1. Chapter 11 Chemical Bonding
Forces that hold atoms together

2. The Nature of Bonding
There are several major types of bonds. Ionic, covalent and metallic bonds are the three most common types of bonds.
Covalent bonds – electrons are shared between atoms.
Ionic bonds – electrons are transferred between atoms, creating cations and anions.
Metallic bonds – two or more metals bonded together.

3. The Nature of Covalent Bonding
There are two different types of covalent bonds, polar covalent and nonpolar covalent.
polar covalent – electrons are not shared equally between the two bonded atoms. The electrons are pulled toward the more electronegative of the elements.
nonpolar covalent – electrons are shared equally between the two bonded atoms.

4. Electronegativities 2

5. The formation of a bond between two hydrogen atoms.
Source: Andrey K. Geim/High Field Magnet Laboratory/University of Nijmegen

6. Probability representations of the electron sharing in HF
Probability representations of the electron sharing in HF. (a) What the probability map would look like if the two electrons in the H–F bond were shared equally. (b) The actual situation, where the shared pair spends more time close to the fluorine atom than to the hydrogen atom.

7. The Nature of Covalent Bonding
Ionic bonds are formed when there is an electronegativity difference (DEN) greater than 2.0.
Polar covalent bonds form when there is a DEN between 0.5 and 1.7.
Nonpolar covalent bonds form when there is a DEN between 0 and 0.49.

9. The Nature of Covalent Bonding
If the DEN is between 1.7 and 2.0, an ionic bond will form if a metal is one of the elements, and a polar covalent bond will form if only nonmetals or metalloids are present.

10. The Nature of Covalent Bonding
What type of bond is formed between the following elements?
N and O K and F
Mg and Cl P and F
C and H

11. Bond Polarity
Covalent bonding between unlike atoms results in unequal sharing of the electrons
One end of the bond has larger electron density than the other
The result is bond polarity
The end with the larger electron density gets a partial negative charge
The end that is electron deficient gets a partial positive charge
H F •   6

12. The three possible types of bonds: (a) a covalent bond formed between identical atoms; (b) a polar covalent bond, with both ionic and covalent components; and (c) an ionic bond, with no electron sharing.

13. Dipole Moment
Bond polarity results in an unequal electron distribution, resulting in areas of partial positive and partial negative charge
Any molecule that has a center of positive charge and a center of negative charge in different points is said to have a dipole moment (two different poles of charge). 8

14. Dipole Moment
If a molecule has more than one polar covalent bond, the areas of partial negative and positive charge for each bond will partially add to or cancel out each other
The end result will be a molecule with one center of positive charge and one center of negative charge
The dipole moment effects the attractive forces between molecules and therefore the physical properties of the substance

15. (a) The charge distribution in the water molecule
(a) The charge distribution in the water molecule (b) The water molecule behaves as if it had a positive end and a negative end, as indicated by the arrow.

16. (a) Polar water molecules are strongly attracted to positive ions by their negative ends. (b) They are also strongly attracted to negative ions by their positive ends.

17. Polar water molecules are strongly attracted to each other.

18. Electron Configuration in Ionic Bonding
Metals tend to lose their valence electrons, leaving a complete octet in their next-lowest energy level.
Sodium – (1 valence electron) loses 1 electron and becomes Na+1.
Na ([Ne]3s1)  1e- + Na+1([Ne])
Calcium – (2 valence electrons) loses 2 electrons and becomes Ca+2.
Ca ([Ar]4s2)  2e- + Ca+2([Ar])

19. Electron Configuration in Ionic Bonding
Nonmetals tend to gain or share valence electrons to complete an octet in their highest energy level.
Oxygen – (6 valence electrons) gains two electrons to become O-2 .
O ([He]2s22p4) + 2e-  O-2 ([He] 2s22p6)
Phosphorus – (5 valence electrons) gains three electrons to become P-3.
P ([Ne]3s23p3) + 3e-  P-3 ([Ne] 3s23p6)

20. Formation and Properties of Ionic Compounds
Ionic bonds – forces of attraction that bind cations and anions together.
Ionic compound – consists of electrically neutral group of ions joined by electrostatic forces.
Example: Sodium chloride

21. Formation and Properties of Ionic Compounds
At room temperature, most ionic compounds are crystalline solids, where ions are arranged in various 3-D patterns.
Because of the large attractive forces of the ions to each other the compounds become very stable and have high melting points.

