1. Chemical Bonding and Interactions
The Ties That Bind
Chemical Bonding and Interactions

2. Chemical Bonding and Interactions
Stable Electron Configurations
Electron-Dot (Lewis) Structures
Drawing, Rules for Drawing
The Octet Rule
Some Exceptions to the Rule
Ionic Bonding
Naming ionic compounds
Drawing
Covalent Bonding
Naming covalent compounds
Electronegativity and Polar Covalent Compounds
Molecular Shapes and the VSEPR Theory
Intermolecular Forces of Attraction
H-bonds, Dipole-Dipole, Ion-Dipole, London Dispersion Forces

3. It’s all about stability…
Fact: Noble gases are inert (undergo few, if any, chemical reactions)
Theory: The inertness of these gases is a result of their electronic structure
If elements could alter their electron structures like those of the noble gases, they would be less reactive.

4. Electron Configurations of Cations and Anions
Of Representative Elements
Na [Ne]3s1
Na+ [Ne]
Atoms lose electrons so that cation has a noble-gas outer electron configuration.
Ca [Ar]4s2
Ca2+ [Ar]
Al [Ne]3s23p1
Al3+ [Ne]
H 1s1
H- 1s2 or [He]
Atoms gain electrons so that anion has a noble-gas outer electron configuration.
F 1s22s22p5
F- 1s22s22p6 or [Ne]
O 1s22s22p4
O2- 1s22s22p6 or [Ne]
N 1s22s22p3
N3- 1s22s22p6 or [Ne]

5. Cations and Anions Of Representative Elements+1 +2 +3 -3 -2 -1

6. Na+: [Ne]
Al3+: [Ne]
F-: 1s22s22p6 or [Ne]
O2-: 1s22s22p6 or [Ne]
N3-: 1s22s22p6 or [Ne]
Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne
What neutral atom is isoelectronic with H- ?
H-: 1s2
same electron configuration as He

7. Valence electrons are the outer shell 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

8. (Electron Dot Structures)
*We generally place the electrons one four sides of a square around the element symbol.
*Octet rule: we know that s2p6 is a noble gas configuration. We assume that an atom is stable when surrounded by 8 electrons (4 electron pairs).

9. A violent bond: Sodium and Chlorine
Sodium is a soft, reactive metal.
Chlorine is a greenish-yellow gas which is toxic (WWI poison)
Sodium chloride is used for our food.

10. Na(s) + ½Cl2(g)  NaCl(s)
Ionic Bonding
Consider the reaction between sodium and chlorine:
Na(s) + ½Cl2(g)  NaCl(s)

11. Solar Evaporation for Salt
Chemistry In Action:
Sodium Chloride
Mining Salt
Solar Evaporation for Salt

12. The Ionic Bond - - - Li+ F Li + F 1s22s1 1s22s22p5 [He] 1s2 1s22s22p6
[Ne] Li
Li+ + e-
e- + F - F -
Li+ +
Li+
IONIC BONDING
-bonding due to electrostatic attraction arising from an exchange of electrons.
9.2

13. The ionic compound NaCl
ionic compounds consist of a combination of cations and an anions
the formula is always the same as the empirical formula
the sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero
The ionic compound NaCl
2.6

14. Ionic Bonding
*NaCl forms a very regular structure in which each Na+ ion is surrounded by
6 Cl- ions.
Similarly, each Cl- ion is surrounded by six Na+ ions.
There is a regular arrangement of Na+ and Cl- in 3D.
Note that the ions are packed as closely as possible.
Note that it is not easy to find a molecular formula to describe the ionic lattice.

15. SIDEBAR: Electrostatic (Lattice) Energy
Lattice energy (E) is the energy required to completely separate one mole of a solid ionic compound into gaseous ions.
Q+ is the charge on the cation
E = k
Q+Q- r
Q- is the charge on the anion
r is the distance between the ions
cmpd
lattice energy
MgF2
MgO
2957
3938
Q= +2,-1
Q= +2,-2
Lattice energy (E) increases as Q increases and/or
as r decreases.
LiF
LiCl
1036
853
r F < r Cl
9.3

16. 9.3

17. cation – ion with a positive charge
An ion is an atom, or group of atoms, that has a net positive or negative charge.
cation – ion with a positive charge
If a neutral atom loses one or more electrons
it becomes a cation. Na
11 protons
11 electrons
Na+
11 protons
10 electrons
anion – ion with a negative charge
If a neutral atom gains one or more electrons
it becomes an anion.
Cl-
17 protons
18 electrons Cl
17 protons
17 electrons
2.5

18. A monatomic ion contains only one atom
Na+, Cl-, Ca2+, O2-, Al3+, N3-
A polyatomic ion contains more than one atom
OH-, CN-, NH4+, NO3-
2.5

