Covalent Compounds – Basics

Covalent bonds were presented In the Intro to Covalent bonding in the Elements section as the sharing of valence electrons between nonmetal atoms.  Compounds H2, O2, N2, F2, Cl2, Br2, and I2 discussed in that section are covalent compounds.  These compounds will be used to expand the discussion regarding covalent compounds.

In the Intro to Covalent bonding and earlier in the Compound section, the statement was made that the more valence electrons in the valence (outer) energy level, the greater the stability of that atom.  The most stable naturally occurring atoms are the noble gas atoms (Family 18) and with the exception of He, they all have 8 valence electrons.  This gives rise to the “rule of octet” which says all atoms want 8 electrons in their valence energy level.  As the case with all rules, there are exceptions and you will encounter them.  But for the majority of covalent compounds, this rule is followed.

Before continuing there are several concepts and tools that need to be introduced.

A molecule is the basic entity of a covalent compound   See MoleculeAnswers : molecule, and Marvelous Molecules.

Lewis structures are used to help visualize the arrangement of atoms and valence electrons in covalent compounds   Check the following links to see what Lewis structures are and how to use them, Lewis Structure, Lewis Diagrams and Structures, A Sure-Fire Way to Draw Lewis structures.  Select item H. Covalent and Lewis Structures.  Then select “Simple Lewis structures” and follow the discussion.

Electronegativity is a calculated value that indicates the ability of an atom to attract electrons towards itself in a covalent bond.  Smaller atoms are more electronegative than larger atoms.  Using your current Internet browser, open a second window.  In Chrome, Firefox, and Safari, under the “File” drop down menu there is a ”New Window” choice.  Click this to open a new window.  Copy and paste the following address: http://www.ptable.com/#   When this site is available, click on “Properties” tab.  Then click on “electronegativity” in the list of properties listed above the periodic table.  Note the most electronegative atom in a covalent compound is fluorine (F) with a value of 3.98.  All other elements have lower electronegative values.  Open a third window with your browser.   Copy and paste the URL http://www.crystalmaker.com/support/tutorials/crystalmaker/atomicradii/index.html . Compare the sizes of atoms in this graphic to the electronegative values in the Ptable window.  The student should observe and remember the electronegative trend in the periodic table.   Electronegative values increase as one moves right in a Period and up in a Family.

Molecular shape, electron arrangement, and symmetry are important in determining the chemical and physical properties of molecules as will be discussed below.

Examples
To show how the above concepts and tools are applied to covalent compounds, we will start with the diatomic compounds H2, Cl2 O2 and N2 and proceed to compounds with more complex molecules.  You should still have the PTable open (see above instructions) so you can see the number of valence electrons.

H2
Hydrogen is an exception to the octet rule because it is a much smaller atom.  Each hydrogen atom has one valence electron.  The Lewis structure for a hydrogen atom is

H lewis

The H in the symbol represents the nucleus of the hydrogen atom and the dot represents the valence electron. (For atoms of all other elements, the symbol for the element represents the nucleus and all lower energy level electrons.  The dots represent the valence electrons.)

When two hydrogen atoms share the two electrons, the Lewis structure becomes:

H2 lewis

Each dot represents a valence electron from each of the hydrogen atoms.  A single line is also used to represent two electrons shared between two atoms.  The bond is said to be a “single bond” because two electrons are shared by the two atoms.  Since there are only two atoms, the structure of a hydrogen molecule is linear. The two atoms are identical so the electronegativity is the same in each atom and the electrons spend equal time around each hydrogen atom.  H2 molecules are symmetrical with respect to both the placement of the charge and the atoms.  This molecule is said to have “non-polar covalent bonding” and is a “non-polar covalent molecule”.

Cl2
Chlorine atoms have 7 valence electrons.  The Lewis structure for a Cl atom is:

Cl atom Lewis

The Cl symbol represents the nucleus and all lower level energy electrons.  The 7 dots represent the 7 valence electrons.

When two chlorine atoms share two valence electrons, the Lewis structure becomes:

Cl2 molec Lewis

There are two electrons shared by the two Cl atoms so there is a single bond between the Cl atoms.  Each chlorine has 7 valence electrons plus the one it is sharing for a total of 8 valence electrons.  So each Cl atom will have 8 valence electrons as a result of sharing.  This is the rule of octet.

