1. Structure and Properties of Ionic and Covalent Compounds
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Chapter 4
Structure and Properties of Ionic and Covalent Compounds
Denniston Topping Caret
4th Edition

2. 4.1 Chemical Bonding
Chemical Bond - the force of attraction between any two atoms in a compound.
Interactions involving valence electrons are responsible for the chemical bond.

3. 4.1 Chemical Bonding Lewis Symbols
Lewis symbol (Lewis structure) - a way to represent atoms (and their bonds) using the element symbol and valence electrons as dots.
4.1 Chemical Bonding

4. Principal Types of Chemical Bonds: Ionic and Covalent1
4.1 Chemical Bonding
Ionic bond - a transfer of one or more electrons from one atom to another.
forms attractions due to the opposite charges of the atoms.
Covalent bond - attractive force due to the sharing of electrons between atoms.

5. 4.1 Chemical Bonding IONIC BONDING Let’s examine the formation of NaCl
Na + Cl  NaCl
Chlorine has a high
electron affinity.
When chlorine gains
an electron, it gains
the Ar configuration.
4.1 Chemical Bonding
Sodium has a
low ionization energy
it readily loses this
electron .
When Sodium loses
the electron, it gains
the Ne configuration.
Na  Na+ + e-

6. 4.1 Chemical Bonding Essential Features of Ionic Bonding
Atoms with low I.E. and low E.A. tend to form positive ions.
Atoms with high I.E. and high E.A. tend to form negative ions.
Ion formation takes place by electron transfer.
The ions are held together by the electrostatic force of the opposite charges.
Reactions between metals and nonmetals (representative) tend to be ionic.
4.1 Chemical Bonding

7. 4.1 Chemical Bonding COVALENT BONDING
Let’s look at the formation of H2
H + H  H2
Each hydrogen has one electron in it’s valance shell.
If it were an ionic bond it would look like this:
4.1 Chemical Bonding
However, both hydrogen atoms have the same tendency to gain or lose electrons.
Both gain and loss will not occur.

8. Instead, each atom gets a noble gas configuration by sharing electrons.
4.1 Chemical Bonding
Each Hydrogen atom now has two electrons around it and has a He configuration
The shared
electrons
pair is a
Covalent Bond

9. 4.1 Chemical Bonding HOFBrINCl Features of Covalent Bonds
Covalent bonds tend to form between atoms with similar tendency to gain or lose electrons.
The diatomic elements have totally covalent bonds (totally equal sharing.)
H2, N2, O2, F2, Cl2, Br2, I2
4.1 Chemical Bonding
HOFBrINCl
Each fluorine has eight electrons around it. Ne’s configuration.

10. Other examples:
4.1 Chemical Bonding

11. Polar Covalent Bonding and Electronegativity1
The Polar Covalent Bond
Ionic bonding involves the transfer of electrons.
Covalent bonding involves the sharing of electrons.
Polar covalent bonding - bonds made up of unequally shared electron pairs.
4.1 Chemical Bonding

12. 4.1 Chemical Bonding These two
somewhat positively charged
somewhat negatively charged
These two
electrons
are not shared
equally.
The electrons spend more time with fluorine.
This sets up a polar bond
4.1 Chemical Bonding
A truly covalent bond can only occur when both atoms are identical.
Electronegativity is used to determine if a bond is polar and who gets the electrons the most.

13. Electronegativity - a measure of the ability of an atom to attract electrons in a chemical bond.
4.1 Chemical Bonding
electronegativity increases
electronegativity increases

14. The greater the difference in electronegativity between two atoms, the greater the polarity of a bond.
Which would be more polar, a H-F bond or a H-Cl bond?
H-F … = 1.9
H-Cl… = 0.9
Therefore, the HF bond is more polar than the HCl bond.
4.1 Chemical Bonding

15. 4.2 Naming Compounds and Writing Formulas of Compounds
Nomenclature - the assignment of a correct and unambiguous name to each and every chemical compound.
We will learn two systems
one for naming ionic compounds and
one for naming covalent compounds.

