Covalent Compounds

Covalent compounds exhibit the greatest variety among compounds.
Most substances in biological cells are covalent compounds, including water, proteins, carbohydrates, DNA, amino acids, and RNA.
These biological molecules largely consist of covalent compounds.
A covalent compound comprises identical molecules where nonmetal atoms are joined by covalent bonds.

Nature and Formation of Covalent Bonds

Covalent bonds link two or more nonmetal atoms.
Formation and nature of covalent bonds.
Writing molecular formulas and nomenclature for covalent compounds.

Molecules of Identical Nonmetal Atoms

Some elements exist as molecules made of identical nonmetal atoms, such as diatomic gases: hydrogen (H2), nitrogen (N2), oxygen (O2), fluorine (F2), chlorine (Cl2), bromine (Br2), iodine (I_2).
These are elements with covalent bonds because they consist of the same element.
Covalent compounds generally consist of two or more different nonmetals.

Examples of Covalent Compounds

Hydrogen fluoride (HF) consists of hydrogen and fluorine, both nonmetals, making it a covalent compound.
Covalent compounds vary in size; HF is a small molecule.
Proteins, large covalent compounds, contain thousands of carbon, hydrogen, oxygen, nitrogen, and sulfur atoms.

Formation of Covalent Bonds

Recall ionic bonds where metals lose electrons to achieve a noble gas electron configuration, filling energy levels.
Nonmetals gain these electrons to complete their octet.
Covalent bonds occur between two nonmetals, both needing to gain electrons.
Neither atom wants to lose electrons, so stability is achieved through sharing.

Sharing of Electrons

Hydrogen has one valence electron and needs to gain one electron to achieve stability (like Helium with two electrons).
In the absence of a metal to donate electrons, nonmetals share electrons.
Hydrogen shares its electron with another hydrogen atom.
By sharing, each hydrogen atom effectively has a full energy level.
Sharing creates an interaction or force known as a covalent bond.
Covalent bonds are formed by sharing electrons.

Representation of Covalent Bonds

A line represents a covalent bond, indicating that two atoms are sharing one electron each.
A covalent bond can be depicted with a line or two dots, where each dot represents an electron.
Each electron belongs to one atom, but they are shared to achieve stability, keeping the atoms close and forming the covalent bond.

Nomenclature of Binary Covalent Compounds

Naming covalent compounds is simpler than naming ionic compounds.
Binary covalent compounds consist of two different nonmetal elements.
If a compound contains a metal, it is ionic, not covalent.

Prefixes in Covalent Compound Nomenclature

Prefixes indicate the number of atoms in the compound:
- 1: mono
- 2: di
- 3: tri
- 4: tetra
- 5: penta
- 6: hexa
- 7: hepta
- 8: octa
- 9: nona

Naming Formula

General formula: prefix + name of first nonmetal + prefix + name of second nonmetal with "-ide" ending.
Example: N2O5 is dinitrogen pentoxide.

Exceptions and Tricks

When two vowels are adjacent, they combine (e.g., pentoxide instead of pentaoxide).
The prefix "mono" is not used for the first element in the compound.

Practice Naming Covalent Compounds

CO2: Carbon dioxide (mono is skipped for the first element).
SF_6: Sulfur hexafluoride (mono is skipped for the first element).
N_2O: Dinitrogen monoxide (monoxide instead of monooxide).

Writing Molecular Formulas

Writing molecular formulas from names is straightforward.
Example: Nitrogen trifluoride: NF_3 (tri means three fluorine atoms).
Example: Sulfur trioxide: SO_3 (tri means three oxygen atoms).
Example: Diphosphorus pentoxide: P2O5 (di means two phosphorus atoms, penta means five oxygen atoms).

Identifying Ionic vs. Covalent Compounds

Determine if a formula represents an ionic or covalent compound.
Ionic compounds are formed from a metal and a nonmetal.
Covalent compounds are formed from two nonmetals.
If the first element in the formula is a metal, the compound is ionic; if it is a nonmetal, the compound is covalent.
Ionic compounds do not use prefixes like di, tri, or tetra in their names.

Examples

AlCl_3: Aluminum chloride (ionic, no prefixes).
CS_2: Carbon disulfide (covalent, uses prefixes).
FeO: Iron(II) oxide (ionic, no prefixes, iron is a transition metal, indicating oxidation state).
N2O4: Dinitrogen tetroxide (covalent, uses prefixes).

Chemical Structures

Molecular formulas indicate the number of atoms of each element but not their arrangement.
Chemical structure provides information about the arrangement of atoms in a molecule.

Lewis Dot Structures

Lewis dot structures determine the chemical structure of simple molecules.
Bonds form between atoms to fill their valence energy levels.
Most elements need eight electrons in their valence energy level (octet rule), except for hydrogen, which needs two.
Covalent bonds are formed when two electrons are shared.
One covalent bond equals two shared electrons.

Valence Electrons and Lewis Dot Structures

For main group elements, valence electrons are easily determined by the group number.
Hydrogen (Group 1A) has one valence electron.
Carbon (Group 4A) has four valence electrons.
Oxygen (Group 6A) has six valence electrons.

Determining Number of Bonds

Atoms form bonds to achieve a full valence energy level (octet rule).
Carbon needs four more electrons, so it makes four bonds.
Oxygen needs two more electrons, so it makes two bonds.
Elements in Group 5A make three bonds, and elements in Group 7A make one bond.

Lewis Structure of Methane (CH_4)

Carbon (four valence electrons) shares with four hydrogen atoms (one valence electron each).
Each hydrogen atom shares one electron with carbon, forming a bonding pair.
Carbon achieves an octet, and each hydrogen completes its duet, resulting in stability for both.

