Without Drawing A Lewis Structure, Do You Think That Co Contains A Single, Double, Or Triple Bond?

When elements combine, there are two types of bonds that may form between them:

• Ionic bonds result from a transfer of electrons from one species (usually a metal) to another (usually a nonmetal or polyatomic ion).

• Covalent bonds result from a sharing of electrons by two or more atoms (usually nonmetals).

Lewis theory (Gilbert Newton Lewis, 1875-1946) focuses on the valence electrons, since the outermost electrons are the ones that are highest in energy and farthest from the nucleus, and are therefore the ones that are most exposed to other atoms when bonds form.

Lewis dot diagrams for elements are a handy way of picturing valence electrons, and especially, what electrons are available to be shared in covalent bonds. The valence electrons are written as dots surrounding the symbol for the element: one dot is place on each side first, and when all four positions are filled, the remaining dots are paired with one of the first set of dots, with a maximum of two dots placed on each side. Lewis-dot diagrams of the atoms in row 2 of the periodic table are shown below:

Unpaired electrons represent places where electrons can be gained in ionic compounds, or electrons that can be shared to form molecular compounds. (The valence electrons of helium are better represented by two paired dots, since in all of the noble gases, the valence electrons are in filled shells, and are unavailable for bonding.)

Covalent bonds generally form when a nonmetal combines with another nonmetal. Both elements in the bond are attracted to the unpaired valence electrons so strongly that neither can take the electron away from the other (unlike the case with ionic bonds), so the unpaired valence electrons are shared by the two atoms, forming a covalent bond:

The shared electrons act like they belong to both atoms in the bond, and they bind the two atoms together into a molecule. The shared electrons are usually represented as a line (�) between the bonded atoms. (In Lewis structures, a line represents two electrons.)

Atoms tend to form covalent bonds in such a way as to satisfy the octet rule, with every atom surrounded by eight electrons. (Hydrogen is an exception, since it is in row 1 of the periodic table, and only has the 1s orbital available in the ground state, which can only hold two electrons.)

The shared pairs of electrons are bonding pairs (represented by lines in the drawings above). The unshared pairs of electrons are lone pairs or nonbonding pairs.

All of the bonds shown so far have been single bonds, in which one pair of electrons is being shared. It is also possible to have double bonds, in which two pairs of electrons are shared, and triple bonds, in which three pairs of electrons are shared:

Multiple bonds are shorter and stronger than their corresponding single bond counterparts.

Rules for Writing Lewis Structures

1. Count the total number of valence electrons in the molecule or polyatomic ion. (For example, H₂O has 2x1 + 6 = 8 valence electrons, CCl₄ has 4 + 4x7 = 32 valence electrons.) For anions, add one valence electron for each unit of negative charge; for cations, subtract one electron for each unit of positive charge. (For example, NO₃ ^- has 5 + 3x6 + 1 = 24 valence electrons; NH₄ ⁺ has 5 + 4+1 � 1 = 8 valence electrons.)
2. Place the atoms relative to each other. For molecules of the formula AX_n, place the atom with the lower group number in the center. If A and X are in the same group, place the atom with the higher period number in the center. (This places the least electronegative atom in the center.) H is NEVER UNDER ANY CIRCUMSTANCES a central atom.
3. Draw a single bond from each terminal atom to the central atom. Each bond uses two valence electrons.
4. Distribute the remaining valence electrons in pairs so that each atom obtains eight electrons (or 2 for H). Place the lone pairs on the terminal atoms first , and place any remaining valence electrons on the central atom. The number of electrons in the final structure must equal the number of valence electrons from Step 1.
5. If an atom still does not have an octet, move a lone pair from a terminal atom in between the terminal atom and the central atom to make a double or triple bond. Use the formal charge as a guideline for placing multiple bonds:

Formal charge = valence � (� bonding e-) � (lone pair e-)

• The formal charge is the charge an atom would have if the bonding electrons were shared equally.
• The sum of the formal charges must equal the charge on the species.
• Smaller formal charges are better (more stable) than larger ones.
• The number of atoms having formal charges should be minimized.
• Like charges on adjacent atoms are not desirable.
• A more negative formal charge should reside on a more electronegative atom.

