions attract (ionic bond)

Cl Cl ClCl Atoms must stay together to share valenceelectons (covalent bond)

Chem 220 Notes Structure and Bonding of Organic Molecules


C. The type of bond formed depends on the electronegativity difference of the atoms

� Electronegativity: the strength with which an atom attracts electrons. Electronegativity is a periodic property which increases from left to right and from bottom to top of the periodic table:

� Atoms with an electronegativity difference larger than about 2.0 form ionic bonds.This usually happens between atoms from opposite sides of the periodic table (between metals and non-metals).

II. Lewis Structures

� Lewis structures show the connections between atoms that form covalent bonds. They are created by connecting atoms so that they share electrons to have octets.

� In explicit Lewis structures, all atoms are shown using their atomic symbols, all covalent bonds are shown as dashes, and all unshared electrons are shown as dots.

� To be correct, a Lewis structure must have the correct number of valence electrons, and in most cases should not violate the octet rule.

� There are many ways to predict Lewis structures from the formula of a compound.One method will be presented in Chem 20.

electronegativity increases

H2.1

Li1.0

Be1.6

B2.0

C2.5

N3.0

O3.5

F4.0

Na0.9

Mg1.2

Al1.5

Si1.8

P2.1

S2.5

Cl3.0

K0.8

Br2.8

I2.4



III. Formal Charges

Formal charges must be calculated for each atom of a given Lewis structure. There are patterns for the formal charges of the important atoms in organic chemistry, and eventually we will begin to predict them without calculating.

To calculate the formal charge on an atom in a Lewis structure:

1. Draw a complete Lewis structure of the compound of interest. Include all lone pairs and unpaired electrons.

2. For each atom, determine the number of valence electron that the atom “wants” to own. This will be equal to the group number of that atom in the periodic table, or the number of valence electrons on the neutral atom.

3. Determine the number of valence electrons that the atom actually “owns” in the Lewis structure:

� An atom owns two electron for each of its lone pairs.

� An atom owns one electron for each bond connected to it (a double bond counts as two bonds; a triple bond counts as three bonds).

� An atom owns one electron for each unpaired electron.

Using these rules:

[# of electrons “owned”] = [2 x (# of lone pairs)] + [1 x (# of bonds)] + [1 x (# of unpaired electrons)]

4. The formal charge is calculated by subtracting the number of electrons that the atom “owns” from the number of electrons that the atom “wants” (remember: “wants” minus “owns”):

Formal Charge = [# of electrons atom “wants”] – [# of electrons atom “owns”]

5. Write the formal charge, in a circle, next to the atom with the charge. This procedure is repeated for each atom in the Lewis structure.

6. The overall charge on the structure will equal the sum of all the formal charges on the individual atoms. When the overall charge is already known, this is a good way to check that all of the formal charges have been calculated correctly.



� Some Common Patterns for Formal Charges

IV.Octet Rule Violators

Atoms with Open Shells (Incomplete Octets)

� If we work out the Lewis structure of BF3 using the 6N + 2 method, we might conclude that there is a multiple bond in the structure as shown below:

� The actual most likely structure is shown above. It has all single bonds between boron and fluorine, and only six electron pairs (three bonds) around the central boron atom.In this case, the boron has an incomplete octet (an open shell).

� The second structure is considered a better representation because it has fewer formal charges on the atoms in the molecule, which implies that the atoms are closer to having their desired number of electrons.

� In general, atoms in Group III (third column of the periodic table), such as boron and aluminum, will have an open shell when there is a = 2 in the 6N + 2 calculation. This will allow the structure to have the fewest formal charges.

Formal Charge Carbon Nitrogen Oxygen

Neutral(no charge)

+1 Charge

–1 Charge

BF3 Proposed Structure Actual Most Likely Structure

valence electrons = 24

6N + 2 = 26

= 2 (one multiple bond?)

