In this chapter, we will learn to:

- Understand the concept of chemical instability
- Distinguish between two types of chemical bonds: covalent and ionic
- Predict the type of a bond
- Compare the two bonds

note that this chapter requires a knowledge of electron configurations and distribution. To review that chapter, CLICK HERE.

1. Chemical Stability

     The connection between the microscopic electron configuration and the macroscopic property of the element lies in the number of electrons found on the outer energy level of the atom of the element.

 

1.1 Stability of Inert Gases

     Studies of inert gases revealed that their atoms contain eight electrons in their outer energy level ( except for helium, which contains two electrons). Given that fact, it can be deduced that they are chemically stable and do not react readily with other elements.

1.2 Valence Electrons

     The electrons found on the outer energy level (called valence level) of the atom of an element have an important role in determining the properties of the element. They are called valence electrons. In fact, the number of valence electrons of an element is related to its position in the periodic table - it is equal to the unit digit of the column of a main-group element. For example all of the elements in column 1A have 1 valence electron, while elements of column 6Ahave 6 valence electrons. It is also important to mention that the valence of an element indicates the number of electrons involved in bonding to acheive the electron configuration of the closest inert gas. It is also known as valency.

1.3 Lewis Electron-dot Symbol

     As Lewis electron-dot symbol consists of the symbol for the element surrounded by dots, one for each valence electron. By convention, the first four electron dots are placed as single dots around the four sides of the element symbol. When an atom has more than four valence electrons, the first four single dots are placed around the symbol of the element, and then each single dot is paired respectively, until all available electrons are used. The number of dots placed around the symbol of the elements is equal to the number of valence electrons.

1.4 The Octet Rule

     The inert gases as mentioned earlier do not combine readily with other elements. One explanation for their lack of reactivity is that their eight valence electrons correspond to a particularly stable arrangement. Inert gases have a stable octet of valence electrons (four pairs of electrons), except for helium, which is stable with two valence electrons on its K level. The formation of chemical bonds to acheive the electron arrangement of inert gases is called the octet rule (stability with 2 electrons is called the duet rule), which states that:
       Atoms tend to gain, lose or share electrons in order to acquire a stable octet of valence electrons.

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2. Covalent and Ionic Bonding

     Just as the same peices of a Lego set can be used to construct different objects by fitting these pieces together in different ways, the 90 naturally occuring elements can associate in many different forms to yield infinite substances. A chemical bond is an attractive force that holds together atoms or ions. During the chemical bond formation, the valence electron rearrange to acheive stable octet. There are only 2 possibilities to consider. First, is a sharing of valence electrons between atoms. The second possibility is a transfer of valence electrons between atoms.

 

2.1 Covalent Bond

     The bond that holds atoms of the same or different non-metal elements together by sharing pairs of electrons is called a covalent bond. Atoms joined by covalent bonds form molecules.Depending on the number of shared pairs of electrons between two atoms, the covalent bond can be described as a single or a multiple covalent bond.

I - Single Covalent Bond:

     When two atoms can share a pair of electrons to acheive a stable state of configuration, the bond is called single covalent bond.

 

 

** the lewis dot structure is also known as the lewis dot diagram or electron dot formula.

** only valence electrons are shown in a lewis dot structure.

** note that some valence electrons are not involved in covalent bond formation. these are cakked non-bonding electron pairs.

** a lewis dot structure is a representation in which shared electron pairs are shown either as lines or as pairs of dots between two atoms and lone pairs (non-bonding pair) are shown as pairs of dots on individual atoms.

