1. Chapter 6
Chemical Bonding

2. Introduction to Chemical Bonding
Section 6-1
Introduction to Chemical Bonding

3. Types of Chemical Bonding
Very rarely do atoms exist by themselves.
They usually exist in combinations with each other.
These combinations are held together by chemical bonds.

4. What is a Chemical Bond?
It is an electrical attraction between the nuclei of one atom and the valence electrons of another that binds the atoms together.
It occurs to increase stability.

5. There are Two Types of Chemical Bonds.
Ionic Bonding – The bonding that occurs between large numbers of metal cations and nonmetal anions. It is the attraction of opposite charges.
Covalent Bonding – The bonding that occurs between two elements that share electron pairs.

6. Usually, a Bond is neither purely ionic or covalent.
The ionic or covalent character of a bond is determined by electronegativity.
What is electronegativity?

7. Electronegativity
The ability of an atom to attract electrons to itself.
If the electronegativity difference is less than 0.5, it is covalent.
If it is between 0.6 and 1.9, it is polar.
If it is greater than 2, it is ionic.

8. What is a Polar Covalent Bond?
It is a bond where the electrons are pulled more closely toward the electronegative element.

9. Polar Covalent Bonds, cont.
This gives the electronegative atom a slightly negative charge, and the electropositive atom a slightly positive charge.
It is a bond where electrons are shared unequally.

10. Covalent Bonding and Molecular Compounds
Section 6-2
Covalent Bonding and Molecular Compounds

11. Molecular Compounds Many compounds are made up of molecules.
A molecule is a neutral group of atoms held together by covalent bonds.
A chemical compound whose simplest units are molecules is called a molecular compound.

12. Chemical Formulas
A chemical compound is represented by a chemical formula.
Chemical formulas indicate the number of each type of atom in a chemical compound by using atomic symbols and numerical subscripts.
A molecular formula represents a molecular compound.

13. A Sample Molecular Formula
The molecular formula H2O is the formula for water.
It indicates that water is composed of two atoms of hydrogen, and one atom of oxygen.

14. Diatomic Molecules
There is a special class of molecule called a diatomic molecule.
These are formed by atoms that cannot exist alone, so if they are pure, they must be in pairs.
IAS for Chemistry – C.1.28: Explain that chemical bonds between atoms in molecules, such as H2, N2, O2, F2, Cl2, I2, and Br2 and many large biological molecules are covalent.

15. Diatomic Molecules There are seven diatomic molecules.
These are: H2, N2, O2, F2, Cl2, Br2, and I2
These molecules are held together by purely covalent bonds.

16. Every Covalent Bond has Two Properties
Bond Length
Bond Energy

17. Bond Length The average distance between two bonded atoms.
Measured in picometers (pm)

18. Bond Energy The energy required to break a chemical bond.
Measured in kilojoules per mole. (kJ/mol)
These two properties are characteristic of every bond.
Table 6.1 on page 168 shows some common values.

19. Noble Gases Noble Gases do exist by themselves.
This is because they have 8 valence electrons.
Other atoms must achieve this through covalent or ionic bonding.
This helps them satisfy the octet rule.

20. The Octet Rule
Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons has 8 valence electrons in its outer energy level.
The exceptions are H, Li, Be, and B, which can have less than 8.

21. Extended Valence
Some elements can have more than 8 valence electrons when they combine with elements like F, O, or Cl.
These cases of extended valence are rare, and are not important for us.

22. Electron Dot Notation
Covalent bonding usually only involves an atom’s valence electrons.
We use electron dot notation to keep track of these electrons.
Electron dot notation is an electron configuration in which valence electrons are shown as dots around the element’s symbol.
The number of dots is equal to the group number for all group A elements.
This can also be used for molecules.

23. Electron Dot Notation
In a molecule, a pair of dots between two atoms represents a shared pair of electrons, or what we call bonding electrons.
A pair of electrons around an element that are not shared with another element are non – bonding electrons or unshared electron pairs.

24. Lewis Structures
The bonding pair of electrons can also be drawn as a dash.
This then becomes a Lewis Structure.
Lewis structures are formulas in which bonding electrons are drawn as a dash and nonbonding electrons are drawn as dots.