CN O

C N O

C N O

F B F

F

F B F

F



Atoms with Expanded Octets

� If we work out the Lewis structure of SO42– using the 6N + 2 method, we would

conclude that the structure has all single bonds and formal charges on adjacent sulfur and oxygen atoms:

� The actual most commonly preferred structure has double bonds between the sulfur and two of the oxygens. This puts more than eight electrons around the sulfur atom, which gives it an expanded octet. The other two oxygens still have one bond to sulfur, three lone pairs, and a –1 formal charge.

� The second structure is commonly preferred because it minimizes the opposite formal charges on adjacent atoms. The violation of the octet rule on sulfur is allowed because it is a third row atom. Atoms in the third row or higher have extra orbitals (d orbitals) available to hold extra electrons, which permits them to have expanded octets.

� Expanded octets are generally used when a third row (or higher) atom with a formal charge is adjacent to an atom with the opposite formal charge. In that case, remove a lone pair from the negatively charged atom and draw a multiple bond between the atoms. This will form an expanded octet, and will reduce the number of formal charges in the structure.

� First and second row atoms (including carbon, nitrogen, and oxygen) never have expanded octets!

SO42– Proposed Structure Actual Most Likely Structure

valence electrons = 32

6N + 2 = 32

= 0 (no multiple bonds)

O S O

O

O

2+

O S O

O

O



V. Constitutional Isomers

� Isomers are molecules with the same molecular formula that have different structures. We will discuss several different types of isomers in this course. Most involve differences in three-dimensional structures (which we will look at in Chapters 3 and 5).

� Constitutional isomers are molecules with the same formula but different connectivity between atoms.

� Connectivity means the order that atoms are connected in the molecule. Moleculeshave different connectivity when they have different Lewis structure skeletons(ignoring multiple bonds and lone pairs).

� Constitutional isomers are also called structural isomers. They are so common that many chemists will use the term isomers to indicate constitutional isomers unless another type of isomer is explicitly indicated.

� Lewis structures do not show three-dimensions. They only really indicate the connectivity between atoms. Therefore, changing the orientation of a Lewis structure does not create an isomer. For example, the following two Lewis structures represent the same compound (they are not constitutional isomers).



VI.Classification of Carbons

� When we start to examine constitutional isomers, we can see that often different carbons in a compound can have different numbers of other carbon atoms directly attached.

� We will distinguish betweens these different types or classes of carbons based on the number of carbons atoms directly attached. These classifications apply only to carbons that are not part of a multiple bond (only have single bonds attached).

� Later we will see that different classes of carbons can have different reactivity and chemical behavior (even in the same compound).

VII.Drawing Chemical Structures

� There are three widely used alternatives to Lewis structures. The primary advantages of these structures is that they save time and space. They all have the disadvantage that they are less explicit than a Lewis structure, which can at times be confusing and cause us to make unintentional errors.

A. Partially Condensed Structures

� In partially condensed structures, the bonds between carbon and hydrogen are not drawn out. Instead, the hydrogens attached to a given carbon are listed after that carbon in the structure. An example is shown below:

# Carbon Atoms Directly Attached

0 1 2 3 4

Classification methylprimary

(1o)secondary

(2o)tertiary

(3o)quaternary

(4o)

H C C C C

H

H

H

H

H

H

H

H

H

CH3 CH2 CH2 CH3becomes

Lewis Structure Partially Condensed Structure



� The bonds between carbons and to other atoms are still explicitly shown. This makes it easy to see the structure of the carbon skeleton and any branches.

� This is a convenient way to save time and space when writing structures, and we will use it most of the time.

B. Fully Condensed Structures

� In fully condensed structures, all bond dashes are omitted. All the atoms or groups attached to a carbon are listed immediately after that carbon. Complicated, polyatomic groups and branches are written in parentheses to show that they are together. An example is shown below:

� Fully condensed structures save a great deal of space, and they are easy to type out.However, they are most useful only for relatively simple structures, and they are very confusing when the carbon chain has complicated branches.

� This type of structure will only be used for simple molecules with few branches.

C. Bond-Line Structures

� In bond-line structures, all of the carbon and hydrogen atoms are omitted, and only the bonds between carbons are drawn. All the vertices represent carbon atoms, and they are assumed to be bonded to enough hydrogens to complete their octet. Also, the end of any line that is not attached to a written out atom is also a carbon.