Examples
1.	CH4 (methane)
	8 valence electrons (4 + 4x1)
	Place the C in the center, and connect the four H�s to it: This uses up all of the valence electrons. The octet rule is satisfied everywhere, and all of the atoms have formal charges of zero.
2.	NH3 (ammonia)
	8 valence electrons (5 + 3x1)
	Place the N in the center, and connect the three H�s to it: This uses up six of the eight valence electrons. The last two electrons cannot go on the H�s (that would violate the octet rule for H), so they must go on the N: All of the valence electrons have now been used up, the octet rule is satisfied everywhere, and all of the atoms have formal charges of zero.
3.	H2O (water)
	8 valence electrons (2x1 + 6)
	Place the O in the center, and connect the two H�s to it: This uses up four of the valence electrons. The remaining four valence electrons cannot go on the H�s, so they must go on the O, in two pairs: All of the valence electrons have now been used up, the octet rule is satisfied everywhere, and all of the atoms have formal charges of zero.
4.	H3O+ (hydronium ion)
	8 valence electrons (3x1 + 6 � 1)
	Place the O in the center, and connect the three H�s to it: This uses up six of the valence electrons. The remaining two valence electrons must go on the oxygen: All of the valence electrons have been used up, and the octet rule is satisfied everywhere. The formal charge on the oxygen atom is 1+ (8 � ��6 � 2):
5.	HCN (hydrogen cyanide)
	10 valence electrons (1 + 4 + 5)
	Place the C in the center, and connect the H and N to it: This uses up four of the valence electrons. The remaining six valence electrons start out on the N: In the structure as shown, the octet rule is not satisfied on the C, and there is a 2+ formal charge on the C (4 � ��4 � 0) and a 2- formal charge on the N (5 � ��2 � 6): The octet rule can be satisfied if we move two pairs of electrons from the N in between the C and the N, making a triple bond: The octet rule is now satisfied, and the formal charges are zero.
6.	CO2 (carbon dioxide)
	16 valence electrons (4 + 2x6) Place the C in the center, connect the two O�s to it, and place the remaining valence electrons on the O�s: This uses up the sixteen valence electrons The octet rule is not satisfied on the C, and there are lots of formal charges in the structure: The octet rule can be satisfied, and the formal charges diminished if we move a pair of electrons from each oxygen atom in between the carbon and oxygen atoms: The octet rule is satisfied everywhere, and all of the atoms have formal charges of zero.
7.	CCl4 (carbon tetrachloride)
	32 valence electrons (4 + 4x7)
	Place the C in the center, and connect the four Cl�s to it: This uses up eight valence electrons The remaining 24 valence electrons are placed in pairs on the Cl�s: Now, all of the valence electrons have been used up, the octet rule is satisfied everywhere, and all of the atoms have formal charges of zero.
8.	COCl2 (phosgene or carbonyl chloride)
	24 valence electrons (4 + 6 + 2x7)
	Place the C in the center, and connect the O and the two Cl�s to it. (The relative placement of the O and the Cl�s does not matter, since we are not yet drawing a three-dimensional structure.) Place the remaining valence electrons on the O and Cl atoms: The octet rule is not satisfied on the C; in order to get eight electrons around the C, we must move a pair of electrons either from the O or one of the Cl�s to make a double bond. Making a carbon-chlorine double bond would satisfy the octet rule, but there would still be formal charges, and there would be a positive formal charge on the strongly electronegative Cl atom (structure 2). Making a carbon-oxygen double bond would also satisfy the octet rule, but all of the formal charges would be zero, and that would be the better Lewis structure (structure 3):

Examples (continued from section B)
9.	O3 (ozone)
	18 valence electrons (3x6)
	Place one O in the center, and connect the other two O�s to it. Drawing a single bond from the terminal O�s to the one in the center uses four electrons; 12 of the remaining electrons go on the terminal O's, leaving one lone pair on the central O: We can satisfy the octet rule on the central O by making a double bond either between the left O and the central one (2), or the right O and the center one (3): The question is, which one is the �correct� Lewis structure?