25. Structural Formulas
Lewis structures can also be drawn without dots around them.
These are structural formulas.

26. Multiple Bonds
Up until now, every molecule we have looked at has a single pair of shared electrons between atoms.
It is possible to have more than one pair that are shared between the same two atoms. These result in multiple bonds.
If two pairs of electrons are shared, that results in a double bond.

27. Multiple Bonds, cont.
If three pairs of electrons are shared, it is a triple bond.
The octet rule must still be followed in multiple bonded molecules.
Anytime C, N, or O are in a molecule, multiple bonds are possible.

28. Resonance
Some molecules cannot be accurately shown with one structure.
These are molecules with alternating single and double bonds.
Neither form really exists. The true form is a combination of the two.
This is called resonance.

29. Resonance Each structure is called a resonance structure.
The true form of the molecule is called a resonance hybrid.
A double headed arrow between Lewis structures indicates that they are resonance forms.

30. Ionic Bonding and Ionic Compounds
Section 6-3
Ionic Bonding and Ionic Compounds

31. Ionic Compounds
To this point, we have only discussed molecular compounds. A second type of compound is an ionic compound.
This is a compound made up of positive and negative ions that combine into a neutral substance.
These exists as crystals of many positive and negative ions.

32. Ionic Crystals
A three – dimensional arrangement of ions in a regular, repeating pattern.
Instead of writing an entire crystal, we usually represent an ionic compound with a formula unit.

33. Formula Unit
A formula unit is the simplest collection of atoms from which an ionic compound’s formula can be determined.
This ratio is determined by ionic charge.

34. Crystal Lattice An ionic crystal is known as a crystal lattice.
For example, in NaCl, each sodium attracts six negative chlorines, and every chlorine attracts six positive sodiums.
All of these additional bonds makes a crystal very strong and stable.

35. Lattice Energy
It takes a great deal of energy to break down an ionic compound.
This energy is known as lattice energy.
Lattice energy is the energy released when one mole of a crystal forms.
IAS for Chemistry – C.1.8: Use formulas and laboratory investigations to classify substances as metal or nonmetal, ionic or molecular, acid or base, and organic or inorganic.

36. What are the differences between Ionic and Molecular Compounds?
Ionic Compounds:
Have High Melting Points.
Have High Boiling Points.
Are Very Hard and Brittle.
Conduct Electricity in water.
Molecular Compounds:
Have Low Melting Points.
Have Low Boiling Points.
Are Soft.
Do not Dissolve in Water.
IAS for Chemistry – C.1.35: Infer and explain the physical properties of substances, such as melting points, boiling points, and solubility, based on the strength of molecular attractions.

37. Why do Ionic Compounds Conduct Electricity?
Scientists have known for a long time that ionic substances conduct electricity when they are dissolved in water. In fact, another term for an ionic substance is an electrolyte.
The problem was, no one knew how or why this occurred.

38. Svante August Arrhenius (1859 – 1927)
Arrhenius was one of the scientists studying electrolytes.
He concluded that electrolytes split, or dissociated in water into positive and negative ions.
The degree to which this happened depended on the ionic character of the substance, and its concentration.

39. Svante August Arrhenius (1859 – 1927)
The work of Arrhenius not only described how ionic compounds are put together, but also greatly affected both chemistry and biology, allowing us to describe simple reactions like why adding ionic compounds increases freezing and boiling points, to a greater understanding of digestion and absorption.
IAS for Chemistry – C.2.5: Explain how Arrhenius’ discovery of the nature of ionic solutions contributed to the understanding of a broad class of chemical reactions.

40. Polyatomic Ions
A charged group of covalently bonded atoms are known as a polyatomic ion.
There are many polyatomic ions, and you will be responsible for memorizing a few of them.

41. Polyatomic Ions you Should Know
NH4+1 – Ammonium
H3O+1 – Hydronium
CO3-2 – Carbonate
ClO3-1 – Chlorate
ClO2-1 – Chlorite
CrO4-2 – Chromate
HCO3-1 – Hydrogen Carbonate
HSO4-1 – Hydrogen Sulfate
OH-1 – Hydroxide
NO3-1 – Nitrate
NO2-1 – Nitrite
PO4-3 – Phosphate
SO4-2 – Sulfate
SO3-2 - Sulfite

42. Section 6-4
Metallic Bonding

43. Metallic Bonding
There is a third type of bonding that is different from ionic or covalent bonding. This is metallic bonding.
Metallic bonding accounts for the different properties metals have.