� Bond-line structures focus attention on the shape of the carbon skeleton, and they also save time and space. They are used especially for compounds that contain rings or complicated carbon skeletons. We will use these mostly for rings, and use partially condensed structures for most other compounds.

H C C C C

H

H

H

H

H

H

H

H

H

becomes

Lewis Structure

CH3CH2CH2CH3

Fully Condensed Structure

H C C C C

H

H

H

H

H

H

H

H

H

becomes

Lewis Structure Bond-Line Structure



IX.Resonance Structures

� The Lewis structure model works for many molecules, but it needs to be extended to adequately describe some molecules and ions.

� Resonance structures occur when there are several different possible arrangements of multiple bonds and lone pairs on a given skeleton for a given formula.

For example, there are two possible structures for ozone (O3), both of which have the same underlying skeleton, but different locations of the multiple bond and lone pairs (and consequently the formal charges).

Neither structure can be called more correct; they both seem valid for the molecule.Therefore, we draw both structures and relate them with the special double-headed arrow shown above. These two possible structures are called resonance structures, and the double-headed arrow is called a resonance arrow.

A. Significance of Resonance Structures

� There may be more than two resonance structures for a molecule.

� None of the individual resonance structures completely describe the actual molecule; instead the actual molecule is viewed as a mixture or hybrid of the individual structures.

� The resonance structures are not in equilibrium with each other. Each resonance structure is a partial, incomplete picture of the actual molecule. The actual molecule has one structure that can’t be drawn accurately using the Lewis structure rules.

� The underlying skeleton does not change between resonance structures. Changing the skeleton involves moving atoms in the structure. The only things that change (or move) between resonance structures are multiple bonds and lone pairs.

� The actual molecule will have a lower energy (be more stable) than any of the individual resonance structures. One important feature of resonance is that it adds stability to a molecule when present.

O O O OOO



B. Relative Stability and Contribution of Resonance Structures

� Not all resonance structures contribute equally to the overall hybrid of the molecule.Instead, each will contribute a certain percent amount to the overall hybrid.

� In general, the more stable resonance structures will make a greater contribution to the hybrid.

� The order of stability of resonance structures depends on several factors, which are listed below (most stable to least stable):

1. Structures with the most bonds tend to be more stable than structures with fewer bonds.

2. Octet rule followers tend to be more stable than structures with incomplete octets.

3. Structures with fewer numbers of formal charges on atoms tend to be more stable.

4. Structures with positive (+) formal charges on atoms with low electronegativity and negative (–) formal charges on atoms with high electronegativity are more stable.

5. Structures with negative (–) formal charges on atoms with low electronegativity and positive (+) formal charges on atoms with high electronegativity are very unstable and are disfavored. Usually they are not written because they make such a small contribution. These structures can be said to violate the natural polarity of the bond.

C. Pushing Electrons to Predict Resonance Structures

� We can use curved arrows to show how electrons in multiple bonds ( bonds) and lone pairs could be rearranged to form additional resonance structures from a given Lewis structure. This method is called electron pushing.

� The movements of electrons in this method are imaginary. Electrons don’t actually move in the real structure, and resonance structures are not in equilibrium.Instead, the real structure is a mixture or hybrid of the resonance structures. We use resonance structures because the real hybrid structure can’t be correctly represented by Lewis structure rules.



Curved-Arrow Notation� We use curved-arrow notation to represent the movement of electrons in Lewis

structures.

� This electron movement may be imaginary (or conceptual)—like in resonance structures—or it can be real movement that occurs during chemical reactions.

� A simple example of curved-arrow notation is shown below. There is one key concept in curved-arrow notation: a curved-arrow always starts at the electrons and points to where they are moving.

� A double-headed arrow is used to show the movement of two electrons (a pair of electrons). Later, we will see single-headed “fishhook” arrows that are used for movement of single electrons.

Pushing Electrons to Generate Resonance Structures� There are five basic patterns of electron movement that are seen in resonance.

