Covalent Bonding and Molecular Structure | General Chemistry 1

Covalent bonds:

Covalent bonds are chemical bonds formed by the sharing of pairs of electrons between atoms. This sharing allows each atom to achieve a stable electron configuration, typically resembling the nearest noble gas.

Bond length and bond strength:

Bond length: 

• Bond length is the distance between the nuclei of 2 covalently bonded atoms. 
• Bond length decreases as the bond multiplicity (number of shared electron pairs) increases: single bond > double bond > triple bond. It also depends on the size of the atoms involved; larger atoms form longer bonds

Bond strength: 

• Bond strength is the energy required to break a bond.
• Bond strength increases with bond multiplicity: single bond < double bond < triple bond. For bonds of the same type, shorter bonds are generally stronger.

Covalent compounds:

Covalent compounds, also known as molecular compounds, consist of individual molecules where atoms share electrons through covalent bonds.

Properties:

• Covalent compounds do not form ions in solution.
• They are poor conductors of electricity in both solid and liquid states.
• They often have lower melting and boiling points compared to ionic compounds.

Chlorine molecule (Cl₂):

Cl: Z = 17 ⇒ 1s² 2s² 2p⁶ 3s² 3p⁵ ⇒ 7 valence electrons

If 2 chlorine atoms each share an electron:

•  they have 8 electrons in their outer shell
• they reach the electron configuration of argon 

⇒ formation of Cl₂

Binary molecular compounds:

1. The name of the first element in the formula is given first, using the element's full name.
2. The second element's name is modified to end with the suffix "-ide".
3. Prefixes are used to indicate the number of atoms of each element in the compound: mono- = 1, di- = 2, tri- = 3, tetra- = 4 …. The prefix mono- is generally omitted for the first element.

The final "a" or "o" of a prefix is often dropped if the element name begins with a vowel to make the name easier to pronounce.
 

CO₂ = carbon dioxide
PCl₅ = phosphorus pentachloride
N₂O₄ = dinitrogen tetroxide

Binary acids:

Binary acids are composed of hydrogen and a nonmetal anion. These acids are named based on the nonmetal anion they contain:

• by adding the prefix hydro- in front of the name of the anion
• by replacing the -ide ending of the anion with -ic
• by adding the word acid at the end
   

HCl = hydrochloric acid
HBr = hydrobromic acid

Oxoacids:

Oxoacids are acids that contain hydrogen, oxygen, and another element (the central atom). These acids produce hydrogen ions (H⁺) and oxoanions when dissolved in water.

• If the oxoanion ends in -ate, the acid name ends in -ic acid.
• If the oxoanion ends in -ite, the acid name ends in -ous acid.

NO₃^- = nitrate ⇒ HNO₃ = nitric acid
NO₂^- = nitrite ⇒ HNO₂ = nitrous acid
SO₄^2- = sulfate ⇒ H₂SO₄ = sulfuric acid

Octet rule:

The octet rule states that main-group elements tend to form molecules by gaining, losing, or sharing electrons to achieve a stable electron configuration with eight electrons in their valence shell, resembling the electron configuration of a noble gas.

The octet rule is crucial for understanding the bonding behavior of elements, especially those in the second period of the periodic table, which includes carbon, nitrogen, and oxygen. This rule is foundational in the study of organic chemistry.

Carbon (C):

Atomic Number (Z): 6 ⇒ Electron Configuration: 1s² 2s² 2p² ⇒ Valence Electrons: 4 ⇒ Carbon needs 4 more electrons to complete its octet. It achieves this by forming 4 covalent bonds with other atoms, making it tetravalent.
Example: In methane (CH₄), carbon forms 4 single covalent bonds with hydrogen atoms to complete its octet.

Oxygen (O):

Atomic Number (Z): 8 ⇒ Electron Configuration: 1s² 2s² 2p⁴ ⇒ Valence Electrons: 6 ⇒ Oxygen needs 2 more electrons to complete its octet. It achieves this by forming 2 covalent bonds.
Example: In water (H₂O), oxygen forms 2 single covalent bonds with hydrogen atoms, and it has 2 lone pairs of electrons.

Lewis structure:

A Lewis structure is a visual representation of the arrangement of atoms and the distribution of all valence electrons in a molecule or polyatomic ion. It helps in understanding how atoms are bonded in a molecule and how electrons are shared or distributed.

• Dot: Represents one non-bonding electron.
• Pair of dots: Represents a lone electron pair (lone pair).
• Line: Represents a pair of shared electrons (a bond).

• Lewis structure of NH₃: 

N: 5 valence electrons ⇒ needs 3 shared electrons ⇒ 3 covalent bonds​​​.
H: 1 valence electron ⇒ needs 1 shared electron ⇒ 1 covalent bond.

 

• Lewis structure of CO₂:

C: 4 valence electrons ⇒ needs 4 shared electrons ⇒ 4 covalent bonds.
O: 6 valence electrons ⇒ needs 2 shared electrons ⇒ 2 covalent bonds.