In this example, we can draw two Lewis structures that are energetically equivalent to each other � that is, they have the same types of bonds, and the same types of formal charges on all of the structures. Both structures (2 and 3) must be used to represent the molecule�s structure. The actual molecule is an average of structures 2 and 3, which are called resonance structures. (Structure 1 is also a resonance structure of 2 and 3, but since it has more formal charges, and does not satisfy the octet rule, it is a higher-energy resonance structure, and does not contribute as much to our overall picture of the molecule.) Structures 2 and 3 in the example above are somewhat �fictional� structures, in that they imply that there are �real� double bonds and single bonds in the structure for ozone; in reality, however, ozone has two oxygen-oxygen bonds which are equal in length, and are halfway between the lengths of typical oxygen-oxygen single bonds and double bonds � effectively, there are two �one-and-a-half� bonds in ozone. The real molecule does not alternate back and forth between these two structures; it is a hybrid of these two forms. (This is analogous to describing a real person as having the characteristics of two or more fictional characters � the fictional characters don�t exist, but the real person does. Another analogy is to consider a mule: a mule is a cross or hybrid between a horse and a donkey, but it doesn�t alternate between being a horse and a donkey.)

The ozone molecule, then, is more correctly shown with both Lewis structures, with the two-headed resonance arrow () between them:

In these resonance structures, one of the electron pairs (and hence the negative charge) is �spread out� or delocalized over the whole molecule. In contrast, the lone pairs on the oxygen in water are localized � i.e., they�re stuck in one place. Resonance delocalization stabilizes a molecule by spreading out charges, and often occurs when lone pairs (or positive charges) are located next to double bonds. Resonance plays a large role in our understanding of structure and reactivity in organic chemistry. (A more accurate picture of bonding in molecules like this is found in Molecular Orbital theory, but this theory is more advanced, and mathematically more complex topic, and will not be dealt with here.)

As a general rule, when it�s possible to make a double bond in more than one location, and the resulting structures are energetically equivalent to each other, each separate structure must be shown, separated from each other by resonance arrows.

Examples
10.	CO3 2- (carbonate ion)
	24 valence electrons (4 + 3x6 + 2)
	Place the C in the center, with three lone pairs on each of the O�s: We can satisfy the octet rule and make the formal charges smaller by making a carbon-oxygen double bond. Since there are three energetically equivalent ways of making a C=O, we draw each of the three possible structures, with a resonance arrow between them: Once again, structure 1 is a resonance structure of 2, 3, and 4, but it is a higher energy structure, and does not contribute as much to our picture of the molecule. Since the double bond is spread out over three positions, the carbon-oxygen bonds in carbonate are �one-and-a-third� bonds.

Molecules with more than one central atoms are drawn similarly to the ones above. The octet rule and formal charges can be used as a guideline in many cases to decide in which order to connect atoms.

Examples
11.	C2H6 (ethane)
12.	C2H4 (ethylene)
13.	CH3CH2OH (ethyl alcohol)

A number of species appear to violate the octet rule by having fewer than eight electrons around the central atom, or by having more than eight electrons around the central atom. Once again, the formal charge is a good guideline to use to decide whether a �violation� of the octet rule is acceptable.

• Electron deficient species, such as beryllium (Be), boron (B), and aluminum (Al) can have fewer than eight electrons around the central atoms, but have zero formal charge on that atom. Molecules with electron deficient central atoms tend to be fairly reactive (many electron-deficient species act as Lewis acids).
• Free radicals contain an odd number of valence electrons. As a result, one atom in the Lewis structure will have an odd number of electrons, and will not have a complete octet in the valence shell. These species are extremely reactive. When drawing these compounds, optimize the placement of bonds and the odd electron to minimize formal charges; there are often several possible resonance structures than can be drawn.
• Expanded valence shells are often found in nonmetals from period 3 or higher, such as sulfur, phosphorus, and chlorine. These species can accommodate more than 8 electrons by shoving �extra� electrons into empty d orbitals. For example, sulfur's valence shell contains 3s, 3p, and 3d orbitals (since sulfur is in row 3 of the periodic table, the valence shell is n=3); however, since there are only 16 electrons on a neutral sulfur atom, the 3d orbitals are unoccupied.  When sulfur forms a compound with another element, the empty 3d orbitals can accommodate additional electrons.  Note that period 2 elements CANNOT have more than eight electrons, since the n=2 shell has no d orbitals to put �extra� electrons in.