1. Push lone pairs into adjacent incomplete octets



2. Push multiple bonds ( bonds) toward adjacent incomplete octets

3. Push lone pairs into adjacent multiple bonds ( bonds)

4. Pull electrons from a multiple bond ( bond) onto an electronegative atom

5. Push multiple bonds ( bonds) into adjacent multiple bonds ( bonds) around a ring



X. Atomic Orbitals

A. What are atomic orbitals?

� Atomic orbitals are areas in three-dimensional space around an atom where electrons with a specific energy are allowed to exist.

� Atomic orbitals are defined by special mathematical functions called wavefunctions.Wavefunctions are often symbolized as (x, y, z).

� When the (x, y, z) coordinates are substituted into the wavefunction, a number will be calculated. We can plot the wavefunction on a three-dimensional (x, y, z) coordinate system by using different size dots to represent different size numbers. An example is shown below for the 1s wavefunction.

B. Significance of Wavefunctions

� Wavefunctions relate the energy of an electron to its location in space around an atom. There is a one-to-one correspondence between the energy of an atoms electrons and their location in space.

� The numbers that a wavefunction (x, y, z) produces relate to the probability of finding an electron with the given energy at that (x, y, z) location in space. Specifically,

2(x, y, z) gives the probability of finding the electron at that location. A larger number indicates a higher probability.

� Each different allowed energy value has different, unique function. These functions define the different atomic orbitals for an atom. If the energy of a given electron changes, then it will be described by a different function and will move to a different area of space.



C. Types of Atomic Orbitals

� The shapes of the two most important types of atomic orbitals are shown below:

� The s orbitals are spherical, with the center on the nucleus of the atom. The numbers that are produced by (x, y, z) for an s orbital are all positive numbers.

� The p orbitals have two lobes that exist on opposite sides of the nucleus. The (x, y, z) produces positive numbers for (x, y, z) coordinates in one lobe, and negative numbers for the other lobe. These mathematical signs are indicated by shading the lobes in opposite colors. The sign of the function does not indicate the electrical charge.

� There are always three different p orbitals with the same energy. Orbitals with the same energy are called degenerate orbitals.

� A given function can only describe a maximum of two electrons at a time. In other words, a given atomic orbital can only hold a maximum of two electrons. This is known as the Pauli Exclusion Principle.



XI.Molecular Orbital Theory

A. Forming Bonds with Orbitals (Orbital Overlap)

� In Lewis bonding theory, atoms shared valence electrons to form covalent bonds. In quantum theory, valence electrons are restricted to particular areas of space, defined by atomic orbitals. Molecular orbital theory joins these two concepts to explain how atoms form covalent bonds using electrons in orbitals.

� In molecular orbital theory, electrons become shared in covalent bonds when the atomic orbitals containing the electrons approach close enough to overlap in space.

� This overlap forms a new molecular orbital that allows the electrons to exist simultaneously in the vicinity of both atoms in the bond (in other words, to be shared):

� In molecular orbital theory, orbitals which can’t overlap will not form covalent bonds.

B. Forming Molecular Orbitals

� In molecular orbital theory, molecular orbitals are formed by mathematically combining the functions that describe the individual atomic orbitals in the bond into new wavefunctions.

� The new wavefunctions describe new orbitals that have the following properties:

1. They extend over two (or more) atoms simultaneously.

2. They have a specific energy (just like atomic orbitals).

3. They can hold a maximum of two electrons (just like atomic orbitals).

� It takes at least two atomic orbitals (a.o.’s) to form a molecular orbital (m.o.). The two original atomic orbitals can hold a total of for electrons (two each). Therefore, each pair of atomic orbitals must form exactly two m.o.’s, so that there is still room for a total of four electrons. In general, when a given number of atomic orbitals are combined, they must form the same number of new molecular orbitals.



� For simple pairs of atomic orbitals, the new molecular orbitals are made by adding and subtracting the wavefunctions of the atomic orbitals.

� Adding the two a.o.’s is sometimes referred to as overlapping the orbitals in-phase.The lobes of the orbitals that are overlapped will have the same mathematical sign; therefore, in the overlap region the numbers produced by each function will add up.