Lewis Structure of Ammonia (NH_3)

Nitrogen (Group 5A) has five valence electrons; three are unpaired, and one is a lone pair.
Three hydrogen atoms share their electrons with the unpaired electrons of nitrogen.
Nitrogen forms three bonding pairs with hydrogen and retains a lone pair of electrons.
Nitrogen achieves an octet, and each hydrogen has a duet. So we have three covalent bonds and one lone pair.

Lewis Structure of Water (H_2O)

Oxygen (Group 6A) has six valence electrons, with two unpaired electrons.
Two hydrogen atoms share their electrons with the unpaired electrons of oxygen.
Oxygen forms two bonds (bonding electrons) and has two nonbonding pairs (lone pairs).
There are two bonding pairs and two lone pairs around the oxygen atoms.

Single, Double, and Triple Bonds

Single bond: one shared pair of electrons.
Double bond: two shared pairs of electrons.
Triple bond: three shared pairs of electrons.

Single Bonds

Chlorine needs to share one electron, forming a single bond with hydrogen to create HCl.

Double Bonds

Oxygen needs to share two electrons. If no hydrogen is available, two oxygen atoms share electrons and form bonds (O_2).
Two shared pairs of electrons form a double bond.
Each oxygen atom has eight electrons in the overlap region.

Triple Bonds

Nitrogen (Group 5A) has five valence electrons and is not stable alone.
Two nitrogen atoms share three pairs of electrons, forming a triple bond in N_2.
Triple bonds are stronger than double bonds, and double bonds are stronger than single bonds. However, the relationship is not linear (i.e., a double bond is not twice as strong as a single bond).

Examples of Double and Triple Bonds

Formaldehyde: Contains a carbon-oxygen double bond.

Number of Bonds and Nonbonding Electrons

Atoms in the same group generally form the same number of bonds and have the same number of nonbonding electrons.
Carbon typically forms four bonds with no nonbonding electrons.
Oxygen typically forms two bonds with two nonbonding pairs of electrons.

Writing Lewis Structures

Problem: Write the Lewis structure of carbon tetrachloride (CCl_4).

Determine the total number of valence electrons in the molecule.
- Carbon: 1 * 4 = 4 valence electrons.
- Chlorine: 4 * 7 = 28 valence electrons.
- Total = 32 valence electrons.
Arrange the molecules symmetrically: Carbon is the central atom, surrounded by four chlorine atoms.
Connect each atom with a single covalent bond (two electrons each).
- 32 - (4 * 2) = 24 electrons remaining.
Ensure the central atom (carbon) has an octet.
Distribute the remaining electrons to the side atoms (chlorine) until they achieve an octet.

Step to Completion

Use the remaining 24 electrons to complete the octets of the chlorine atoms (3*2 * 4 = 24).

Each chlorine atom now has eight electrons and carbon has eight electrons.

Therefore, the Lewis structure is complete.
Bonding electrons are shown as covalent bonds, and nonbonding electrons are shown as dots.

Central Atom Short of Octet

If the central atom is short of an octet, a nonbonding pair of electrons from a side atom can be used to make a double or triple bond.
Write the Lewis structure of selenium dioxide (SeO_2).

Count valence electrons:
- Selenium (Group 6A): 6 valence electrons.
- Oxygen (Group 6A): 2 * 6 = 12 valence electrons.
- Total = 18 valence electrons.
Arrange the molecule symmetrically: Selenium is the central atom with two oxygen atoms on the sides.
Connect with single bonds:
- Use two electrons for each Se-O bond. So there are 4 electrons in total.
- 18 - 4 = 14 electrons remaining.
Give remaining electrons to side atoms until their octets are complete (one pair at a time).
Distribute 12 of the 14 remaining electrons around the oxygen atoms so those have 8 electrons.
Selenium is still short of an octet (only six electrons around it).
Form a double bond by of the lone pairs of electrons. Oxygen does not lose it’s octet, but gains its octet.
*Now all the atoms in structure has complete octet. So it’s a stable structure.

Lewis Structure of Ammonia (Again)

Ammonia (NH_3):
- Nitrogen (Group 5A): 5 valence electrons.
- Hydrogen (Group 1A): 3 * 1 = 3 valence electrons.
- Total = 8 valence electrons.
Nitrogen is the central atom, surrounded by three hydrogen atoms.

Use two electrons to make a covalent bond on each side and two additional electrons on the nitrogen.
Nitrogen is stable because Nitrogen has 8 valence electrons and the two Hydrogens each have 2 valence electrons.
All elements are stable.
This concludes the Lewis structure.

Lewis Structure of Hydrogen Sulfide (H_2S)

Two hydrogen atoms (1 electron each):
- Total: 2 * 1 = 2 valence electrons
Sulfur (Group 6A): 6 valence electrons
Total: 8 valence electrons
Sulfur is the central atom, flanked by two hydrogen atoms.
Each hydrogen is bonded bonded to sulfur via a single covalent bond and four electrons are on the sulfur.

So Hydrogen forms a duet with sulfur and Sulfur has two lone pairs and two covalent bonds.All elements are stable.

Octet Rule Exceptions

Nonmetals in period two (C, N, O, F) always follow the octet rule.
Nonmetals from period three (phosphorus and sulfur) can form compounds that do not obey the octet rule.
Phosphorus and sulfur can have more than eight valence electrons (expanded octets).
Adenosine triphosphate (ATP), DNA, and RNA contain phosphorus atoms with expanded octets.