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Lewis Dot Structure of Hydrogen Fluoride.
Drawing Lewis Structures Sum the valence electrons from all atoms in the species. Write the atomic symbols for the atoms involved so as to show which atoms are connected to which. Draw a single bond between each pair of bonded atoms Complete the octets of the atoms bonded to the central atom (i.e. the peripheral atoms) Place leftover electrons on the central atom, even if it results in the central atom having more than an octet If there are not enough electrons to give the central atom an octet, form multiple bonds by pulling terminal electrons from a peripheral atom and placing them into the bond with the central atom
Question:
1. Draw the Lewis structure for ammonia, NH 3. Solution: Since each H can form only one covalent bond, the arrangement of atoms must be:
From the periodic table we see that N has five valence electrons. These, plus one electron from each H, give us a total of eight. Bonding the atoms in the molecule requires the use of six valence electrons, as
The remaining two valence electrons are then assigned to N to complete its octet:
2. Write the Lewis structure for ethane, C 2 H 6. Solution: A little thought will reveal that the two C atoms must be bonded to each other. (Remember: Hydrogen forms only one bond.) Keeping in mind the octet rule, we predict that besides bonding, to the other C atom, each C forms three bonds to H atoms:
Each C contributes four valence electrons. Bonding all atoms in the molecule with pairs of electrons uses all of these. The Lewis structure is thus
3. Write the Lewis structure for ethylene, C 2 H 4 Solution: Here, as in the previous example, the two C atoms must be bonded to each other.
(b) how the valence electrons are assigned in the molecule, that is, which pairs are bonding pairs and which are lone pairs.
4. Write the Lewis structure for the ammonium ion NH 4 + Solution: The arrangement of the atoms is
The total number of valence electrons is (5 + 4) — 1 = 8. (Here we have subtracted one electron from the total provided by one N and four H atoms, because the ion has a positive charge and therefore has one less electron than a (neutral) molecule. The Lewis structure is
In the electron counting process one electron of a bonding pair often appears to have come from one of the two bonded atoms, and the other electron from the other atom. Thus we have for the formation of the three covalent bonds in NH 3.
or
To keep track of electrons origins small x’s and o’s are sometimes used in Lewis structures. Writing x's for the electrons from the H atoms and o's for those from N, we get for NH 3
A covalent bond in which each electron of the pair appears to have come from each bonded atom is called a normal covalent bond. In the ammonium ion, again using the x and o symbolism, we can write
in this Lewis structure one of the H atoms appears to be bonded to the N by a pair of electrons, both of which originated with the N atom. Such a bond is called a coordinate covalent bond, or sometimes, a dative bond. But all four N—H bonds in NH 4 +^ are identical in all measurable properties. So a coordinate covalent bond is in no way different from a normal covalent bond. There appears to be a difference only when we keep track of the electrons' origins.
5. Draw the Lewis structure for PCl 3. Answer:
bonded to the central atom. BUT : There are only 6 valence electrons left. If we put them on N we do not achieve an octet at C !!
The octet rule and Lewis structures Whatever possible, Lewis structures should show the octet rule to be obeyed. There are some molecules, however, in which the octet rule is clearly violated. Consider the molecule of phosphorus pentachloride, PC1 5. In this molecule phosphorus atom is bonded covalently to five chlorine atoms. The total number of valence electrons is 40 (5 from the P plus 35 from the five Cl atoms). Since the P forms five bonds, the Lewis structure is
Here the valence shell of the phosphorus atom is said to have been expanded in order to accommodate five electron pairs. The expansion of the valence shell of an atom is possible only if the atom has nd or {n — 1} d orbitals which can be added to the ns and three np orbitals normally constituting its valence shell. In the case of PCL 5 the 10 bonding electrons are accommodated in the valence shell of phosphorus which has been expanded by the addition of one of phosphorus 3d orbitals. The valence shells of
Step 4: Try using multiple bonding to share the electrons between C and N. A triple bond is required to give an octet to each atom
atoms of periods 1 and 2 cannot be expanded because they contain no \d or 2d orbitals. (The 3d orbital is unavailable for these atoms because it is of such high energy.) Sometimes the valence shell of an atom in a molecule contains less than an octet. This is the case with boron trifluoride, BP 3. Its Lewis structure is written as
Here the valence shell of boron holds only three pairs of electrons, and again the octet rule is violated. The octet rule is a handy generalization, but exceptions to it are numerous. It must be violated in molecules having an odd number of valence electrons. The Lewis structure for nitric oxide, NO, can be shown as
Bond polarity Identical atoms have identical electronegativities. In the H molecule
the hydrogen atoms attract the electron pair equally. The electronic charge distribution is symmetrical with respect to the two nuclei; that is, it is not pulled closer to one atom than the other. Since one end of the bond is electrostatically just like the other, the bond is said to be nonpolar. (This just means that it does not have different poles, or ends.) For the same reason the bond in
nonpolas by writing two dots exactly in the middle between the symbols A and B. A : B (nonpolar covalent bond) Now we turn the dial on the remote controller and gradually decrease the electronegativity of A. We make A less electronegative, or more electropositive. This decrease is pulling on the pair of electrons, so that their average position moves closer to B creating partial charges, +and -, on A and B, respectively: A+: B-^ (polar covalent bond) The bond has now become polar. It becomes increasingly polar as we further decrease the electronegativity of A, until finally the probability of finding the electron pair on A is very low and on B. very high. The electron pair now largely "belongs" to B. This gives B a net negative charge and leaves A with a positive charge, for A has now transferred an electron to B: A+^ [:B]-^ (ionic bond) The ionic bond can be seen to be an extremely polar bond, one in which there is essentially no sharing of electrons.
Pauling electronegativities : The concept of electronegativity was originally proposed in 1932 by the American chemist Linus Pauling. He pointed out that the distribution of the electronic charge cloud of a bonding pair of electrons should be related to the strength of the bond. A bond which is highly polar (has a high
degree of ionic character) should be very strong, as the attraction between the partial negative charge built up on one
Linus Pauling
atom and the partial positive charge left on the other should augment the bond strength.Using measured values of bond energies (bond energy is the energy necessary to break a bond), Pauling devised a set of electronegativity values for most of the elements.
Numerical values of electronegativities are useful for estimating the polarity or degree of ionic character of a bond. The dividing line between predominately ionic and predominately covalent character works out to be an electronegativity difference of about 1.7. This fact is sometimes useful for deciding whether to write a covalent Lewis structure or an ionic one. For HCl, for example, we
a group (as radius and the number of inner shells both increase). Thus we find that the most electronegative atom is fluorine (F), at the upper right of the periodic table, and the least electronegative (or most electropositive) is firancium (Fr), at the lower left. One of the chemical characteristics of a typical metal is a low eleclronegativity. Thus we find the best metals on the left and the best nonmetals on the right in the periodic table. Note that the transition elements all have fairly low electronegativities; they are metals. The change from metallic to nonmetallic properties occurs to the right of these elements in the periodic table. Lastly, note that it is the metals (at the left) which tend to form positive, simple (monatomic) ions, and the nonmetals (at the right) which tend to form negative ones. (Develop the habit of mentally equating metallic with electropositive and nonmetallic with electronegative.)