� This means that in the new m.o. function, the numbers will be larger in the overlap region. Remember that the numbers represent the probability of finding the electron at that point in space, so the electrons will be denser in the overlap region.

� The m.o.’s formed by in-phase overlap are often called bonding orbitals, and they usually have a lower energy than the original a.o.’s that formed them.

� Subtracting the two a.o.’s (to create another m.o.) is sometimes referred to as out-of-phase overlap. The overlapping lobes will have opposite mathematical signs; therefore, the numbers in the overlap region will be smaller (sometimes they completely cancel out to zero).

� In this new m.o., there will be only a small probability of finding the electrons in the overlap region between the two atoms (low electron density in the overlap region).

� The m.o.’s formed by out-of-phase overlap are often called antibonding orbitals, and they usually have a higher energy than the original a.o.’s that formed them.



� A special type of diagram is used to show the formation of molecular orbitals and their energies:

� The initial atomic orbitals are placed on the outside of the diagram, and the new molecular orbitals are drawn in the middle.

� In the above diagram, the antibonding orbital goes up in energy more than the bonding orbital goes down.

C. Filling Molecular Orbitals with Electrons

� In general chemistry, we learned how to draw electron configuration diagrams that showed atomic orbitals with their electrons. For example, we could draw the electron configuration of oxygen:

� The orbitals are filled from the bottom, two electrons at a time. Degenerate orbitals (at the same energy) each receive one electron before any gets two.



� We can fill the newly formed m.o.’s using the electrons that were in the original atomic orbitals. For example, in H2, each hydrogen 1s orbital starts with one electron. When the m.o.’s are formed, the electrons move down into the lower energy bonding orbital:

� This explains why bonding is favored. The two electrons have a lower total energy when they are in the bonding m.o. than they would if they were separated back into their original atomic orbitals. Therefore, it is energetically favorable for the atoms to stay close, with overlapped orbitals. They are bonded; they will move together without separating to preserve the overlapped m.o.

XII.Valence Bond Theory (Orbital Hybridization)

� For more complicated molecules than H2, it quickly became apparent that bonding m.o.’s could not be formed from simple atomic orbitals. For example, if bonding m.o.’s in CH4 were formed from the carbon 2p orbitals, the bond would only be 90o

apart (because the 2p orbitals are 90o apart). However, we know that CH4 is tetrahedral, and the actual bonds are 109.5o apart.



� To solve this problem, we can mathematically combine the functions for the original atomic orbitals to form new atomic orbitals (with new functions). These new atomic orbitals are called hybrid orbitals, because they are by mixing the original atomic orbitals. The process of forming these new hybrid orbitals is called hybridization.

� The new hybrid atomic orbitals can then be used to form covalent bonds. The theory of using hybrid orbitals to form covalent bonds is called Valence Bond Theory.

A. Forming Hybrid Orbitals (Hybridization)

� Hybrid orbitals are formed by adding and subtracting atomic orbitals, similar to the away that m.o.’s are formed. The only difference is that in hybrid orbitals, all the atomic orbitals come from the same atom.

� When a given number of atomic orbitals are hybridized, they must form the same number of hybrid orbitals.

� An example of hybrid orbitals is shown below, using the 2s orbital and a 2p orbit:

� The hybrid orbitals are named sp orbitals because the are formed from one s orbital and one p orbital.

� Notice that when the 2s and 2px orbitals are used in hybridization, the other two 2p orbitals are left unhybridized.



� The key feature of hybrid orbitals is their geometry. Hybrid orbitals have the same geometry as the same number of electron pairs in VSEPR. This is shown below for sp2

and sp3 hybridized carbon:

B. Valence Bond Orbital Pictures

� In Valence Bond Theory, covalent bonds are still formed by overlapping atomic orbitals. However, in Valence Bond Theory, antibonding orbitals are usually ignored.

