1. Introduction to Chemical Bonding
Ch. 6 Chemical Bonding
Introduction to Chemical Bonding

2. Section 1: Introduction to Chemical Bonding
Mutual electrical attraction between the nuclei and valence electrons of different atoms that bind the atoms together.
Most atoms are at high potential energy (PE) ->not stable
By bonding most atoms decrease in PE -> become stable

3. Atoms form compounds by gaining, losing or sharing electrons
1) ionic bonding: transfer of electrons. Chemical bonding that results from the electrical attraction between cations and anions.
2) covalent bonding: results from the sharing of electron pairs between two atoms.
3) ionic or covalent? Whether the bonding is covalent or ionic can be estimated by calculating the difference in the elements’ electronegativities.

4. Fig. 3.11 – Periodic Table of Electronegativities

5. Sharing evenly, unevenly or transferring of electrons
Difference of the atoms’ electronegativities involved in the bonding determines the type of bond.
Ionic bonds: >1.7 to 3.3
Polar-covalent): 0.3 to 1.7
Nonpolar-covalent 0-0.3

6. Nonpolar-covalent & polar-covalent
Nonpolar-covalent: covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in balanced distribution of electrons: H2
Polar-covalent: uneven sharing of electrons, resulting in an uneven distribution of charge:
HCl (electron cloud is bigger around Cl than H) H H H Cl

7. Practice Problem A: Classifying Bonds (p. 167)
Classify bonding between sulfur, S, (Electronegativity of 2.5) and the following electrons:
(2.1) H: 0.4 Polar-covalent
(0.7) Cs: 1.8 Ionic
(3.0) Cl: 0.5 Polar-covalent
Practice (p. 167) Classify bonding between chlorine, Cl, (3.0) and the following elements: And, indicate the more negative atom in each pair.
(1.0) Ca: 2.0 Ionic Cl
(3.5) O: 0.5 Polar-covalent O
(2.8) Br: 0.2 Nonpolar-covalent Co

8. Section 2: Covalent Bonding and Molecular Compounds
Molecule: a neutral group of atoms that are held together by covalent bonds.
molecule is a single unit of a chemical compound that has the same properties as the compound.
Eg: O2 : molecule of oxygen (diatomic molecule: 2 atoms)
H2O: molecule of water
Chemical formula: indicates the relative #s of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts.

9. Cont.
The chemical formula of a molecular compound is referred to as a molecular formula.
Molecular formula: shows the types and #s of atoms combined in a single molecule of a molecular compound.
H2O – molecular formula
O2 – diatomic molecule (contains only 2 atoms)
Covalent bonds formed from shared electrons.
Lowest potential energy is when repulsion of nuclei and of electrons is balanced by attraction of nuclei of one atom to electrons of other atom.

10. Bond Length & Energy
Bond length: the average distance between 2 bonded atoms.
Bond energy: the energy required to break a chemical bond and form neutral isolated atoms
kJ (kilojoules: unit of energy)/mol: indicates the energy required to break one mole of bonds in isolated molecules.
Example: 436 kJ/mol is required to break the hydrogen-hydrogen bonds in one mole of hydrogen (H2) molecules and form 2 moles of separated hydrogen atoms.
1 mol H2 (436 kJ/mol) -> 1 mol H + 1 mol H

11. Atoms tend to form bonds to follow the octet rule
Octet rule: chemical compounds tend to form so that each atom, by gaining, losing or sharing electrons has an octet (8) of electrons in its highest occupied energy level.
F2:
HF:
Exceptions to the octet rule: most main group elements tend to form covalent bonds according to the octet rule. Hydrogen only needs 2 electrons to be stabel, Boron has 3 valence electrons, can form BF3, which only gives it 6 valence electrons. BF3 easily attracts another molecule that ha an unbonded pair of electrons, such as NH3. The bond formed between these molecules is called a coordinate covalent bond.)

