Chemical Bonding : Octet Rule, Lewis Dot Structure, Ionic Bonding & Covalent Bonding

Chemical bonding is the attraction between atoms at varying strength that allows the formation of a new substance.

The Stability of Noble Gases

Noble gases are defined as the group of chemical elements with low chemical reactivity or un-reactive at all. . The noble gases have all their electronic shells saturated and therefore cannot readily make compounds. Generally, Noble gases have eight electrons in their valence shell, thus, they are in accordance with the octet rule.

Octet Rule
It refers to the tendency of atoms to prefer to have eight electrons in the valence shell. When atoms have fewer than eight electrons, they tend to react and form more stable compounds.
In octet rule,  d or f electrons are not considered. Only the s and p electrons are involved in the octet rule. An octet corresponds to an electron configuration ending with s2p6.

Atoms will react to get in the most stable state possible. A complete octet is very stable because all orbitals will be full. Atoms with greater stability have less energy.
The other tendency of atoms is to maintain a neutral charge. Only the noble gases  have zero charge with filled valence octets. All of the other elements have a charge when they have eight electrons all to themselves.

           Atoms are more stable when they have no charge, or a small charge

Exceptions to the Octet Rule
1. Hydrogen and helium have only one electron shell. The first shell has only one s orbital and no p orbital, so it holds only two electrons. Therefore, these elements are most stable when they have two electrons.
Lithium, with three protons and electrons, is most stable when it gives up an electron.

2 Other notable exceptions are aluminum and boron, which can function well with six valence electrons ( less than an Octet ). Consider BF3. The boron shares its three electrons with three fluorine atoms. The fluorine atoms follow the octet rule, but boron has only six electrons. Although atoms with less than an octet may be stable, they will usually attempt to form a fourth bond to get eight electrons. BF3 is stable, but it will form BF4- when possible. Most elements to the left of the carbon group have so few valence electrons that they are in the same situation as boron: they are electron deficient. Electrons deficient elements often show metallic rather than covalent bonding.

3. In Period 3, the elements on the right side of the periodic table have empty d orbitals. The d orbitals may accept electrons, allowing elements like sulfur and phosphorus to have more than an octet. Compounds such as PCl5 and SF6 can form. These compounds have 10 and 12 electrons around their central atoms, respectively.

Xenon hexafluoride uses d-electrons to form more than an octet. This compound shows another exception: a noble gas compound.

4. Some elements, like nitrogen, have an odd number of electrons and will form somewhat stable compounds. Nitric oxide has the formula NO. No matter how electrons are shared between the nitrogen and oxygen atoms, there is no way for nitrogen to have an octet. It will have seven electrons instead. A molecule with an unpaired electron is called a free radical and radicals are highly reactive

Nitrogen dioxide has an unpaired electron. (Note the positive charge above the N).

Forming ions

Ions From Metal Elements
Generally metals lose electrons to form cations with a positive charge equal to the group number. This ties into the octet rule because the metals revert to the octet rule for the previous row in the periodic table.
Examples:
* Calcium ( Ca ) - forms into Ca2+ cation  since Ca belongs to Group 2A
                              and has the tendency to give up its 2 electrons
* Aluminum (Al ) - forms into Al3+ cation since Al belongs to Group 3A
                  and has the tendency to give up or lose its 3 electrons
* Potassium (K ) - forms into K1+ cation since K belongs to Group 1A and 
                              has the tendency to lose its 1 electron

Ions from nonmetal elements
Generally nonmetal atoms gain electron to form ions with a negative charge. These are called anions. This ties into the octet rule because the nonmetals gain electrons to fill to the octet for their row in the periodic table. The charge equals the group number minus 8.

Example:   What charge is expected for the ion formed by chlorine, Cl ?
Solution  : Charge = Group number – 8 Note:  Cl is under Group 7
                    = 7 – 8
         = - 1

Transition elements are in the d-block metals. They generally do not follow the octet rule so it is more difficult to predict their charge ( oxidation number).  They have electrons in the  “ (n-1)d ns” sublevels.
For example iron has 8 electrons in the 3d64s2 subshells. The number of electrons lost is variable. Iron can be stripped of the "4s" electrons and one of the "3d" electrons or only the "4s" electrons. The other transition metals like copper, titanium, etc show similar behavior

Lewis Dot Structure
A Lewis structure ( named after Gilbert N. Lewis ), also known as “electron dot diagram”, is a diagram showing how the valence electrons are arranged around an atom, ion or molecule.

Lewis Dot structures for Elements
For Main Group (Representative) elements, the number of valence electrons is the same as the Group A number . All Group 1A elements have one valence electron, all Group 2A elements have two valence electrons and so on all the way over to the Noble Gases in Group 8A with 8 valence electrons. The table below shows the Lewis Dot structure for the elements in Period (row) 2. It would be noticed that the number of valence electrons for each Group matches its Group A number.