� This allows us to draw special pictures that show the orbital overlapping to form a molecule. An example is shown below for ethane (C2H6):



C. Multiple Bonds ( Bonds)

� Valence Bond Theory also explains how multiple bonds are formed:

� In a multiple bond, such as a double bond of an alkene, the first bond is formed between sp2 hybrid orbitals. This bond lies directly along the line connecting the two atoms; bonds along the line between two atoms are called sigma bonds ( bonds).

� The second bond is formed using the unhybridized p orbitals on the two atoms. These orbitals are perpendicular to the line between the two atoms, and they overlap side-to-side, rather than head on.

� This forms an orbital with two separate lobes—one above the bond and one below the bond. Although there are two lobes, this is still one orbital and it is still considered one bond. Bonds that have two lobes above and below the line between the atoms are called bonds.

� Triple bonds are formed from one bond and two bonds. The two bonds of a triple bond are perpendicular to each other:



� In the molecular orbital theory, p orbitals overlap in-phase and out-of-phase to form a pi orbital bonding orbital ( ) and a pi antibonding orbital ( *):

E. Hybridization and Bond Length

� P orbitals project farther out from the nucleus than s orbitals.

� As the contribution of p orbitals in a hybrid orbital increases, the hybrid orbital gets longer and projects farther from the nucleus. The contribution of a p orbital to a hybrid orbital is indicated as the percentage of p character in the hybrid orbital.

� As the percentage of s character increases in a hybrid orbital, the hybrid orbital gets shorter and closer to the nucleus.

� The percentages of s and p character in the three types of hybrid orbitals are summarized in the following table:

� The length of a bond depends on the length of the orbitals used to make the bond.

hybrid orbital percent s character percent p character

sp 50% 50%

sp2 33% 66%

sp3 25% 75%



� As the s character in the orbitals of the bond increases, the bond becomes shorter:

XIII.VSEPR: Molecular Geometry and Shape

� VSEPR: Valence Shell Electron Pair Repulsion. Groups of valence electrons arrange themselves around an atom as far apart as possible.

� Geometry: the arrangement of all groups of electrons (bonds and lone pairs) around a central atom.

� Shape: the arrangement of groups of atoms around a central atom. Lone pairs are ignored when naming the shape.

� Electron Group: a single bond, a double bond, a triple bond, or a lone pair each counts as one electron group. Double and triple bonds act as one electron group because all the electrons in the bond must stay in approximately the same area of space, so they cannot repel themselves apart and they act together to repel other electron groups.

� Bonding Groups: electron groups that are single, double, or triple bonds are bonding groups. Lone pairs are not bonding groups.

� A table of the common molecular geometries and shapes is on the next page.

Compound BondBond

Length, Å Orbitals in Bond

C–C 1.53 C(sp3)–C(sp3)

C–H 1.10 C(sp3)–H(1s)

C–C 1.50 C(sp3)–C(sp2)

C=C 1.34 C(sp2)–C(sp2)

C–H 1.10 C(sp3)–H(1s)

C–H 1.08 C(sp2)–H(1s)

C–C 1.46 C(sp3)–C(sp)

C C 1.21 C(sp)–C(sp)

C–H 1.10 C(sp3)–H(1s)

C–H 1.06 C(sp)–H(1s)

CH3 CH2 CH3

CH3 CH CH2

CH3 C CH



B. Drawing Three-Dimensional Structures

� In planar and linear structures, all of the atoms lie in a plane. It is relatively easy to accurately represent the three-dimensional (3-d) structure of the molecule by just drawing all the atoms as they are arranged in the plane:

� However, tetrahedral shapes are harder to represent because not all the atoms are in the same plane. Tetrahedral atom centers will have only three of the five atoms groups in the same plane; the other two atom groups will be arranged with one group sticking out of the plane, and the other group sticking back into the plane:

� We use wedge bonds to represent bonds sticking out of the plane toward us. We use dashed bonds to represent bonds sticking into the plane away from us:



C. Orbital Hybridization and Geometry

� The hybridization of an atom can be predicted from its VSEPR geometry:

� Resonance structures can add a complication. When the shape/geometry of an atom is different in different resonance structures, the actual atom geometry will be the one with the lowest number of electron groups. For example, if the atom is tetrahedral in one resonance structure, and trigonal planar in another, we generally observe that the actual shape of the atom is trigonal planar (tetrahedral = 4 electron groups and trigonal planar = 3 eletron groups).