I - Multiple Covalent Bond:

     When two atoms share two or three pairs of electrons to aheive the stable octet of electrons, double and triple covalent bond form. A double covalent bond is the sharing of two pairs of electrons. It is represented by two dashes =. A triple covalent bond consists of three pairs of shared electrons; it is represented by three dashes similar to the double covalent bond.

i - Double Covalent Bond

     

i - Triple Covalent Bond

 

 

     An important issue to note in this section is the concept of polarity that is associated with electronegativity. In principle, the elements invloved in a covalent bond are sharing their electrons; however, this sharing of electrons is not done equally always. That is, some elements attract electrons more strongly than others and so the bonding pair is distorted towards one of the elements which shows greater attraction and hence has an excess of negative charge. The other element will have a deficit of negative charge or an excess of positive charge. These charges are called partial charges and designated by d- and d+. In such cases, the covalent bond is said to be polar. Consider the following water molecule:

2.2 Ionic Bond

     The transfer of electrons from one atom to another causes electrically neutral atoms to become ions. In an ionic crystal, oppositely charged ions are in direct contact; on the other hand, ions of the same charge are not in direct contact. An electrostatic attraction occurs between oppositely charged ions, and an electrostatic repulsion occurs between ions of the same charge. This electrostatic force of attraction between these oppositely charged ions is called ionic bond. Since the ions having opposite charges are in close contact, the attractive force overcomes the repulsive one. This gives the crystal a very strong cohesion which is ensured by a type of bond called ionic bond.

     In general, non-metals in the presence of metals tend to gain one, two, or three electrons from metal atoms to form negative ions, while the metal ions form positive ions (For more details, CLICK HERE to review ions). Each of the resulting ions will have stable electron arrangement of an inert gas which is closest to it in the periodic table.

2.2.1 Formation of ionic bond

     To better describe the formation of an ionic bond, let us give you an example. Here is the same compound, in 2 different views.

 

 

 

2.2.3 Mole of Ions,Formula Unit, and Formula Mass

 

1. A mole of one kind of ion is the quantity of matter which contains 6.02 x 10²³ i.e. 1 mole = 6.02 x 10²³ ions.
2. An ionic compound is represented by a formula called formula unit, which shows the proportions of each of its constituents. A formula unit DOES NOT stand for a molecule. For example, the formula unit for sodium chloride is NaCl. That of calcium chloride is CaCl₂
3. The sum of the atomic masses of all atoms present in the formula unit of an ionic compound is called the formula mass. The molar formula mass or simply molar mass of an ionic compound is the mass of a mole of this compound. For example,
   
        The formula mass of MgO ( the formula unit of magnesium oxide) is: 24 + 16 = 40 u
        whereas its molar mass is 40g.

2.2.4 Crystal Lattice

     The electrostatic force of attraction in an ionic bond is very strong. The oppositely charged ions in an ionic compound group together in a regular repeating pattern arrangement called a crystal lattice. Each ionic compund has a characteristic crystal lattice.The ionic compund, as a whole, is electrically neutral. The sum of the net charges of all the ions is zero. In general, the crystal lattices of ionic compounds can be simple cubic, body-centered cubic, or face-centered cubic.

 

Simple cubic	Body - centered cubic	Face - centered cubic

 

 

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3. Predicting Types of Bonds

     We can expect the type of bonds formed in a compund based on the electron arrangement of the atoms forming the compund.

•  Elements at the left and in the center of the periodic table are metals, which lose electrons easily. Those that are at the right of the periodic table (except elements of column 8A or 18) tend to gain electrons readily.
•  Bonds between metals and non-metals WILL be ionic. Compunds formed from metals and non-metals are ionic compounds.
•  Bonds between non-metals will be covalent. Compounds formed from non-metals are covalent molecular compounds.

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4. Some Differences Between Ionic and Molecular Compounds

     Interparticle forces make the difference in the properties of ionic and molecular compounds

 

Ionic Compounds	Covalent Compounds
solids at room temperature	not as hard as ionic compounds
have high melting points	in general, have low melting points
conduct electricity when molten or when dissolved in water	are less soluble in water than ioninc compounds and molecular solutions dont conduce electricity
are electrolytes	are non-electrolytes
made up of oppositely chrged ions	they are molecular compounds made up of molecules

 

Look at these funny animations describing the different bonds we've discussed:

 

 

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