Lewis Dot Structure ( LDS ) for Ions
The LDS can also be drawn for ions, recognizing that a negative charge means that extra electrons have been acquired and that a positive charge means that electrons have been lost. Lewis Dot Structures for ions are always enclosed in brackets with the charge indicated outside. Below are examples:

Ionic Bonding
Ionic bonding is the transfer of valence electron(s) between atoms. It is a type of chemical bond that generates two oppositely charged ions. In ionic bonds, the metal loses electrons to become a positively charged cation, whereas the nonmetal accepts those electrons to become a negatively charged anion.

By losing electrons, metals can achieve noble gas configuration and satisfy the octet rule. Similarly, nonmetals that have close to 8 electrons in their valence shells tend to readily accept electrons to achieve noble gas configuration. In ionic bonds, the net charge of the compound must be zero. 
Illustration 1:  Ionic bonding between Na and Cl
                    2 s2                                           2s2 3s2         4s2        
  Na       1s2           3s1                         Cl  1s2 3p6         4p5 
                    2p6                          2p6         3d10 


This sodium molecule donates the lone electron in its valence orbital in order to achieve octet configuration. This creates a positively charged cation due to the loss of electron. This chlorine atom receives one electron to achieve its octet configuration, which creates a negatively charged anion.

Illustration 2:  Ionic bonding between Mg and Chlorine Using the Lewis Dot
                      Structure
 
In this example, the magnesium atom is donating both of its valence electrons to chlorine atoms.  Each chlorine atom can only accept 1 electron before it can achieve its noble gas configuration; therefore, 2 atoms of chlorine are required to accept the 2 electrons donated by the magnesium.  Notice that the net charge of the compound is 0.

Lewis Dot Structure for Some  Ionic Compounds
Cations are formed when atoms lose electrons, represented by fewer Lewis dots, whereas anions are formed by atoms gaining electrons. The total number of electrons does not change. The table below shows how an ionic compound is illustrated by Lewis dot structure.

.
Ionic Compounds  and  Properties

Molecules that consist of charged ions with opposite charges are called ionic.
Ionic compounds are formed from metal and non-metal elements.

Examples of  ionic compound are :
      NaCl         BaO      Al2S3
       Sodium Chloride         Barium oxide       Aluminum sulfide

Some properties of ionic compound are  : they form crystal lattices ,have higher enthalpies of fusion and vaporization, hard, brittle, have high melting and boiling points, and conduct electricity in aqueous solution.

Formula of Ionic compound

The strong attraction between positive and negative ions often produces crystalline solids that have high melting points. Ionic bonds form instead of covalent bonds when there is a large difference in electronegativity between the ions. The positive ion, called a cation, is listed first in an ionic compound formula, followed by the negative ion, called an anion.
A balanced formula has a neutral electrical charge or net charge of zero.
Predicting  the  Type of Chemical Bond  Based on Electronegativity
If the electronegativity values of two atoms are:
• similar...
• Metallic bonds form between two metal atoms.
• Covalent bonds form between two non-metal atoms.
*Nonpolar covalent bonds form when the electronegativity values are     very similar.
     *Polar covalent bonds form when the electronegativity values are a       little further apart.
• different...
Ionic bonds are formed.


Determining the Formula of an Ionic Compound

A stable ionic compound is electrically neutral, where electrons are shared between cations and anions to complete outer electron shells or octets. The correct formula for an ionic compound is attained when the positive and negative charges on the ions are the same or "cancel each other out".
In writing ionic compound formula, determine first the charge on an ion. For the representative elements, the charge of the ion is related to the column or group that the element is in.
• Group IA elements have only one valence electron, so when they lose that electron they will have a +1 charge.
• Group IIA elements have two valence electrons. When they lose their two valence electrons they will have a +2 charge.
• Group IIIA elements have three valence electrons. They lose their three electrons to form +3 ions.
• Group IVA elements are somewhat of an exception to the trend. Tin (Sn) and lead (Pb) can lose multiple electrons to form differently charged ions. Carbon, silicon and germanium rarely form ions.
• Group VA elements have five valence electrons. Instead of losing these five electrons, Group VA elements will gain three electrons to have a -3 charge.
• Group VIA elements have six valence electrons and gain two electrons to have a -2 charge.
• Group VIIA elements have seven electrons in their outer shell. These elements gain one electron to have a -1 charge.

Transition metals are elements that are in groups IB to XB. These metals are capable of losing different numbers of electrons and can take multiple ionic forms. The names of transition metal ions contain Roman numerals to indicate the ions' charge.
For example, lead (II) nitrate contains a +2 lead ion (Pb2+ ). Vanadium (IV) oxide contains a +4 vanadium ion ( V4+).