XIV.Polar Covalent Bonds

� Electrons are not always shared equally in a covalent bond.

� If the two atoms in the bond are not identical, then the more electronegative atom will pull a greater share of the electrons toward itself.

� A good example of this is the H–O bond. Oxygen has a higher electronegativity than hydrogen, and consequently it pulls a larger share of the electrons in the bond toward itself. We can represent this by picturing the electrons as an elliptical cloud surround the two atoms. The oxygen end of the cloud would be larger and denser than the end near the hydrogen.

� When we calculated formal charges, we assumed that the electrons in a bond were equally shared, and that each atom in the bond “owned” exactly one of the electrons.

VSEPR Shape Electron Pair Geometry HybridizationOrbitals on

Central Atom

linear linear sp sp, sp, p, p

trigonal planar trigonal planar sp2 sp2, sp2, sp2, p

tetrahedral tetrahedral sp3 sp3, sp3, sp3, sp3

trigonal pyramidal tetrahedral sp3 sp3, sp3, sp3, sp3

bent tetrahedral sp3 sp3, sp3, sp3, sp3



� If the electrons in the bond are not equally shared, then one atom actually “owns” slightly more than one electron, and the other “owns” slightly less than one electron. If both atoms were initially determined to have zero formal charge, then unequal sharing of electrons would cause the more electronegative atom to have a slight excess of negative charge, and the less electronegative atom would have a slight positive charge.

� We represent these slight charges using the symbols + (which stands for partial positive charge) and – (which stands for partial negative charge). An H–O bond could therefore be represented as shown below, with a partial positive charge on hydrogen and a partial negative charge on oxygen. Effectively, the charges in the bond are unbalanced, and it has a positive and negative end (oppositely charged ends).

� Covalent bonds with oppositely charged ends are called polar covalent bonds (or polar bonds).

� Polarity is the general property of have oppositely charged ends.

� To represent the polarity of a bond, we use an arrow that points to the negative end of the bond and has a cross on the positive end of the bond.

� The size of the polarity in a bond is called the dipole moment, and it depends on both the difference in electronegativity and the length of the bond. The bond polarity arrow is considered a symbol for the dipole moment of a bond.



XV.Polar Molecules (Dipoles)

� When a molecule has oppositely charged ends, it is called a polar molecule or a dipole.

� Molecule polarity is more complex than bond polarity, because it involves both the polarity of bonds and the three-dimensional shape of the molecule.

� To determine if a molecule is polar, we first need to know the shape of the molecule.We then must add all the bond dipoles using vector addition.

� To vectors are arrows that indicate the size and direction of a number. Bond dipole moment arrows are vectors because the indicate the size and direction of the dipole moment.

� To add vectors, we line them up with the arrow of one vector pointing to the tail of the next vector. When all the vectors are lined up this way, the sum is an arrow from the start of the first vector to the arrowhead of the last vector. For example, look at the vector sum of the bond dipoles for water (H2O):

The vector sum of the bond dipoles is another vector, with a length greater than zero. We say that water has a net dipole moment. The dipole moment points from the positive end of the molecule to the negative end.

� A molecule can have polar bond and not be a polar molecule. This occurs when the bond dipoles balance each other out in three-dimensional space.

� A molecule will not be polar when the vector sum of its bond dipoles is zero (the last vector ends up at the starting point of the first vector. In this case we say the molecule has no net dipole moment and is non-polar.



� One example of a non-polar molecule is carbon dioxide, CO2. Although the two C=O bonds are polar bonds, the point in exactly opposite directions. Therefore, the bond dipoles balance each other out in space.

� Carbon dioxide is a non-polar molecule. If we put the + and – charges on the molecule, we can see that it doesn’t have oppositely charged ends. Instead, the ends of the molecule are both negatively charged (and the middle is positive). This is another way of visualizing the polarity of a molecule.

chem 220 notes page 1 structure and bonding of organic ... · patterns for the formal charges of...

Documents