Bonding and Structure

AP REVIEW: Bonding and Structure

Bonds form due to electrostatic attractions; negatively charged electrons of one atom are attracted to the positively charged nucleus of another atom.

Atoms bond together to form molecules, and this is known as intramolecular bonding. The atoms are driven to bond in a way that they can achieve the most stable electron configuration possible. They are most stable with a half filled or filled principle quantum level. There are two types of intramolecular bonding:
1. Ionic Bonding – one atom completely transfers an electron to another, forming two oppositely charged ions that are attracted to each other. According to Coulombs Law, bond strength (and lattice energy) is directly proportional to the size of the charges, and indirectly proportional to the inter-ionic distance. Ionic bonds form when there is a large difference (about 2 or greater) in electronegativity between two atoms, and most often involve a metal and non-metal.
2. Covalent Bonding – two atoms share electrons, there is no complete exchange. Each atom counts the electrons as part of its valence shell to become more stable. The number of electrons in an atom’s valence shell is the same as the number of bonds it can form.
>Sigma (σ) bonds are the first covalent bond formed between two atoms, and all single bonds are sigma bonds
>Pi (π) bonds are additional bonds (as in the second in double bonds and second and third in triple bonds). Pi bonds are generally weaker than sigma bonds
A net increase in ▲H for a reaction means the bonds formed are weaker than the bonds broken, and a net decrease means the bonds formed are stronger than the bonds broken.
▲H° = Σ(bond energies of bonds broken) – Σ(bond energies of bonds formed)

Polarity occurs in molecules with covalent bonding when the atoms have differing electronegativity. The more electronegative element has a greater force on the electrons, creating an unequal distribution of the negative charge in the molecule, and end of the molecule with the more electronegative atom becomes negatively charged while the other is more positively charged. Other factors to polarity are the presence of lone pairs of electrons on the central atom, and molecular symmetry.

Lewis Dot Diagrams
Used to illustrate how elements combine to form molecules, as well as for predicting the shapes of the molecules. Apply the following steps:
1. Draw the skeletal structure of the molecule, generally placing the atom with lowest electron affinity in the center (note that H is always a terminal atom)
2. Count the valence electrons in the molecule, if it is an ion add or subtract electrons for the charge of the ion.
3. Draw single bonds between the atoms, each of which represents one lone pair.
4. Assign the remaining electrons to terminal atoms until the octet rule is satisfied for each atom, or all the electrons are used. Any remaining pairs are assigned to the central atom.
5. Use the octet rule to determine multiple bonds. If a central atom does not have a full octet, adjust the diagram by removing lone pairs from the terminal atoms and forming a double or triple bond.
The octet rule calls for atoms to be surrounded by 8 electrons (4 pairs of electrons), and the atoms of C, N, O, and F always have an octet. However there are many exceptions to the octet rule.
Incomplete Octet: Some atoms can have a complete outer shell with less than 8 electrons such as atoms from groups 1 and 2.
Expanded Octet: Central atoms in molecules with available d subshells (starting with the third period) can form more than 4 bonds (they can have more than 8 valence electrons).
Odd number of electrons: On rare occasions molecules will have an odd number of electrons, making it impossible to have a complete pair, most often involving N.
Resonance forms are the forms of a molecule that are identical except for where the multiple bond(s) in the atoms are placed. All resonance forms of a molecule are considered to exist simultaneously.
Sometimes there is more than one possible Lewis Structure possible for a molecule, so formal charge must be taken into account to determine the best one. Formal charge is calculated by subtracting the number of electrons assigned to the atom in the Lewis structure from number of valence electrons in the isolated atom. The structure with a formal charge closest to zero is most favorable.