Steps for writing and balancing the formula:

1. Identify the cation ( the portion with a positive charge). It is the least electronegative (most electropositive) ion. Cations include metals and they are often located on the left-hand side of the periodic table.
2. Identify the anion ( the portion with a negative charge). It is the most electronegative ion. Anions include halogens and nonmetals. Keep in mind, hydrogen can go either way, carrying either a positive or negative charge.
3. Write the cation first, followed by the anion.
4. Adjust the subscripts of the cation and anion so the net charge is 0. Write the formula using the smallest whole number ratio between the cation and anion to balance charge.
5. If the charges of the cation and anion are equal (e.g., +1/-1, +2/-2, +3/-3), then combine the cation and anion in a 1:1 ratio. An example is potassium chloride, KCl. Potassium (K+) has a 1- charge, while chlorine (Cl-) has a 1- charge. Note that you do not ever write a subscript of 1.
6. If the charges on the cation and the anion are not equal, add subscripts as needed to the ions to balance the charge. The total charge for each ion is the subscript multiplied by the charge. Adjust the subscripts to balance charge. An example is sodium carbonate, Na2CO3. The sodium ion has a +1 charge, multiplied by the subscript 2 to get a total charge of 2+. The carbonate anion (CO3-2) has a 2- charge, so there is no additional subscript.
7. If you need to add a subscript to a polyatomic ion, enclose it in parentheses so it is clear the subscript applies to the entire ion and not to an individual atom. An example is aluminum sulfate, Al2(SO4)3. The parenthesis around the sulfate anion indicates three of the 2- sulfate ions are needed to balance 2 of the 3+ charged aluminum cations.
Examples of ionic compounds
Many familiar chemicals are ionic compounds. A metal bonded to a nonmetal is a dead giveaway that you're dealing with an ionic compound. Examples include salts, such as table salt (sodium chloride or NaCl) and copper sulfate (CuSO4).

Covalent Bonding

Covalent bonding is the sharing of valence electrons between atoms. This type of bonding occurs between two atoms of the same element or of elements close to each other in the periodic table. This bonding occurs primarily between nonmetals; however, it can also be observed between nonmetals and metals. Thus, covalent or molecular compound is formed by the sharing of electrons between atoms.

If atoms have similar electronegativities (the same affinity for electrons), covalent bonds are most likely to occur. Because both atoms have the same affinity for electrons and neither has a tendency to donate them, they share electrons in order to achieve octet configuration and become more stable. In addition, the ionization energy of the atom is too large and the electron affinity of the atom is too small for ionic bonding to occur.  For example: carbon does not form ionic bonds because it has 4 valence electrons, half of an octet. To form ionic bonds, Carbon molecules must either gain or lose 4 electrons. This is highly unfavorable; therefore, carbon molecules share their 4 valence electrons through single, double, and triple bonds so that each atom can achieve noble gas configurations. Covalent bonds include interactions of the sigma and pi orbitals; therefore, covalent bonds lead to formation of single, double, triple, and quadruple bonds.

Covalent bonds form between non-metal atoms. Each bond consists of a shared pair of electrons, and is very strong.

Lewis Dot Structure of  Molecules

The Lewis structure is used to represent the covalent bonding of a molecule or ion. Covalent bonds are a type of chemical bonding formed by the sharing of electrons in the valence shells of the atoms. Covalent bonds are stronger than the electrostatic interactions of ionic bonds.  Most bonding is not purely covalent, but is polar covalent (unequal sharing) based on electronegativity differences.
The atoms in a Lewis structure tend to share electrons so that each atom has eight electrons (the octet rule). The octet rule states that an atom in a molecule will be stable when there are eight electrons in its outer shell.  Lewis structures display the electrons of the outer shells because these are the ones that participate in making chemical bonds.
For simple molecules, the most effective way to get the correct Lewis structure is to write the Lewis diagrams for all the atoms involved in the bonding and adding up the total number of valence electrons that are available for bonding. The element which is the least electronegative is the center of the compound.

Example 1:  Draw the Lewis Dot  structure for water ( H2O )

1.  Hydrogen  atoms are  always placed on the outside of  the molecule  making
                oxygen as the central atom.
2.  One dot ( 1 valence e- ) is assigned for each hydrogen and 6 dots ( 6 valence
                e- ) for oxygen.  There are now  a  total of  8 electrons  which  conform  to  the
                octet rule.




3.  Each o f the  two  unpaired  electrons  of  the  oxygen  atom  will  form a bond                 
                 with one of the unpaired electrons of the hydrogen atoms. The bonds  formed     
                by the shared electron pairs can  be represented by  either two  closely places
                dots  between  two  element  symbols or  more  commonly   by  a  straight  line
                between element symbols:
   



Example 2:  Draw  the Lewis structure for methane (CH4).