44. Metallic Bonding Most metals have few valence electrons.
The s-block metals have empty p subshells, and the transition metals have empty p and d subshells. These empty orbitals tend to overlap.
This allows valence electrons to freely move across the metal from atom to atom.

45. The Sea of Electrons These electrons are said to be delocalized.
The bonding that results from the attraction of metal atoms to the sea of electrons is metallic bonding.

46. Metallic Bonding Accounts for Several Unusual Properties of Metals
Conductivity
Malleability
Ductility
The electrons surround each atom, and keep opposite charges from contacting each other.
Allows the atoms to slide past each other, without repelling.

47. Section 6-5
Molecular Geometry

48. The properties of molecules depends not only on their bonding, but also on their molecular geometry.
Molecular geometry is determined mostly by molecular polarity – the uneven distribution of molecular charge.

49. VSEPR Theory
The theory used to determine molecular geometry is VSEPR Theory.
VSEPR stands for Valence Shell Electron Pair Repulsion.
It states that the valence level electron repulsion around an atom causes the sets of electrons to be oriented as far apart as possible.
There are six geometric shapes molecules form that we will deal with in this class.

50. Shape 1
Any molecule made of only two atoms is called linear diatomic.

51. Shape 2
Any molecule made of three atoms with no pair of free electrons on the center atom is linear triatomic.
An example would be Calcium Fluoride, CaF2.

52. Shape 3
Any molecule with three atoms and a pair of free electrons on the center atom is Bent Triatomic.
An example of this would be water, H2O.

53. Shape 4
Any molecule with four atoms and no free electron pairs on the center atom is trigonal planar.
An example of this is Aluminum Chloride, AlCl3.

54. Shape 5
Any molecule with four atoms and a pair of free electrons on the central atom is called pyramidal.
An example of this is ammonia, NH3.

55. Shape 6 Any molecule made of five atoms is tetrahedral.
An example of this is methane, CH4.
There are other types of VSEPR shapes, but they only apply to molecules with extended valence, and will not be studied in this class.

56. Hybridization
The VSEPR shapes of molecules are possible because the orbitals in the molecules become rearranged.
This is a process called hybridization.
Hybridization is the blending of two or more different atomic orbitals to create orbitals of identical shape and energy.

57. There are three types of hybridization
Sp3 – one s and three p orbitals in an energy level combine to form four equal orbitals.
Sp2 – one s and two p orbitals in an energy level combine to form three equal orbitals.
Sp – one s and one p orbital in an energy level combine to form two equal orbitals.

58. Sp3 Hybridization The sp3 hybridization results in all single bonds.
It forms the tetrahedral, pyramidal, or bent VSEPR shapes.

59. Sp2 Hybridization
The sp2 hybrid results in either single bonds, or a double bond.
It forms the trigonal planar VSEPR shape.

60. Sp Hybridization
The sp hybrid allows for single, double, or triple bonding.
It results in molecules with a linear VSEPR shape.

61. There are two types of covalent bonds.
Sigma Bonds – form from the head on overlap of two atomic orbitals.
Pi Bonds – form from the sideways overlap of two unhybridized p orbitals.
A single bond is one sigma bond.
A double bond is one sigma and one pi bond.
A triple bond is one sigma and two pi bonds.

62. Intermolecular Forces
These are forces between two or more molecules.
These are relatively weak forces, but in large numbers, can affect things like boiling and melting points.

63. There are three types of Intermolecular Forces.
Dipole – Dipole Forces
Hydrogen Bonding
London Dispersion Forces

64. Dipole – Dipole Forces Forces of attraction between polar molecules.
Polar molecules have equal but opposite charges separated by a short distance.
The negative region of one polar molecule attracts the positive region of another.
We use arrows to point to the negative pole. Not all polar molecules are dipoles.

65. Hydrogen Bonding A very strong dipole – dipole force.
A hydrogen ion in one polar molecule is attracted to a strongly electronegative element (O, Cl, F, N) in another.
Hydrogen bonding is what causes ice to expand when it is frozen. It also causes high boiling and melting points in water.
IAS for chemistry – (C.1.36): Describe the nature of ionic, covalent, and hydrogen bonds, and give examples of how they contribute to the formation of various types of compounds.

66. London Dispersion Forces
Intermolecular attraction resulting from the constant motion of electrons.
Creates instantaneous dipoles.
This acts only on noble gases and nonpolar molecules.