Molecular Geometry
The valence shell electron-pair repulsion model (VSEPR) is used to predict the geometries of molecules. A basic concept behind the geometries is that electrons repel each other, so when the atoms form a molecule it will take on a shape that allows the electron pairs to be as far apart as possible. The number of electron pairs on the central atom determines the molecule’s shape.
Atomic orbitals mix to form new, hybrid orbitals through the process of hybridization. Hybrid orbitals have a different shape than the original atomic orbitals, however the total number of hybrid orbitals is the same as the number of original atomic orbitals.
The hybrid orbitals orient themselves so that they are the maximum distance apart from each other because electron pairs repel each other. The number of hybrid orbitals needed is determined by adding the number of lone pairs and number of bonds on the central atom (multiple bonds only count for one hybrid orbital). Each lone pair and each bond will require a hybrid orbital. When considering molecular geometry and bond angles, remember:
1. Lone (unbonded) electron pairs exert a more a powerful repulsive force than bonded pairs, they will occupy more space than bonding pairs.
2. Multiple bonds have slightly more repulsive strength than single bonds and will therefore occupy more space.
Refer to the table for molecular/basic geometries based on number of hybrid orbitals and lone pairs.

Review Questions

1. (a) Draw the Lewis Dot Diagram for NF3.
(b) Determine the hybridization of the orbitals in the NF3 molecule.
(c) Explain why the angle between the N-F bonds in NF3 is smaller than the angle between the B-F bonds in BF3.


2. Use the VSEPR model to predict the molecular geometry of each of the following molecules (no cheating and looking at the chart!!): (a) H20 (b) XeF2 (c) BrF5 (d) NH3


3. Calculate ▲H for the reaction between 2 moles of hydrogen gas and 1 mole of oxygen gas. Use the following values:
Bond
Bond Energy (kJ/mol)
H-H
436
O=O
499
O-H
463


Answers & Explanations
1. (a)


(b) sp3 hybridization
The number of hybrid orbitals is determined by adding the number of bonds + the number of lone electron pairs on the central atom. The central atom (N) has 3 single bonds and 1 lone electron pair, so its hybrid orbital number is 4 and the molecule requires 4 hybrid orbitals. Therefore the hybridization must be sp3.
(c) NF3 has three bonds and one lone electron pair on the central atom. The basic shape of the molecule is tetrahedral, so the bond angles are theoretically 109.5°. The additional repulsion of the lone electron pair makes the N-F bonds even closer together so the angle between them is less that 109.5°. BF3 on the other hand has 3 bonds and NO lone electron pairs, so the basic structure is identical to the molecular structure of trigonal planar which has bond angles of 120°.

2. Refer to tables 9.2 & 9.3 for complete illustrations
(a) bent
The electron domain structure of the water molecule is tetrahedral; however there are two bonding pairs and two lone electron pairs, making the geometric structure bent.
(b) linear
The electron domain structure of the molecule is trigonal bipyramidal; however there are 3 lone electron pairs and two bonding pairs, so the molecular geometry becomes linear.
(c) square pyramidal
The electron domain structure of the molecule is octahedral, but there is a lone electron pair 5 bonding pairs, so the molecular geometry is square pyramidal.
(d) Trigonal pyramidal
The electron domain structure of the molecule is tetrahedral, but there is one lone electron pair and three bonding pairs, so the molecular geometry is trigonal pyramidal.

3. -481 kJ
The balanced reaction that occurs is:
2H2(g) + O2(g) à 2H20(g)
Using the equation:
▲H° = Σ(bond energies of bonds broken) – Σ(bond energies of bonds formed)
We can solve for ▲H by substituting the bond energies in:
▲H° = Σ[(2)(H-H)+1(O-O)] – [(4)(O-H)]
▲H° = Σ[(2mol)(436kJ/mol)+(1mol)(499kJ/mol)] – [(4mol)(463kJ/mol)]
▲H° = (1371 kJ) – (1852 kJ)
▲H° = -481 kJ

Not only do bonds form between atoms in molecules, but they form between molecules as well. The bonds that hold molecules together are known as intermolecular bonds, some of which are the same as the bonds that hold atoms/ions together in a molecule.
1. Ionic bonds hold together the ions in an ionic solid with the same electrostatic attractions that hold together the ions that make up the salt. Ionic solids tend to bond in a very ordered, structured manner to form a crystal lattice, and the ionic bonds are very strong.
2. Network bonds, also known as covalent bonds, hold atoms together in a lattice of covalent bonds.








3. Metallic bonds