AP Chemistry 2014-2015

CH 8 BONDING

Bonds are attractive forces that hold groups of atoms together and make them function as a unit.Being bound requires less energy than existing in the elemental form. Energy is released when a bond is formed; energy is required to break a bond. The energy required to break a bond is called the bond energy (endothermic).

BONDS

Ionic bond: electrostatic attraction between oppositely-charged ions; results when a metal reacts with a nonmetal

Covalent bond: type of bond in which bonding electrons are shared Unequal sharing results in polar covalent bonds; equal sharing results in nonpolar (or “pure” covalent bonds. We can categorize a bond qualitatively by looking at the periodic table, or quantitatively using electronegativity (more on this later). Most bonds are somewhere between purely ionic and purely covalent.

The following electrostatic potential maps correspond to H2, LiH, and HF. Determine which map belongs to each, and explain why.

EXERCISE 1 BOND TYPES

The map on the left belongs to LiH. LiH is ionic, so the sharing of electron density will be very unequal.The map in the middle belongs to H2, because it is nonpolar, so the electron density is shared equally.The map on the right belongs to HF. Its electron density is also shared unequally, but because it is polar covalent, the sharing is more equal than for LiH.

Coulomb’s Law is used to calculate the energy of an ionic bond (see equation on the right; k = 2.31 x 10 -19 Jnm (joules x nanometers), Q = charge on each ion, r = distance between ion centers in nanometers)The energy of an ionic bond will be negative; it indicates an attractive force so that the ion pair has lower energy than the separated ions.

IONIC BONDING

The final result of ionic bonding is a solid, regular array of cations and anions called a crystal lattice.

Lattice energy or lattice enthalpy: energy required to decompose an ion pair (from a lattice) into ions; a measure of th e strength of the ionic bond (related to Coulomb’s law)

The energy of attraction depends directly on the magnitude of the charges and inversely on the distance between them (related to the size of the ion).

EXERCISE 2 COMPARING LATTICE ENERGY

Which compound in each pair will have the higher lattice energy?

NaF or RbF MgO or LiCl

Coulomb’s Law can also be used to calculate the repulsive energy when two like-charged ions are brought together. The energy of this type of interaction would be positive.In the top scenario, the energy of the bond is negative (an attraction) because of the opposite charges.

In the bottom scenario, the energy of the “bond” is positive (a repulsion) because both charges are negative.

Continuing from the idea of positive repulsions and negative attractions… bond length is the distance where the sum of the attractive energy forces (negative sign) and repulsive energy forces (positive sign) is at a minimum. The attractive forces are due to the attraction of each nucleus for the electrons in the other atom in the bond. The repulsive forces are due to electron-electron repulsion and proton-proton repulsion.

ELECTRONEGATIVITY

Electronegativity is the ability of an atom in a molecule to attract shared electrons to itself.Fluorine is the most electronegative (4.0) due to highest effective nuclear charge (Zeff) and smallest radius.

Francium is the least (0.7) due to lowest Zeff and largest radius.

This atomic trend is only used when atoms form molecules. Ionic: ΔEN >1.67 Covalent: ΔEN < 1.67 Nonpolar covalent: ΔEN < 0.4

Electronegativity determines polarity since it measures a nucleus’s attraction or “pull” on the bonded electron pair. When two nuclei are the same, the sharing is equal and the bond is described as nonpolar. When two nuclei are different, the electrons are not shared equally, setting up slight +/- poles, and the bond is described as polar. When the electrons are shared very unequally, the bond is described as ionic.

BOND POLARITY AND ELECTRONEGATIVITY

Order the following bonds according to polarity: H-H, O-H, Cl-H, S-H, and F-H

EXERCISE 3 RELATIVE BOND POLARITIES

H-H S-H Cl-H O-H F-H

δ- δ- δ- δ-δ+

δ+

δ+

δ+

Most compounds are covalently bonded, especially carbon compounds.

Localized electron (LE) bonding model: assumes that a molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. Electron pairs are assumed to be localized on a particular atom [lone pairs] or in the space between two atoms [bonding pairs]. Lewis structures describe the valence

electron arrangement Geometry of the molecule is predicted

with VSEPR (valence shell electron pair repulsion)

COVALENT BONDING

Atoms of main-group elements tend to combine in such a way that each atom has eight electrons in its valence shell. This configuration can be achieved by sharing or transferring electrons.

OCTET RULE

Single bond: one pair of electrons shared (sigma σ bond)

Double bond: two pairs of electrons shared (one sigma bond, one pi bond)

Triple bond: three pairs of electrons shared (one sigma bond, two pi bonds)

Pi bonds are weaker than sigma bonds but never exist alone. Triple bonds are stronger than double bonds are stronger than single bonds.

