Chapter 4: Covalent Compounds, Formulas, and Structures

Introduction to Covalent Bonds

In contrast to ionic bonding, where electrons are transferred from atom to atom, covalent bonding occurs when electrons are shared between two or more atoms.

• Whereas ionic bonding is simply electrical attraction of oppositely charged ions, in covalent compounds the atoms are physically attached to each other.

Molecules - Compounds composed of covalently bonded groups of atoms are called molecules.

Lewis Dot Diagrams - Chemists draw structures of covalent molecules using Lewis electron-dot structures to demonstrate the sharing of electrons in covalent bonding.

• In Lewis structures, valence electrons are represented by dots arranged around the atomic symbol.

Covalent Bonds - When electrons are shared between two atoms, one atom donates one electron while the other atom donates the second electron and the shared pair of electrons represents a covalent bond.

• Double Bond - If two pairs of electrons are shared between two atoms, a double bond exists.
• Triple Bond - When two atoms share three pairs of electrons, a triple bond exists.
• Single bonds have the lowest energy, double bonds have a higher energy (A double bond is almost twice as strong as a single bond),  and triple bonds have the highest energy.
• Single bonds are the longest, double bonds are shorter and triple bonds are the shortest
• Covalent compounds are governed by the octet rule.
• In a covalent bond each atom contributes the same number of electrons to the bond.
• In a covalent molecule each atom considers the shared electrons (those electrons between the two atoms) as its own.

Octet Rule - The octet rule dates that the noble gas configuration will be achieved if the Lewis structure shows 8 electrons around each atom.

• Hydrogen is an exception; its "octet" is two electrons.

Paramagnetic - Substances that contain unpaired electrons are said to be paramagnetic.

• Magnetic properties of atoms, ions, and molecules arise from unpaired electrons. The electron spin generates a small magnetic field which is canceled when electrons are paired.
• When they are unpaired, they impart magnetic properties to the substance.
• The more unpaired electrons, the stronger the magnetic effect.

Diamagnetic - Diamagnetic substances have all of their electrons paired and have very small magnetic fields.

With small molecules and polyatomic ions, Lewis structures can be drawn using the basic octet rule. But with larger molecules it is necessary to first determine a general skeleton that consists of a central atom with surrounding atoms bound into it.

1. Carbon is usually a central atom in the structure and in compounds with more than one carbon atom, the carbon atoms are joined in a chain to start the skeleton
2. Hydrogen is never a central atom because it can form only one covalent bond.
3. Halogens form only a single covalent bond when oxygen is not present and therefore a halogen will generally not be a central atom.
4. Oxygen forms only two covalent bonds and is rarely a central atom. However it may link to carbon atoms in a carbon chain.
5. In the simpler molecules, the atom that appears only once in the formula will often be the central atom.

Bonding Pairs - When a pair of electrons is placed between two atoms in the skeletal structure, they represent a covalent bond.

• To simplify a Lewis Dot Structure, the bonding pairs of electrons can be replaced by a line representing a bond. A single line represents a single bond, two lines represent a double bond and three Lines represent a triple bond.

Nonbonding Pairs - The remaining electrons that are used to complete the octet of all the other atoms in the skeleton are referred to as the nonbonding pairs of electrons.

• Nonbonding pairs are often referred to as lone pairs.

All valence electrons of the atoms within a neutral covalent molecule must add up to a zero.

• If the substance is a polyatomic ion, we must take into account that the electrons are used to form the charge of the ion. All the valence electrons of the atoms should add up to the charge of the ion.

When writing Lewis structures for ions, it is necessary to enclose the ion in brackets and indicate the charge of the ion as shown.

Some compounds have formulas in which the total number of valence electrons is an odd number, making it impossible to construct a Lewis structure with an octet around each atom.

Free Radicals - Molecules that have Lewis structures with an unpaired electron are often called free radicals.

• The unpaired electron makes the molecule unusually reactive.

Formal Charge - The formal charge is the difference between the number of electrons an atom has in a Lewis structure and its number of valence electrons.

• Formal charges are used to determine the validity of a Lewis structure.
• The smaller the formal charge, the more stable the Lewis structure is.
• When one has several possible Lewis structures for a singular molecule, formal charge can be used to determine which one is the most stable and therefore the most correct structure.

Resonance Structures - Resonance structures are equally probable Lewis structures of a given molecule.

• In the case of resonance structures, The Lewis structures can be totally equivalent, even down to the formal charges on the atoms.
• Most resonance structures are very similar for a given substance, usually differing only in the geometry of the molecule or ion.
• It is found experimentally that None of the resonance Lewis structures properly describe the molecule; the two properties of the substance are found by blending all of the resonance structures together.

Polar Bond - If the electrons are not shared equally by two atoms, they will spend more time localized near one atom or the other, creating a polar bond with a positive and a negative end.

• This imbalanced sharing of the electrons within a covalent bond is caused by electronegativity; the more electronegative atom attracts more electrons and experiences a more negative dipole moment.

Nonpolar Bond - If the atoms have the same electronegativity (usually when they are the same element), then the bond is considered to be nonpolar.

• Electrons are only shared equally in a nonpolar covalent bond.

The atom that attracts electrons will have a partial negative charge and the other atom will have a partial positive charge.

Bond Order - Bond order is a term that refers to the average number of bonds that an atom makes in all of its bonds to other atoms .

Bond Energy - The strength of a covalent bond is expressed as a bond energy.

Molecular Geometry

Once a valid Lewis structure has been determined, the overall geometry and three-dimensional shape of covalently bonded molecules can be found using the VSEPR theory.

Valence-shell electron-pair repulsion theory (VSEPR theory) - Electron pairs will repel each other since all electrons carry a negative charge. In fact they will repel each other so that they are as far apart as well.

• Bonding Domain - The region in space occupied by a bonding pair (or pairs for double and triple bonds).
• Nonbonding Domain - The region in space occupied by a non-binding electron pair.

In the case of structures that have no nonbonding domain, there are six basic geometric shapes found around an atom that may have one to six atoms bonded to a central atom.

Derived Structures - Derived structures have one or more of the bonding domains replaced with nonbonding domains.