                 Hydrogen atoms are always placed on the outside of the molecule, so carbon  should be the central atom.


There is now a total of 8  valence electrons  [4 from carbon + 4(1 from each hydrogen] .

Lewis Structures of Polyatomic Ions

Building the Lewis Structure for a polyatomic ion can be done in the same way as with other simple molecules, but we have to consider that we will need to adjust the total number of electrons for the charge on the polyatomic ion. If the ion has a negative charge, the number of electrons that is equal to the charge on the ion should be added to the total number of valence electrons. If the ion has a positive charge, the number of electrons that is equal to the charge should be subtracted from the total number of valence electrons. After writing the structure, the entire structure should then be placed in brackets with the charge on the outside of the brackets at the upper right corner.

Example 1:   Write the Lewis structure for the ammonium ion (NH4+).
            Hydrogen atoms are always placed on the outside of the molecule, so nitrogen should be the central atom.

After counting the valence electrons, we have a total of 9 [5 from nitrogen + 4(1 from each hydrogen)] = 9. The charge of +1 means an electron should be subtracted, bringing the total electron count to 8.
Each hydrogen atom will be bonded to the nitrogen atom, using two electrons. The four bonds represent the eight valence electrons with all octets satisfied, so the structure is complete. (Do not forget to put  brackets and charge on the outside of the brackets)

Example 2:   Write the Lewis structure for the hydroxide ion (OH-).
Since there are only two atoms,  begin with just a bond between the two atoms.

After counting the valence electrons, there is  a total of 7 [6 from Oxygen + 1 from each Hydrogen)] = 7. The charge of -1 indicates an extra electron, bringing the total electron count to 8.
Oxygen will be bonded to the hydrogen, using two electrons. Place the remaining six electrons as three lone pairs on the oxygen atom. All octets are satisfied, so the structure is complete. (Do not forget to put  brackets and  charge on the outside of the brackets)

Example 3:  Draw the Lewis structure  for nitrate ion ( NO3 - )
Nitrogen is the least electronegative atom and should be the central atom.

After counting the valence electrons, we have a total of 23[5 from nitrogen + 3(6 from each oxygen)] = 23. The charge of -1 indicates an extra electron, bringing the total electron count to 24.
Each oxygen atom will be bonded to the nitrogen atom, using a total of six electrons. We then place the remaining 18 electrons initially as 9 lone pairs on the oxygen atoms (3 pairs around each atom).

Although all 24 electrons are represented in the structure (two electrons for each of the three bonds and 18 for each of the nine lone pairs), the octet for the nitrogen atom is not satisfied. To satisfy the octet rule for the nitrogen atom, a double bond needs to be made between the nitrogen and one of the oxygen atoms. Because of the symmetry of the molecule, it does not matter which oxygen atoms is chosen. Because there are three different oxygen atoms that could form the double bond, there will be three different resonance structures showing each oxygen atom with a double bond to the nitrogen atom. Double-headed arrows will be placed between these three structures. (Do not forget your brackets and to put your charge on the outside of the brackets)

Naming and Formulas for Molecular or Covalent Compounds

Covalent compounds are named in different ways than  ionic compounds . Simple covalent compounds are generally named by using prefixes to indicate how many atoms of each element are shown in the formula. Also, the ending of the last (most negative) element is changed to -ide.
Prefixes which indicate the number of atoms of each element are used in the naming of inorganic molecular compounds. Below are the prefixes used in naming covalent compound.

mono- = 1 di- = 2 tri- = 3 tetra- = 4 penta- = 5
hexa- = 6 hepta- = 7 octa- = 8 nona- = 9 deca- = 10

When given a formula the prefixes above are applied to the words that would be used to name the compound as if it were ionic. For example, P2O3 would be named Phosphorous Oxide if it were ionic, but it consists of two nonmetals, so it would be named Diphosphorous Trioxide.

Whenever there is only one atom of the first element in a formula ,  drop the term Mono-. For example CO is Carbon Monoxide, not Monocarbon Monoxide.

To write formulas ,  just interpret the prefixes on the names and write the appropriate symbolic representation. For example, Sulfur Dioxide is SO2.

Examples:
1.  Write the molecular formula of the following:
a.   Carbon Tetrachloride - CCl4
b.   Trinitrogen Pentoxide - N3O5
2.  Give the name of the molecular compound as shown by the molecular formula:
a.   N2O - Dinitrogen Oxide
b.   CO2 - Carbon Dioxide

There are special cases where we use common names for molecular compounds. Examples of these are  Water = H2O and Ammonia = NH3 (not to be confused with the ammonium ion = NH4+1