SINGLE AND MULTIPLE BONDS

Compare the types of bonds found in ethane (C2H6), ethene (C2H4), and ethyne (C2H2). Also determine which molecule would have the strongest carbon-carbon bond.

EXERCISE 4 TYPES OF BONDS

Multiple bonds increase the electron density between two nuclei and therefore decrease the nuclear repulsions while enhancing the attraction between nucleus/electrons; either way, the nuclei move closer together and the bond length is shorter for a double bond than a single bond.

MULTIPLE BONDS ARE MOST OFTEN FORMED BY C, N, O, AND P.

Fewer than eight: H, Be, BExpanded valence: can only happen if the

central atom has d-orbitals and can thus be surround by more than four valence pairs in certain compounds.

Odd-electron compounds: ex. NO, NO2, ClO2

EXCEPTIONS TO THE OCTET RULE

For atomsWrite the symbolDetermine the number of valence electrons in the element

Place the valence electrons (dots) around the symbol—in order

Exceptions—carbon and silicon form hybrid orbitals and do not follow the usual order

DRAWING LEWIS STRUCTURES

Give the Lewis structure for each of the following. Sulfur Chlorine Silicon

Xenon  

Hydrogen Nitrogen Boron Fluorine

EXERCISE 5 WRITING LEWIS STRUCTURES (SINGLE ATOMS)

H is always a terminal atomAtom with lowest EN goes in center (or, the atom

capable of forming the largest number of bonds)Find the total number of valence electrons by adding

together the valence electrons of every atom in the compound For ions, add for negative charges and subtract for positive

chargesPlace one pair of electrons, a sigma bond, between

each pair of bonded atoms.Complete the octets of all atoms with lone pairs.

Leftover pairs are assigned to the central atom if it can accommodate them. Double/triple bonds may need to be used (pi bonds).

Watch out for charges—you might have to give the Lewis structure for an ion

FOR COMPOUNDS

Give the Lewis structure for each of the following.a) HFb) N2

c) NH3

d) CH4

e) CF4

f) NO+

EXERCISE 6 WRITING LEWIS STRUCTURES

Give the Lewis structure for each of the following. a) PCl5

b) ClF3

c) XeO3

d) RnCl2

e) BeCl2

f) ICl4-

EXERCISE 7 LEWIS STRUCTURES FOR MOLECULES THAT VIOLATE

THE OCTET RULE

Ex. ozone has equal bond lengths and equal bond strengths, implying that there are an equal number of bond pairs on each side of the central oxygen atom. The Lewis structure does not agree with this; instead, we have to use a composite to describe the reality. This composite depicts the blending of resonance structures for ozone. Instead of truly having a single bond and a double bond, both of its O-O bonds could be thought of as “a bond and a half”.

RESONANCE STRUCTURES

Resonance structures differ only in the assignment of electron pair positions, never atom positions. They differ in the number of bond pairs between a given pair of atoms.

Note that the resonance structures and composite are drawn with brackets (required for full credit on AP exam).

EXERCISE 8 RESONANCE STRUCTURES

Draw every resonance structure for the carbonate ion. Also draw the composite structure.

Bond order: simply the number of bonding electron pairs shared by two atoms in a molecule. 1 = one shared pair; sigma bond between two atoms 2 = two shared pairs; sigma bond and pi bond 3 = three shared pairs; sigma bond and two pi bonds Fractional for resonance structures (3/2 for ozone, 4/3

for carbonate)Bond order = number of shared pairs

linking X and Y number of X-Y links

BOND PROPERTIES

Bond length: the distance between the nuclei of two bonded atoms Higher bond order = shorter length

Bond energy: the greater the number of electron pairs between a pair of atoms, the shorter the bond. This implies that atoms are held together more tightly when there are multiple bonds, so there is a relationship between bond order and the energy required to break a bond.

BOND PROPERTIES

Formal charge is the difference between the number of valence electrons on a free element, and the number of electrons assigned to the atom once it is in a molecule.

Formal charge = group number – [# of lone electrons – 2(# of bonding electrons)]

FORMAL CHARGE

The ideal of formal charge allows us to determine the most favored structure out of a set of nonequivalent Lewis structures.

Oxidation states of more than +/- are questionable, while formal charges are more realistic.

The sum of the formal charges on an ion must equal the ion’s overall charge.

Use formal charges along with the following to determine resonance structureAtoms in molecules (or ions) should have formal charges that are as small (close to zero) as possible

A molecule (or ion)is most stable when any negative formal charge resides on the most electronegative atom.