Since there are only two atoms, the structure of a Cl2 molecule is linear. The two atoms are identical so the electronegativity is the same in each atom and the electrons spend equal time around each chlorine atom.  Molecules of Cl2 are symmetrical with respect to both the placement of the atoms and the charge.  This molecule is said to have non-polar covalent bonding and is a non-polar covalent molecule.

O2
Oxygen atoms have 6 valence electrons.  The Lewis structure for an O atom is:

Oxyg Lewis

 

When two oxygen atoms share valence electrons, the Lewis structure becomes:

O2 lewis

There are four electrons (2 pairs) shared by the two O atoms so there is a “double bond” between the O atoms.  Each oxygen has 6 valence electrons plus the ones it is sharing for a total of 8 valence electrons.  So each O atom will have 8 valence electrons as a result of sharing.  This is the rule of octet.

Since there are only two atoms, the structure of a O2 molecule is linear. The two atoms are identical so the electronegativity is the same in each atom and the electrons spend equal time around each chlorine atom.  Molecules of O2 are symmetrical with respect to both the placement of the atoms and the charge.  This molecule is said to have non-polar covalent bonding and is a non-polar covalent molecule.

N2
Nitrogen atoms have 5 valence electrons.  The Lewis structure for an N atom is:

N lewis

When two nitrogen atoms share valence electrons, the Lewis structure becomes:

N2 lewis

There are six electrons (3 pairs) shared by the two N atoms so there is a “triple bond” between the N atoms.  Each nitrogen has 5 valence electrons plus the ones it is sharing for a total of 8 valence electrons.  So each N atom will have 8 valence electrons as a result of sharing.  This is the rule of octet.

Since there are only two atoms, the structure of a N2 molecule is linear. The two atoms are identical so the electronegativity is the same in each atom and the electrons spend equal time around each chlorine atom.  Molecules of N2 are symmetrical with respect to both the placement of the atoms and the charge.  This molecule is said to have non-polar covalent bonding and is a non-polar covalent molecule.

The preceding covalent bond examples show how atoms can share electrons to achieve the rule of octet (each atoms wants 8 electrons).  Examples were also used to introduced the existence of single bonds (1 pair of electrons shared between two atoms), double bonds (2 pair of electrons shared between two atoms), and triple bonds (3 pair of electrons shared between two atoms).  The student will encounter each of the above during this course.


HCl (hydrogen chloride)

The H atom has one valence electron and the Cl atom has 7 valence electrons.  The Lewis structure for HCL is:

HCl lewis

H has 2 valence electrons as a result of sharing with Cl and the CL has 8 valence electrons as a result of sharing. Since there are only two atoms, the structure of a HCl molecule is linear.  The electronegativity of H is 2.20 and the electronegativity of Cl is 3.16.   As was mentioned above electrons will spend more time around more electronegative atoms.  Therefore electrons will spend more time around the Cl atom and less time around the H atom.  This results in unsymmetrical placement of charge with the H atom being more positive charged and the Cl atom being more negative charged.  This separation of charges produces a “dipole” (two poles of charge) within the HCl molecule. The HCl molecule is not symmetrical with the placement of either the atoms or charges.  This molecule is said to have “polar covalent bonding” and is a “polar covalent molecule’.  The result of polar covalent bonding can be observed in the table below comparing boiling points. The trend in boiling points for H2, N2, O2, and Cl2 was introduced and explained in the Introduction to Covalent Bonds.

COMPOUND BOILING PT(C) MOLECULAR WEIGHT (G/MOLE)
H2 -295 2
N2 -195.8 28
O2 -183 32
HCL -85 36.5
O2 -34 71

The difference between H2, N2, O2, Cl2 and HCl molecules is that HCl molecules have a permanent dipole.  The positive end of HCl molecules will be attracted to the negative ends of neighboring HCl molecules.  The force of this “intermolecular attraction” requires more energy to separate them and that is reflected in a higher boiling point.

Why would Cl2 a nonpolar compound have a higher boiling point than HCl a polar compound?  The difference is due to Cl2 molecules being larger (approximately 2 times) than HCl molecules. The larger size allows the formation of instantaneous dipoles (scroll to bottom of page ) and a small attraction between neighboring molecules.  In comparison, O2 and HCl are closer in size but there is a large difference in the boiling points due to the dipoles in the HCl molecules.

The foregoing discussion has focused on two atom molecules to illustrate basic principles that will be applied to all molecules.  The discussion now will be expanded to include compounds with larger molecules.