19. Test Yourself: MgO MgBr2
The combination of the ion of magnesium and the ion of nitrogen

20. Formula of Ionic Compounds
2 x +3 = +6
3 x -2 = -6
Al2O3
Al3+
O2-
1 x +2 = +2
2 x -1 = -2
CaBr2
Ca2+
Br-
1 x +2 = +2
1 x -2 = -2
Na2CO3
Na+
CO32-
2.6

21. 2.6

22. 2.7

23. Chemical Nomenclature
Ionic Compounds
often a metal + nonmetal
anion (nonmetal), add “ide” to element name
BaCl2
barium chloride
K2O
potassium oxide
Mg(OH)2
magnesium hydroxide
KNO3
potassium nitrate
2.7

24. Transition metal ionic compounds
indicate charge on metal with Roman numerals
FeCl2
iron(II) chloride
2 Cl- -2 so Fe is +2
FeCl3
3 Cl- -3 so Fe is +3
iron(III) chloride
Cr2S3
3 S-2 -6 so Cr is +3 (6/2)
chromium(III) sulfide
2.7

25. 2.7

26. Covalent Bonding

27. Why should two atoms share electrons?
A covalent bond is a chemical bond in which two or more electrons are shared by two 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

28. Lewis structure of water
single covalent bonds
2e-
8e-
2e- H + O + H O H or
Double bond – two atoms share two pairs of electrons
8e-
8e-
8e-
double bonds O C or O C
double bonds
Triple bond – two atoms share three pairs of electrons
8e-
triple bond N
8e- or N
triple bond
9.4

29. Lengths of Covalent Bonds
Bond Type
Bond Length
(pm)
C-C
154
CC
133
CC
120
C-N
143
CN
138
CN
116
Bond Lengths
Triple bond < Double Bond < Single Bond
9.4

30. 9.4

31. Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms H F
electron rich
region
electron poor
region
e- poor
e- rich F H d+ d-
9.5

32. Electron Affinity - measurable, Cl is highest
Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond.
Electron Affinity - measurable, Cl is highest
X (g) + e X-(g)
Electronegativity - relative, F is highest
9.5

33. 9.5

34. Bond Polarity and Electronegativity

35. Drawing Lewis Structures
1. Add the valence electrons.
2. Identify the central atom (usually the one with the highest molecular mass, lowest electronegativity, or closest to the center of the periodic table).
3. Place the central atom in the center of the molecule and add all other atoms around it.
4.Place one bond (two electrons) between each pair of atoms.
5.Complete the octet for the central atom.
6.Complete the octets for all other atoms. Use double bonds if necessary.
7. Place remaining electrons on the central atom.

36. 5 + (3 x 7) = 26 valence electrons
Write the Lewis structure of nitrogen trifluoride (NF3).
Step 1 – N is less electronegative than F, put N in center
Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)
5 + (3 x 7) = 26 valence electrons
Step 3 – Draw single bonds between N and F atoms and complete
octets on N and F atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons F N
9.6

37. Trial
Write the Lewis stucture for
H2O
CO2
NH3
CH4 N2
BH3

38. Exceptions to the Octet Rule
The Incomplete Octet
Be – 2e-
2H – 2x1e-
4e-
BeH2 H Be
B – 3e-
3F – 3x7e-
24e-
3 single bonds (3x2) = 6
9 lone pairs (9x2) = 18
Total = 24 F B
BF3
9.9

39. Exceptions to the Octet Rule
Odd-Electron Molecules
N – 5e-
O – 6e-
11e- NO N O
The Expanded Octet (central atom with principal quantum number n > 2) S F
S – 6e-
6F – 42e-
48e-
6 single bonds (6x2) = 12
18 lone pairs (18x2) = 36
Total = 48
SF6
9.9

40. Strengths of Covalent Bonds
Bond Enthalpies and Bond Length

41. Classification of bonds by difference in electronegativity
Bond Type
Covalent
 2
Ionic
0 < and <2
Polar Covalent
Increasing difference in electronegativity
Covalent
share e-
Polar Covalent
partial transfer of e-
Ionic
transfer e-
9.5

42. Classify the following bonds as ionic, polar covalent,
or covalent: The bond in CsCl; the bond in H2S; and
the NN bond in H2NNH2.
Cs – 0.7
Cl – 3.0
3.0 – 0.7 = 2.3
Ionic
H – 2.1
S – 2.5
2.5 – 2.1 = 0.4
Polar Covalent
N – 3.0
N – 3.0
3.0 – 3.0 = 0
Covalent
9.5

43. Test Yourself Write the formula and draw the ionic structure
Magnesium (II) bromide
Iron (III) Oxide
Sodium Azide
Ammonium Chloride
Magnesium Nitrate
Ammonium Phosphate
Write the name
FeO
Fe2O3
LiBr
KCrO4
K2Cr2O7