22. The shape of snowflakes results from bonding.
Source: ©Clyde H. Smith/Peter Arnold, Inc.

23. Sodium Chloride Crystals

24. The structure of lithium fluoride.

25. Electron Configuration in Ionic Bonding
Scientists have learned that all of the elements within each group behave similarly because they have the same number of valence electrons.
Valence electrons - # of electrons in the highest occupied energy level of an atom.
The number of valence electrons is related to the group numbers on the periodic table.

26. Electron Configuration in Ionic Bonding
Group 1 elements = 1 valence electron.
Group 2 elements = 2 valence electrons.
Groups 3-12 elements = 2 valence electrons.
Group 13 elements = 3 valence electrons.
Group 14 elements = 4 valence electrons.
Group 15 elements = 5 valence electrons.
Group 16 elements = 6 valence electrons.
Group 17 elements = 7 valence electrons.
Group 18 elements = 8 valence electrons.

27. Determining Valence Electrons for an Ion or a Compound
1. Multiply the number of valence electrons by the number of moles of each element.
2. Add up all the electrons for each of the elements.
3. If there is a charge and it is negative, add that number of electrons to the total.
4. If there is a charge and it is positive, subtract that number of electrons from the total.
Total # of electrons should always be an even number!

28. Determining Valence Electrons Examples
Determine the number of valence electrons in each of the following compounds and ions:
NH4+1
CH2ClBr
PO4-3

29. Electron Configuration in Ionic Bonding
Valence electrons are the only electrons involved in bonding, and are the only ones written when drawing electron dot structures.
In forming compounds, atoms tend to achieve the electron configuration of a noble gas, having 8 valence electrons which as known as having a stable octet (octet for 8 valence electrons).

30. Lewis Symbols of Atoms and Ions
Also known as electron dot symbols
Use symbol of element to represent nucleus and inner electrons
Use dots around the symbol to represent valence electrons
put one electron on each side first, then pair
Elements in the same group have the same Lewis symbol
Because they have the same number of valence electrons
Cations have Lewis symbols without valence electrons
Anions have Lewis symbols with 8 valence electrons
Li• Be• •B• •C• •N• •O: :F: :Ne: • ••
Li• Li+1 :F: [:F:]-1 14

31. The Nature of Covalent Bonding
Structural formula – chemical formulas that show the arrangement of atoms in molecules and polyatomic ions.
Octet rule – atoms gain or lose electrons to acquire the stable electron configuration of a noble gas, usually having 8 valence electrons.

32. Lewis Structures
You can represent the formation of the covalent bond in H2 as follows: H . : +
This uses the Lewis dot symbols for the hydrogen atom and represents the covalent bond by a pair of dots. 2

33. Lewis Structures
The shared electrons in H2 spend part of the time in the region around each atom. : H
In this sense, each atom in H2 has a helium configuration. 2

34. . : : + Lewis Structures H Cl H Cl
The formation of a bond between H and Cl to give an HCl molecule can be represented in a similar way. H . : Cl + : H Cl
Thus, hydrogen has two valence electrons about it (as in He) and Cl has eight valence electrons about it (as in Ar). 2

35. Lewis Structures
Formulas such as these are referred to as Lewis electron-dot formulas or Lewis structures.
An electron pair is either a bonding pair (shared between two atoms) or a lone pair (an electron pair that is not shared).
bonding pair
lone pair : H Cl 2

36. The Nature of Covalent Bonding
Exceptions to the octet rule:
H needs 2 electrons to be stable
Be needs 4 electrons to be stable
B needs 6 electrons to be stable

37. The Nature of Covalent Bonding
Steps for Drawing Lewis-dot structures
Determine the number of valence electrons in the molecule.
- When drawing determining valence electrons for an ion, add electrons if it an anion, and subtract electrons if it is a cation.
The first element in the compound will be the central atom. Exception: hydrogen will never be the central atom.

38. The Nature of Covalent Bonding
Steps for Drawing Lewis-dot Structures
3. Use one pair of electrons to bond each outer or terminal atom to the central atom.
4. Make all outer or terminal atoms stable using the valence electrons.
5. Put any remaining electrons around the central atom as lone pairs.

39. The Nature of Covalent Bonding
Draw the Lewis structure for:
NH3
PO43-
CHFClBr
PF5-2

40. The Nature of Covalent Bonding
Single covalent bond – a bond in which two atoms share a pair of electrons.
Double covalent bond – a bond in which two atoms share two pairs of electrons.
Triple covalent bond – a bond in which two atoms share three pairs of electrons.

41. The Nature of Covalent Bonding
If you have used up all of the valence electrons and you still need two more electrons to make the central atom stable, you must have one double bond.
If you still need four more electrons to make the central atom stable, you must have either one triple bond or two double bonds.
Double and triple bonds exist most commonly between C, N, O, and S atoms.