Ex. There are three possible structures for the sulfate ion shown below (note that these are not resonance structures). The third is the most valid of the three; it results in the fewest (and smallest) formal charges.

Give possible Lewis structures for XeO3, an explosive compound of Xenon. Which Lewis structure or structures are most appropriate according to the formal charges?

EXERCISE 9 FORMAL CHARGES

Molecular shape changes with the number of sigma bonds plus lone pairs about the central atom

Molecular geometry is the arrangement in space of the atoms bonded to a central atom Lone pairs take up more space

around an atom than bonds Each lone pair or bond pair repels all

other electron pairs; they try to avoid each other making as wide an angle as possible. Ex. Water: the two lone pairs on oxygen

“warp” the normal 109.5 angle through repulsion, resulting in a bond angle of 104.5.

VALENCE SHELL ELECTRON PAIR REPULSION THEORY (VSEPR)

Sketch the Lewis dot structureDescribe the structural pair or electronic geometry

(the shape of the molecule considering both its bonds and lone pairs)

Focus on the bond locations (ignore lone pairs) and assign a molecular geometry based on their locations

Molecular geometry and electronic geometry are only the same in the absence of lone pairs on the central atom.

VSEPR works well for elements of the s and p-blocks; does not apply to transition element compounds (exceptions)

TO DETERMINE MOLECULAR GEOMETRY

Molecular shapes for central atoms with normal valenceNo more than 4 structural pairs if the atom obeys the Octet rule

The combination of s and p orbitals provides four bonding sites

Molecular shapes for central atoms with expanded valenceOnly elements with a principal energy level of 3 or higher can expand their valence and violate the octet rule on the high side. This is because d orbitals are needed for expansion to a 5th or 6th bonding location.

Lone pairs “want” to be as far apart as possible. For example, look at the two possible structures for XeF4 on the right. The equatorial configuration is favored (lower energy) because it allows the lone pairs to be as far from each other as possible.

In general, when you need to determine the most stable shape for a molecule, you should think about repulsions. Everything wants to be as far from everything else as possible (“everything” being electrons, of course).

Another example: SF4 has one lone pair and four S-F bonds. Configuration (a) is favored because it minimizes total repulsions (two 90° and two 129° lone-bonding repulsions vs. three 90° and one 180° lone-bonding repulsions)—see saw shape.

Which of the three following arrangements do you predict to be the most stable for I3

-?

EXERCISE 10 PREDICTION OF MOLECULAR STRUCTURE I

Draw and describe the molecular structure of the water molecule. 

EXERCISE 11 PREDICTION OF MOLECULAR STRUCTURE II

When phosphorous reacts with excess chlorine gas, the compound phosphorous pentachloride is formed. In the gaseous and liquid states, this substance consists of PCl5 molecules, but in the solid state it consists of a 1:1 mixture of PCl4

+ and PCl6- ions. Predict the

geometric structures of PCl5, PCl4+, and PCl6

-.

EXERCISE 12 PREDICTION OF MOLECULAR STRUCTURE III

Bonds can be polar while the entire molecule isn’t, and vice versa.

Dipole moment: separation of the charge in a molecule; product of the size of the charge and the distance of separation. Molecules align themselves with an electric field Molecules align with each other in the absence of an

electric field The direction of the “arrow” indicating the dipole moment

always points to the negative pole with the cross hatch on the arrow at the positive pole.

MOLECULAR POLARITY

In water, two lone pairs establish a strong negative pole. Similarly, a negative pole is established by the one lone pair in ammonia (NH3). This negative pole is flipped if we substitute the hydrogen atoms for fluorine to make nitrogen trifluoride.

If the Octet rule is obeyed and all the surrounding bonds are the same, then the molecule is nonpolar since all the dipole moments cancel each other out.Ex. carbon dioxide is nonpolar since the dipole moments cancel

Ex. methane, CH4, and carbon tetrachloride, CCl4, are both nonpolar since the dipole moments in each molecule cancel. However, the molecules “between” these two (shown below) all have net dipoles and are therefore polar.

For each of the following molecules, show the direction of the bond polarities and indicate which ones have a dipole moment: HCl, Cl2, SO3, and H2S.

EXERCISE 13 BOND POLARITY AND DIPOLE MOMENT

Predict the molecular structure of the sulfur dioxide molecule. Is this molecule expected to have a dipole moment?

EXERCISE 14 PREDICTION OF MOLECULAR STRUCTURE IV

1) VSEPR predicts an SbF5 molecule will be which of the following shapes?

A.tetrahedralB.trigonal bipyramidalC.square pyramidD.trigonal planarE.square planar

2) The shortest bond would be present in which of the following substances?

A. I2

B. COC. CCl4

D. O22–

E. SCl2