Examples
14.	BF3 (boron trifluoride)
	24 valence electrons (3 + 3x7)
	The octet rule is not satisfied on the B, but the formal charges are all zero. (In fact, trying to make a boron-fluorine double bond would put a positive formal charge on fluorine; since fluorine is highly electronegative, this is extremely unfavorable.)
15.	NO (nitrogen monoxide, or nitric oxide)
	11 valence electrons (5 + 6)
	In this structure, the formal charges are all zero, but the octet rule is not satisfied on the N. Since there are an odd number of electrons, there is no way to satisfy the octet rule. Nitric oxide is a free radical, and is an extremely reactive compound. (In the body, nitric oxide is a vasodilator, and is involved in the mechanism of action of various neurotransmitters, as well as some heart and blood pressure medications such as nitroglycerin and amyl nitrite)
16.	PCl5 (phosphorus pentachloride)
	40 valence electrons (5 + 5x7)
	The octet rule is violated on the central P, but phosphorus is in the p-block of row 3 of the periodic table, and has empty d orbitals that can accommodate �extra� electrons. Notice that the formal charge on the phosphorus atom is zero.
17.	SF6 (sulfur hexafluoride)
	48 valence electrons (6 + 6x7)
	The octet rule is violated on the central S, but sulfur is in the p-block of row 3 of the periodic table, and has empty d orbitals that can accommodate �extra� electrons. Notice that the formal charge on the sulfur atom is zero.
18.	SF4 (sulfur tetrafluoride)
	48 valence electrons (6 + 6x7)
	The octet rule is violated on the central S, but sulfur is in the p-block of row 3 of the periodic table, and has empty d orbitals that can accommodate �extra� electrons. Notice that the formal charge on the sulfur atom is zero.
19.	XeF4 (xenon tetrafluoride)
	36 valence electrons (8 + 4x7)
	The octet rule is violated on the central Xe, but xenon is in the p-block of row 5 of the periodic table, and has empty d orbitals that can accommodate �extra� electrons. Notice that the formal charge on the xenon atom is zero.
20.	H2SO4 (sulfuric acid)
	32 valence electrons (2x1 + 6 + 4x6)
	Structures 1 and 2 are resonance structures of each other, but structure 2 is the lower energy structure, even though it violates the octet rule. Sulfur can accommodate more than eight electrons, and the formal charges in structure 2 are all zero.

Drawing a Lewis structure is the first steps towards predicting the three-dimensional shape of a molecule. A molecule�s shape strongly affects its physical properties and the way it interacts with other molecules, and plays an important role in the way that biological molecules (proteins, enzymes, DNA, etc.) interact with each other.

The approximate shape of a molecule can be predicted using the Valence-Shell Electron-Pair Repulsion (VSEPR) model, which depicts electrons in bonds and lone pairs as �electron groups� that repel one another and stay as far apart as possible:

1. Draw the Lewis structure for the molecule of interest and count the number of electron groups surrounding the central atom. Each of the following constitutes an electron group:
   • a single, double or triple bond (multiple bonds count as one electron group)
   • a lone pair
   • an unpaired electron
2. Predict the arrangement of electron groups around each atom by assuming that the groups are oriented in space as far away from one another as possible.
3. The shapes of larger molecules having more than one central are a composite of the shapes of the atoms within the molecule, each of which can be predicted using the VSEPR model.

Two Electron Groups

2 bonds, 0 lone pairs

linear

bond angles of 180�

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Three Electron Groups

3 bonds, 0 lone pairs		2 bonds, 1 lone pair
trigonal planar		bent
bond angles of 120�		bond angles of < 120�
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Lone pairs take up more room than covalent bonds; this causes the other atoms to be squashed together slightly, decreasing the bond angles by a few degrees.