How to write a Lewis structure:

1. Count the total number of valence electrons. Add or subtract electrons if you have a negative or a positive charge.
2. Determine the number of covalent bonds / lone pairs the molecule will have by dividing the number of valence electrons by 2. Use the molecular formula to draw the skeletal structure.
3. Distribute the remaining valence electrons to satisfy the octet rule, completing the octet of the more electronegative atoms first. Include double or triple bonds if necessary.
4. Check the number of valence electrons in the drawn molecule.
5. Assign formal charges to all atoms.

Draw a Lewis structure for methanol CH₃OH:

Count the valence electrons:

• Carbon: 4 valence electrons
• Hydrogen: 1 valence electron (4 H atoms ⇒ 4 valence electrons)
• Oxygen: 6 valence electrons
• Total = 14 valence electrons ⇒ 7 pairs (bonds and lone pairs)

Determine the number of covalent bonds: 14 valence electrons ⇒ 7 pairs (bonds and lone pairs)

Arrange the atoms

 

Add bonds ...

5 bonds ⇒ 10 electrons 

... then lone pairs

5 bonds + 2 lone pairs ⇒ 14 e^- 

Formal charge:

The formal charge of an atom in a molecule is the charge assigned to the atom assuming that electrons in all chemical bonds are shared equally between atoms, regardless of relative electronegativity.
 

Calculation of formal charge:

The formal charge (FC) is calculated using the formula:
 

Formal Charges in the CH₃NO₂ Molecule:

• N: 1s² 2s² 2p³ = [He] 2s² 2p³ ⇒ 5 valence electrons in free atom ⇒ V = 5
  0 non-bonding electron ⇒ N = 0
  4 bonds:  8 shared electrons ⇒ B = 8
  Formal charge FC = 5 - 0 - 4  = + 1
• O: 1s² 2s² 2p⁴ = [He] 2s² 2p⁴ ⇒ 6 valence electrons in free atom ⇒ V = 6
  2 lone pairs ⇒ 4 non-bonding electrons ⇒ N = 4
  2 bonds: 4 shared electrons ⇒ B = 4
  Formal charge FC = 6 - 4 - 2 = 0
• O: 1s² 2s² 2p⁴ = [He] 2s² 2p⁴ ⇒ 6 valence electrons in free atom ⇒ V = 6
  3 lone pairs ⇒ 6 non-bonding electrons ⇒ N = 6
  1 bond: 2 shared electrons ⇒ B = 2
  Formal charge FC = 6 - 6 - 1 = - 1

Resonance structures:

Resonance structures are a set of two or more valid Lewis structures for a single molecule that cannot be accurately represented by a single Lewis structure. These structures are used to depict delocalized electrons (π or nonbonded electrons) within molecules, where the actual distribution of electrons is a hybrid of all possible resonance forms. The single bonds remain the same, but the positions of the multiple bonds and nonbonded electrons differ.
 

Key points for resonance structures:

• Same placement of atoms: Resonance structures have the same arrangement of atoms but differ in the placement of electrons, particularly π electrons or nonbonded (lone pair) electrons.
• Valid Lewis structures: Each resonance structure must be a valid Lewis structure, adhering to the rules of valence electrons and octet rule (where applicable).
• Resonance hybrid: The actual molecule is represented by a resonance hybrid, which is a weighted average of all possible resonance structures. 

  The different resonance forms of a substance are not all equal: the form with the most bonds and fewer charges has a higher contribution to the resonance hybrid

Criteria for evaluating resonance structures:

When multiple resonance structures are possible, the following guidelines help in identifying the most significant contributors:

• Minimize formal charges
• Place negative charges on electronegative atoms
• Avoid adjacent formal charges
• Maximize covalent bonds

Incomplete Octet:

Some elements are stable with fewer than eight electrons in their valence shell. These are usually small atoms found early in the periodic table such as boron (B) and beryllium (Be).
 

• Beryllium chloride (BeCl₂)
  Each beryllium atom forms 2 single covalent bonds with chlorine atoms, resulting in only 4 electrons in its valence shell.
• Boron trifluoride (BF₃)
  Boron forms 3 single covalent bonds with fluorine atoms, resulting in 6 electrons in its valence shell.

Odd Number of Electrons (Free Radicals)​​​​​​​:

A free radical is a chemical species with an odd number of valence electrons, formed by homolytic cleavage, containing one or more unpaired electrons. Free radicals are highly reactive because the unpaired electron tends to form a covalent bond with an unpaired electron on another molecule.
 

• Methyl radical (CH₃): Carbon has 7 electrons in its valence shell, with one unpaired electron.
• Nitric oxide (NO): Nitrogen has 7 electrons in its valence shell, with one unpaired electron.

Expanded valence shell

Atoms with a valence shell n $\ge$ 3 can have more than eight electrons in their valence shell. This is because they also have d orbitals that can be used in bonding to form an expanded octet.
 

Phosphorus pentachloride (PCl₅):

P is in the 3^rd period of the periodic table and therefore has 3s, 3p and 3d orbitals. The 3d orbitals of phosphorus in PCl₅ are used and form an expanded octet ⇒ P has 10 electrons in its outer shell.