42. The Nature of Covalent Bonding
Draw Lewis structures for:
NOCl
CO2 N2
SiO3-2

43. The Nature of Covalent Bonding
Resonance structures – molecules or ions that can have two or more different Lewis structures. They must contain a double bond to have any resonance structures.
Resonance structures don’t truly have a single bonds or a double bond, but a hybrid mixture of bonds where the extra bond is spread equally among the other single bonds.

44. Resonance Structures
Draw Lewis structures for:
NOCl
CO2 N2
SiO3-2

45. The Nature of Covalent Bonding
Diamagnetic – substance where all of the electrons of the central atom are paired or bonded with other atoms.
Paramagnetic – substance where all of the electrons of the central atom are not paired or bonded with other atoms.

46. The Nature of Covalent Bonding
Single bonds are longer (length between the atoms) than double and triple bonds.
Double bonds are longer than triple bonds.
Single bonds are not as strong as double bonds, and can be broken much easier than double bonds.
Triple bonds are stronger than double bonds.

47. Bonding Theory
The valence-shell electron pair repulsion (VSEPR) model predicts the shapes of molecules and ions by assuming that the valence shell electron pairs are arranged as far from one another as possible.
To predict the relative positions of atoms around a given atom using the VSEPR model, you first note the arrangement of the electron pairs around that central atom. 2

48. Predicting Molecular Geometry
The following rules and figures will help discern electron pair arrangements.
Draw the Lewis structure
Determine how many bonding pairs are around the central atom. Count a multiple bond as one pair.
Determine how many lone pairs, if any, are around the central atom.
All diatomic molecules have a linear shape. 2

49. Arrangement of Electron Pairs About an Atom
Linear
3 pairs
Trigonal planar
4 pairs
Tetrahedral
5 pairs
Trigonal bipyramidal
6 pairs
Octahedral

52. Molecular Geometry Examples
NH3
PO43-
CHFClBr
PF5
SeF6
NOCl
CO2
SF2 N2
SiO3-2

53. Polar Bonds and Molecules
Nonpolar covalent bond – equal sharing of electrons between two atoms.
Polar covalent bond – unequal sharing of electrons between two atoms.
In polar covalent bonds the electrons are pulled closer to the atom with the larger electronegativity value.
This creates a partial positive and a partial negative pole within the bond.

54. Polar Bonds and Molecules
Polar bonds can create polar or nonpolar molecules and ions.
If the centers of partial positive and partial negative charge are in the same location, the molecule or ion is nonpolar.
If the centers of partial positive and partial negative charge are in different locations, the molecule or ion is polar.

55. Polar Bonds and Molecules
The easiest way to determine if a molecule or ion is polar or nonpolar is to look at the central atom.
If the central atom has lone pairs of electrons, the molecule or ion is polar.
If the central atom does not have any lone pairs of electrons, the molecule or ion is nonpolar.

56. Examples: Polar or Nonpolar?
Determine whether each of the following molecules or ions are polar or nonpolar:
NO2-1 N2
CN-1
CH4
SO3-2

57. Polar Bonds and Molecules
Attractions Between Molecules
Molecules are attracted to one another by a variety of forces.
These intermolecular forces are weaker than ionic or covalent bonds.
These forces are responsible for whether or not a molecular compound is a solid, liquid, or a gas.

58. Polar Bonds and Molecules
van der Waals forces – consist of dispersion forces and dipole interactions (dipole-dipole moments).
Dispersion forces – weakest of all intermolecular forces. They are caused by the motion of electrons. The strength of dispersion forces increases with the increasing number of electrons in a molecule.

59. Polar Bonds and Molecules
All molecules contain dispersion forces.
As molar mass and the number of electrons increase, dispersion forces increase.
Halogens are the most common molecules to have dispersion forces. Fluorine is a gas, Bromine is a liquid and Iodine is a solid.

60. Polar Bonds and Molecules
Dipole interactions – occur when polar molecules or ions are attracted to one another. This occurs when a partial positive charge and a partial negative charge come close to each other.
Dipole interactions are very similar to, but much weaker than ionic bonds.

61. Polar Bonds and Molecules
Hydrogen bonds – force exerted between a hydrogen atom bonded to an F, O, or N atom in one molecule and an unshared pair on another F, O, or N atom in a nearby molecule.
Hydrogen bonds can have a great effect on the boiling point of a substance.

62. Intermolecular Forces Examples