Four Electron Groups

4 bonds, 0 lone pairs		3 bonds, 1 lone pair		2 bonds, 2 lone pairs
tetrahedral		trigonal pyramidal		bent
bond angles of 109.5�		bond angles of < 109.5�		bond angles of < 109.5�
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Five Electron Groups

5 bonds, 0 lone pairs		4 bonds, 1 lone pair		3 bonds, 2 lone pairs		2 bonds, 3 lone pairs
trigonal bipyramidal		seesaw		T-shaped		linear
bond angles of 120� (equatorial), 90� (axial)		bond angles of <120� (equatorial), <90� (axial)		bond angles of < 90�		bond angles of 180�
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The trigonal bipyramidal shape can be imagined as a group of three bonds in a trigonal planar arrangement separated by bond angles of 120� (the equatorial positions), with two more bonds at an angle of 90� to this plane (the axial positions):

Lone pairs go in the equatorial positions, since they take up more room than covalent bonds. In the equatorial position, lone pairs are ~120� from two other bonds, while in the axial positions they would be 90� away from three other bonds.

Six Electron Groups

6 bonds, 0 lone pairs		5 bonds, 1 lone pair		4 bonds, 2 lone pairs
octahedral		square pyramidal		square planar
bond angles of 90�		bond angles of < 90�		bond angles of 90�
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The Lewis structures of the previous examples can be used to predict the shapes around their central atoms:

	Formula	Lewis Structure	Bonding	Shape
1.	CH4		4 bonds 0 lone pairs	tetrahedral
2.	NH3		3 bonds 1 lone pair	trigonal pyramidal
3.	H2O		2 bonds 2 lone pairs	bent
4.	H3O+		3 bonds 1 lone pair	trigonal pyramidal
5.	HCN		2 bonds 0 lone pairs	linear
6.	CO2		2 bonds 0 lone pairs	linear
7.	CCl4		4 bonds 0 lone pairs	tetrahedral
8.	COCl2		3 bonds 0 lone pairs	trigonal planar
9.	O3		2 bonds 1 lone pair	bent*
10.	CO3 2-		3 bonds 0 lone pairs	trigonal planar*
11.	C2H6		4 bonds 0 lone pairs	tetrahedral
12.	C2H4		3 bonds 0 lone pairs	trigonal planar
13.	CH3CH2OH		C: 4 bonds     0 lone pairs O: 2 bonds      2 lone pairs	C: tetrahedral O: bent
14.	BF3		3 bonds 0 lone pairs	trigonal planar
15.	NO			linear
16.	PCl5		5 bonds 0 lone pairs	trigonal bipyramidal
17.	SF6		6 bonds 0 lone pairs	octahedral
18.	SF4		4 bonds 1 lone pair	seesaw
19.	XeF4		4 bonds 2 lone pairs	square planar
20.	H2SO4		S: 4 bonds     0 lone pairs O: 2 bonds      2 lone pairs	S: tetrahedral O: bent

With Lewis structures involving resonance, it is irrelevant which structure is used to determine the shape, since they are all energetically equivalent.

Electronegativity is a measure of the ability of an atom in a molecule to attract shared electrons in a covalent bond. Electronegativity is a periodic property, and increases from bottom to top within a group and from left to right across a period:

Table 1. Periodic Trends in Electronegativity

Table 2. Electronegativity Values (Pauling scale)

When two atoms of the same electronegativity share electrons, the electrons are shared equally, and the bond is a nonpolar covalent bond � there is a symmetrical distribution of electrons between the bonded atoms. (As an analogy, you can think of it as a game of tug-of-war between two equally strong teams, in which the rope doesn�t move.) For example, when two chlorine atoms are joined by a covalent bond, the electrons spend just as much time close to one chlorine atom as they do to the other; the resulting molecule is nonpolar (indicated by the symmetrical electron cloud shown below):

When two bonded atoms have a difference of greater than 2.0 electronegativity units (see Table 2), the bond is an ionic bond � one atoms takes the electrons away from the other atom, producing cations and anions.  For example Na has an electronegativity of 0.93, and Cl is 3.16, a difference of 2.23 units. The Cl atom takes an electron away from the Na, producing a fully ionic bond:

When two bonded atoms have a difference of between 0.4 and 2.0 electronegativity units (see Table 2), the electrons are shared unequally, and the bond is a polar covalent bond � there is an unsymmetrical distribution of electrons between the bonded atoms, because one atom in the bond is �pulling� on the shared electrons harder than the other, but not hard enough to take the electrons completely away. The more electronegative atom in the bond has a partial negative charge ( -), because the electrons are pulled slightly towards that atom, and the less electronegative atom has a partial positive charge ( +), because the electrons are partly (but not completely) pulled away from that atom. For example, in the HCl molecule, chlorine is more electronegative than hydrogen by 0.96 electronegativity units. The shared electrons are pulled slightly closer to the chlorine atom, making the chlorine end of the molecule very slightly negative (indicated in the figure below by the larger electron cloud around the Cl atom), while the hydrogen end of the molecule is very slightly positive (indicated by the smaller electron cloud around the H atom), and the resulting molecule is polar:

We say that the bond has a dipole � the electron cloud is polarized towards one end of the molecule.  The degree of polarity in a covalent bond depends on the electronegativity difference, DEN, between the two bonded atoms:

• DEN 0 - 0.4 = Nonpolar covalent bond

• DEN 0.4 - 2.0  = Polar covalent bond

• DEN > 2.0 = Ionic bond

In a diatomic molecule (X₂ or XY), there is only one bond, and the polarity of that bond determines the polarity of the molecule: if the bond is polar, the molecule is polar, and if the bond is nonpolar, the molecule is nonpolar.

In molecules with more than one bond, both shape and bond polarity determine whether or not the molecule is polar. A molecule must contain polar bonds in order for the molecule to be polar, but if the polar bonds are aligned exactly opposite to each other, or if they are sufficiently symmetric, the bond polarities cancel out, making the molecule nonpolar. (Polarity is a vector quantity, so both the magnitude and the direction must be taken into account.)

For example, consider the Lewis dot structure for carbon dioxide. This is a linear molecule, containing two polar carbon-oxygen double bonds. However, since the polar bonds are pointing exactly 180� away from each other, the bond polarities cancel out, and the molecule is nonpolar. (As an analogy, you can think of this is being like a game of tug of war between two teams that are pulling on a rope equally hard.)

The water molecule also contains polar bonds, but since it is a bent molecule, the bonds are at an angle to each other of about 105�. They do not cancel out because they are not pointing exactly towards each other, and there is an overall dipole going from the hydrogen end of the molecule towards the oxygen end of the molecule; water is therefore a polar molecule:

Molecules in which all of the atoms surrounding the central atom are the same tend to be nonpolar if there are no lone pairs on the central atom. If some of the atoms surrounding the central atom are different, however, the molecule may be polar. For example, carbon tetrachloride, CCl₄, is nonpolar, but chloroform, CHCl₃, and methyl chloride, CH₃Cl are polar:

The polarity of a molecule has a strong effect on its physical properties. Molecules which are more polar have stronger intermolecular forces between them, and have, in general, higher boiling points (as well as other different physical properties).

The table below shows whether the examples in the previous sections are polar or nonpolar. For species which have an overall charge, the term �charged� is used instead, since the terms �polar� and �nonpolar� do not really apply to charged species; charged species are, by definition, essentially polar. Lone pairs on some outer atoms have been omitted for clarity.

