Covalent bonds are either polar or nonpolar.
Nonpolar bonds occur when electrons are equally shared, causing the electron distribution between the two atoms to be symmetrical. An example is H - H. Bonds formed between atoms that differ in electronegativity by 0.4 or less behave like pure nonpolar bonds and are considered nonpolar.
Ionic bonds occur when the electronegativity difference between atoms is 1.7 or greater with the exception of HF. We will not consider these because the question of polarity does not apply.
Polar bonds occur when the electrons forming the bond are not equally shared, causing the electrons to be pulled toward the atom with the higher electronegativity. The electron distribution between the two atoms of the bond is unsymmetrical.
For a molecule to be polar, it must have polar bonds AND be nonsymmetrical. After determining the molecule has polar bonds, then to determine if the molecule is polar, look for pairs of electrons around the central atom that are not shared. These pairs of electrons that are not shared cause the nonsymmetrical shape of the molecule. These molecules are polar.
If the molecule has polar bonds but is symetrical then it is nonpolar. If the molecule does not have polar bonds then it is obvious that the molecule is nonpolar.
The octet rule states that all covalently bonded atoms will tend to surround themselves with four pairs of electrons, or eight electrons. The four pairs correspond to the number of electron pairs found in the valence shells of the most stable elements, the Noble Gases.
There are exceptions to the Octet Rule such as aluminum and boron atoms needing only six electrons in the structure and hydrogen which tries to duplicate helium by surrounding itself with one pair of electrons.
To be able to predict the molecular shapes of a molecule, one must use the VSEPR theory. This theory states that the electron pairs of different valence orbitals repel each other and cause the orbitals to arrange themselves so that they are as far apart as possible. This minimizes the repulsive forces between the electron pairs. To use VSEPR you must draw the Lewis (Electron Dot) Structure for the molecule.
1. Draw the Lewis Structure of the molecule.
2. Count the number of electron pairs around the central atom. Double or triple bonds count as a single bond ( one electron pair).
3. The arrangement of electron pairs that minimize repulsion is called the electron pair geometry.
4. The arrangement of atoms is called molecular geometry.
For the A group elements the number of valence electrons is the same as the group number in the periodic table, with the exception of helium.
1. The first step is to determine the number of electrons to be used to connect the atoms. This is done by adding up the number of valence electrons of the atoms in the molecule. If there is a negative charge add one electron to the total number of valence electrons for every negative charge. If there is a positive charge indicated on the molecular formula subtract one electron from the total number of valence electrons for every positive charge indicated.
2. Determine the central atoms and arrange the other atoms as symmetrical as possible around the central atom (the skeleton structure).
Connect the central atom to the other atoms in the molecule with single bonds. Each bond uses 2 valence electrons.
3. Complete the valence shell of the outer atoms in the molecule. If you run out of electrons at this point, the skeleton structure is wrong. Go back to step 2.
4. Place any remaining electrons on the central atom. Check each atom to see if it obeys the Octet Rule which says that an atom in order to be most stable must have eight valence electrons taking into account the exceptions. If the valence shell of the central atom is complete you have drawn an acceptable Lewis structure. If the valence shell of the central atom is not complete, use a lone pair on one of the outer atoms to form a double bond between that outer atom and the central atom. Continue making multiple bonds until the valence shell of the central atom is complete.
5. Check to make sure you have used to correct number of electrons.
There are three electron geometries possible:tetrahedral, trigonal planar, and linear.
Count the number of electron pairs, BOTH shared and unshared that are around the central atom. Four electron pairs gives an electron geometry of tetrahedral, three electron pairs gives the electron geometry of trigonal planar, and two electron pairs gives the electron geometry of linear.
There are five molecular shapes a molecule can have: tetrahedral, trigonal planar, trigonal pyramidal, bent, and linear. To determine the shape you must first count the number of electron pairs around the central atom that are shared. Ignore all unshared electron pairs. Molecular Geometry considers ONLY the bonding electrons.
A molecule that has a tetrahedral shape has all four pairs of electrons bonded. The four pairs of electrons are repelling each other and to minimize the repulsion the orbitals get as far apart as possible. This causes the tetrahedral shape. Because all four pairs of electrons are bonding pairs, the bond angles are all 109.5o. An example of a tetrahedral shaped molecule is methane.
A molecule with trigonal planar shape has three bonds all of which lie in the same plane. This can include one double bond which acts as a single bond. The bond angles are 120o. An example of trigonal planar molecule is boron trifluoride.
A molecule with trigonal pyramidal shape has four pairs of electrons all repelling each other. Of the four pairs of electrons, 3 are bonding pairs, and one lone pair of electrons. Bond angles for trigonal pyramidal are 109.5o. Ammonia is an example of a molecule with trigonal pyramidal shape.
Molecules with a bent shape has four pairs of electrons, but only two pairs are bonding pairs (two are lone pairs). Because only two pairs are bonding the bond angle is 109.5o. Water is an example of a molecule with a bent shape.
A molecule with a linear shape has two pairs of bonding electrons. Repulsions between these two pairs of electrons will cause them to be as far apart as possible which creates the linear shape. Bond angles for linear molecules are 180o. Carbon disulfide is an example of a molecule with a linear shape (double bonds act as single electron pairs according to VSEPR).
|Total No. of electron pairs||Shared||Unshared||Bond Angle||Electron Pair Geometry||Molecular Geometry||Example|
|3||3||0||120||trigonal planar||trigonal planar||BF3|
Charged molecules have extra electrons that have been added. To draw a lewis structure of these molecules, put brackets around the lewis structure, and include the charge on the outside of the brackets.
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last updated: February 5, 2001