1. Chemical Bonding I: Basic Concepts
Chapter 8
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2. Lewis Dot Symbols
Consists of the symbol of an element and one dot for each
valence electron in an atom of the element.
Note that (except helium) the number of valance electrons
each atom has is the same as the group number of the
element.

3. Ionic Bond
Is the electrostatic force that holds ions together in an ionic compound
The metal gives the electrons to the non metal
Cation is the metal
Anion is the non metal

5. The Covalent Bond
A bond in which two electrons are shared by two atoms.
Only in covalent compounds
Non metal and non metal

6. The Covalent Bond

7. The Covalent Bond
Pairs of valence electrons that are not involved in covalent bond formation are called: Lone pairs
We can only draw Lewis structures for compounds that have covalent bonds

8. Polar Covalent Bonds
Though atoms often form compounds by sharing electrons, the electrons are not always shared equally.
Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does.
Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.

9. Electronegativity
The ability of an atom to attract toward itself the electrons in a chemical bond

10. Electronegativity
The greater the difference in electronegativity, the more polar is the bond

11. Lewis Structures
Is a representation of covalent bonding in which shared electron pairs are shown either as lines or as pairs of dots between two atoms, and lone pairs of dots on individual atoms
Only valence electrons are shown in a Lewis structure

12. Writing Lewis Structures
Find the sum of valence electrons of all atoms in the polyatomic ion or molecule.
If it is an anion, add one electron for each negative charge.
If it is a cation, subtract one electron for each positive charge.
PCl3
Keep track of the electrons:
(7) = 26

13. Writing Lewis Structures
The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds.
Keep track of the electrons:
26 − 6 = 20

14. Writing Lewis Structures
Fill the octets of the outer atoms.
Keep track of the electrons:
26 − 6 = 20; 20 − 18 = 2

15. Writing Lewis Structures
Fill the octet of the central atom.
Keep track of the electrons:
26 − 6 = 20; 20 − 18 = 2; 2 − 2 = 0

16. Writing Lewis Structures
If you run out of electrons before the central atom has an octet…
…form multiple bonds until it does.

17. The "best" Lewis structure for NO3-
1.  Determine the total number of valence electrons in a molecule
Draw a skeleton

18. Of the 24 valence electrons in NO3-, 6 were required to make the skeleton. Consider the remaining 18 electrons and place them so as to fill the octets of as many atoms as possible

19. Are the octets of all the atoms filled
Are the octets of all the atoms filled?   If not then fill the remaining octets by making multiple bonds
Check that you have the lowest FORMAL CHARGES possible for all the atoms, without violating the octet rule; (valence e-) - (1/2 bonding e-) - (lone electrons).

20. Writing Lewis Structures
The best Lewis structure…
…is the one with the fewest charges.
…puts a negative charge on the most electronegative atom.

21. The Concept of Resonance
Resonance structure, is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure
Resonance means the use of two or more Lewis structures to represent a particular molecule

22. Exceptions to the Octet Rule
The expanded octet
Atoms in the 2nd period (that are in families 3A-7A) cannot have more than 8 valance electrons around the central atom
Atoms of elements in and beyond the 3rd period (that are in families 3A-7A) form some compound in which more than 8 electron surround the central atom
Odd- electron molecules
Some molecules contain an odd number of electrons.
We need an even number of electron for complete pairing the octet rule clearly cannot be satisfied with all the atoms in any of these molecules
Incomplete octet
The number of electrons surrounding the central atom in a stable molecule is fewer than eight

23. The modern periodic table.
Figure 2.10
The modern periodic table. +1 +2 +3 NC -3 -2 -1

24. Bond Enthalpy A measure of the stability of a molecule
Thus bond enthalpy is the enthalpy change required to break a particular bond in 1 mole of gaseous molecules
Ho(reaction) = sum of the bond energies of bonds being broken - sum of the bond energies of the bonds being formed.
Ho(reaction) = H(reactant bonds broken) - H(product bonds formed)

25. Example of Using Bond Energies to Calculate Heat (enthalpy) of Reaction
Use the bond energies provided in the table above to calculate the heat (enthalpy) of reaction, Ho, for the reaction: CH4(g) + 4Cl2(g) -----> CCl4(g) + 4HCl(g) Write the balanced chemical equation, with all reactants and products in the gaseous state. CH4(g) + 4Cl2(g) -----> CCl4(g) + 4HCl(g) Write the general equation for the heat (enthalpy) of reaction: Ho(reaction) = H(reactant bonds broken) - H(product bonds formed)

26. Bonds Broken Bonds Formed bond type bond energy bond type bond energy
Substitute bond energy values into the equation and solve for Ho(reaction)
Bonds Broken Bonds Formed
bond type bond energy bond type bond energy
4 x C – H 4 x 413= x C - Cl 4 x 328 = 1312
4 x Cl – Cl 4 x 243 = x H - Cl x 432 = 1728
H(reactant bonds broken) = H(product bonds formed) = 3040
Ho(reaction) = = -416 kJ mol-1