12. Electron-dot notation
Electron-dot notation: an electron-configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element.
Valence electrons are the outer energy level electrons
of an atom. The valence electrons are the electrons
that participate in chemical bonding.

13. Lewis Dot Symbols

14. Valence electrons are the outer energy level electrons
of an atom. The valence electrons are the electrons
that participate in chemical bonding.
Group
# of valence e-
e- configuration 1A 1
ns1 2A 2
ns2 3A 3
ns2np1 4A 4
ns2np2 5A 5
ns2np3 6A 6
ns2np4 7A 7
ns2np5
9.1

15. Practice Electron-dot notations for the following atoms:
H e S
N f Cl
P g Xe Si

16. How is H2 formed?
The hydrogens atoms share their electrons to be isoelectronic with He
The sharing of 2 electrons is called a single covalent bond
Lewis Dot Structure - A symbolic description of the distribution of valence electrons in a molecule.

17. Dash structure - use dash two represent a pair of e-
A covalent bond is a chemical bond in which two or more electrons are shared by two nonmetal atoms.
Why should two atoms share electrons?
1e-
1e-
2e-
2e- H H + H
Lewis structure of H2
single covalent bond
single covalent bond H H
Dash structure - use dash two represent a pair of e-
Dot structure
9.4

18. Why should two atoms share electrons?
A covalent bond is a chemical bond in which two or more electrons are shared by two nonmetal atoms.
Why should two atoms share electrons?
7e-
7e-
8e-
8e- F F + F
Lewis structure of F2
lone pairs F
single covalent bond
single covalent bond F
9.4

19. Lone pair or Unshared pairs – pair of valence electrons that are not shared between the atoms
lone pairs F
single covalent bond
single covalent bond F
Structural formula: indicates the kind, number, arrangement, and bonds but not the unshared pairs of atoms in a molecule.
Examples:
F-F
H-Cl
9.4

20. Sample Problem C (p. 175) Draw the Lewis Structure of idomethane, CH3I
1) Determine the type and # of atoms in the molecule.
1 C
3 H
1 I
2) Write the electron-dot notation for each type of atom in the molecule.
C – 4 valence electrons
H – 1 valence electron
I – 7 valance electrons
C H I
3) Determine the total # of valence electrons.
C 1 x 4e- = 4e-
H 3 x 1 e- = 3e-
I 1 x 7e- = 7e-
14 e-

21. Continued…
4) Arrange the atoms to form a skeleton structure for the molecule. If C is present, its central. Otherwise, the least-electronegative atom is central except H, which is never central. Then connect atoms by e- pairs.

22. Continued…
5) Add unshared pair of electrons to each nonmetal atom (except hydrogen) such that each is surrounded by 8 electrons.
6) Count the electrons in the structure to be sure that the # of valence electrons used equals the # available. Be sure the central atom and other atoms besides hydrogen have an octet.

23. Practice: Draw the Lewis structure for
1) ammonia, NH3
2) hydrogen sulfide, H2S

24. Practice
3) silane, SiH4
4) phosphorus trifluoride, PF3

25. Some atoms can share multiple pairs of electrons
*atoms of some elements can share more than one electron pair, especially C, N & O
Double covalent bond or double bond: a covalent bond in which 2 pairs of electrons are shared between two atoms.
C2H4
Triple covalent bond or triple bond: a covalent bond in which 3 pairs of electrons are shared between 2 atoms.
N3 C2H2

26. Multiple bonds: double and triple bonds.
Double bonds: have more energy and shorter distances between nuclei.
Triple bonds: have even more energy and even shorter distances between nuclei.

27. Resonance structure is...
When more than one valid dot diagram is possible.
Resonance occurs when two or more equally valid electron dot structures can be written for a molecule.
Only differ in the position of the bonding electrons
Consider the two ways to draw ozone (O3)
Which one is it?
It is a hybrid of both, like a mule; shown by a double-headed arrow
Write the resonance structure for CO32-

28. Resonance Structures for CO32-
Carbon and oxygen bond:
1 single bond and 1/3 double bond