	Formula	Lewis Structure	3D Structure Shape  Polarity	Explanation
1.	CH4		tetrahedral nonpolar	The C�H bond is nonpolar, since C and H differ by only 0.35 electronegativity units.
2.	NH3		trigonal pyramidal polar	Since this molecule is not flat, the N�H bonds are not pointing directly at each other, and their polarities do not cancel out. In addition, there is a slight dipole in the direction of the lone pair.
3.	H2O		bent polar	Since this molecule is bent, the O�H bonds are not pointing directly at each other, and their polarities do not cancel out.
4.	H3O+		trigonal pyramidal charged	Since this species is charged, the terms �polar� and �nonpolar� are irrelevant.
5.	HCN		linear polar	Linear molecules are usually nonpolar, but in this case, not all of the atoms connected to the central atom are the same. The C�N bond is polar, and is not canceled out by the nonpolar C�H bond.
6.	CO2		linear nonpolar	The polar C=O bonds are oriented 180� away from each other. The polarity of these bonds cancels out, making the molecule nonpolar.
7.	CCl4		tetrahedral nonpolar	The polar C�Cl bonds are oriented 109.5� away from each other. The polarity of these bonds cancels out, making the molecule nonpolar.
8.	COCl2		trigonal planar polar	Trigonal planar molecules are usually nonpolar, but in this case, not all of the atoms connected to the central atom are the same. The bond polarities do not completely cancel out, and the molecule is polar. (If there were three O�s, or three Cl�s attached to the central C, it would be nonpolar.)
9.	O3		bent polar	Bent molecules are always polar. Although the oxygen-oxygen bonds are nonpolar, the lone pair on the central O contributes some polarity to the molecule.
10.	CO3 2-		trigonal planar charged	Since this species is charged, the terms �polar� and �nonpolar� are irrelevant.
11.	C2H6		tetrahedral nonpolar	Both carbon atoms are tetrahedral; since the C�H bonds and the C�C bond are nonpolar, the molecule is nonpolar.
12.	C2H4		trigonal planar nonpolar	Both carbon atoms are trigonal planar; since the C�H bonds and the C�C bond are nonpolar, the molecule is nonpolar.
13.	CH3CH2OH		C: tetrahedral O: bent polar	The C�C and C�H bonds do not contribute to the polarity of the molecule, but the C�O and O�H bonds are polar, the since the shape around the O atom is bent, the molecule must be polar.
14.	BF3		trigonal planar nonpolar	Since this molecule is planar, all three polar B�F bonds are in the same plane, oriented 120� away from each other, making the molecule nonpolar.
15.	NO		linear polar	Since there is only one bond in this molecular, and the bond is polar, the molecule must be polar.
16.	PCl5		trigonal bipyramidal nonpolar	The P�Cl bonds in the equatorial positions on this molecule are oriented 120� away from each other, and their bond polarities cancel out. The P�Cl bonds in the axial positions are 180� away from each other, and their bond polarities cancel out as well.
17.	SF6		octahedral nonpolar	The S�F bonds in this molecules are all 90� away from each other, and their bond polarities cancel out.
18.	SF4		seesaw polar	The S�F bonds in the axial positions are 90� apart, and their bond polarities cancel out. In the equatorial positions, since one position is taken up by a lone pair, they do not cancel out, and the molecule is polar.
19.	XeF4		square planar nonpolar	The Xe�F bonds are all oriented 90� away from each other, and their bond polarities cancel out. The lone pairs are 180� away from each other, and their slight polarities cancel out as well.
20.	H2SO4		S: tetrahedral O: bent polar	This molecule is polar because of the bent H�O�S bonds which are present in it.

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1. �Electron groups� include bonds, lone pairs, and odd (unpaired) electrons. A multiple bond (double bond or triple bond) counts as one electron group.
2. A multiple bond (double bond or triple bond) counts as one bond in the VSEPR model.
3. A = central atom, X = surrounding atoms, E = lone pairs
4. Molecules with this shape are nonpolar when all of the atoms connected to the central atom are the same. If the atoms connected to the central atom are different from each other, the molecular polarity needs to be considered on a case-by-case basis.
5. Since electrons in lone pairs take up more room than electrons in covalent bonds, when lone pairs are present the bond angles are �squashed� slightly compared to the basic structure without lone pairs.

Martin S. Silberberg, Chemistry:  The Molecular Nature of Matter and Change, 2nd ed.  Boston:  McGraw-Hill, 2000, p. 374-384.

Nivaldo J. Tro, Chemistry:  A Molecular Approach, 1st ed.  Upper Saddle River:  Pearson Prentice Hall, 2008, p. 362-421.

Without Drawing A Lewis Structure, Do You Think That Co Contains A Single, Double, Or Triple Bond?

Source: https://www.angelo.edu/faculty/kboudrea/general/shapes/00_lewis.htm

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