16. Section 4.2 Outline of Nomenclature Topics I. Ionic Compounds
A. Writing Formulas of Ionic Compounds from the Identities of the Component Ions.
B. Writing Names of Ionic Compounds from the Formula of the Compound
C. Writing Formulas of Ionic Compounds from the Name of the Compound
IV. Covalent Compounds
A. Naming Covalent Compounds
B. Writing Formulas of Covalent Compounds
Section 4.2

17. Section 4.2 I. Ionic Compounds
Metals and nonmetals usually react to form ionic compounds.
The metals are the cations and the nonmetals are the anions.
The cations and anions arrange themselves in a regular three-dimensional repeating array called a crystal lattice.
Section 4.2

18. Here is the crystal lattice of sodium chloride (table salt.)
Formula - the smallest whole number ratio of ions in the crystal.
The formula for sodium chloride is NaC.
Section 4.2

19. Section 4.2 A. Writing Formulas of Ionic Compounds
Determine the charge of the ions (usually can be obtained from the group number.)
Cations and anions must combine to give a formula with a net charge of zero
It must have the same number of positive charges as negative charges.
Section 4.2

20. Section 4.2 Predicting Formulas
Predict the formula of the ionic compounds formed from the following combination of ions:
1. sodium and sulfur
2. magnesium and oxygen
3. aluminum and oxygen
4. barium and fluorine
Section 4.2

21. Section 4.2 B. Writing Names of Ionic Compounds from the Formula 2
Stock System:
1. Name cation followed by the name of anion.
2. Give anion the suffix -ide.
Examples:
NaCl is sodium chloride.
AlBr3 is aluminum bromide.
Section 4.2

22. If the cation of an element has several ions of different charges (as with Transition metals) use a Roman numeral after the metal name.
Roman numeral gives the charge of the metal.
Examples:
FeCl3 is iron(III) chloride
FeCl2 is iron(II) chloride
CuO is copper(II) oxide
Section 4.2

23. Section 4.2 Common Nomenclature System
Use -ic to indicate the higher of the charges.
Use -ous to indicate the lower of the charges.
Examples:
FeCl2 is ferrous chloride
FeCl3 is ferric chloride
Cu2O is cuprous oxide
CuO is cupric oxide
Section 4.2

24. Note that the common system requires you to know the common charges of the metals and use the Latin names of the metals.
Monatomic ions - ions consisting of a single atom.
Examples
K+ potassium ion
Ba2+ barium ion
O2- oxide ion
Table 4.2 gives common monatomic cations and anions.
Section 4.2

25. Polyatomic ions - ions composed of 2 or more atoms bonded together.
Within the ion, the atoms are bonded using covalent bonds.
The ions will be bonded to other ions with ionic bonds.
Examples:
NH4+ ammonium ion
SO42- sulfate ion
Section 4.2

26. Section 4.2 Table 4.3 Common Polyatomic Cations and Anions Ion Name
NH4+ ammonium
NO3- nitrate
SO sulfate
OH- hydroxide
CN- cyanide
PO phosphate
CO carbonate
HCO3- bicarbonate
CH3COO- (or C2H3O2-) acetate
Section 4.2
Table 4.3 in your textbook gives a more complete list of polyatomic ion.

27. Let's Practice Nomenclature
Name the following compounds using the Stock System
1. NH4Cl
2. BaSO4
3. Fe(NO3)3
4. CuHCO3
5. Ca(OH)2
Section 4.2

28. Section 4.2 C. Writing Formulas of Ionic Compounds from the Name 2
Determine the charge on the ions.
Write the formula so the compounds are neutral.
Example:
Barium chloride:
Barium is +2, Chloride is -1
Formula is BaCl2
Section 4.2

29. Let's Practice
Give the formula for the following ionic compounds:
1. sodium sulfide
2. copper(I) carbonate
3. magnesium phosphate
4. chromium(II) sulfate
Section 4.2

30. Section 4.2 II. Covalent Compounds
Covalent compounds are usually formed from nonmetals.
Molecules - compounds characterized by covalent bonding.
not a part of a massive three dimensional crystal structure.
Section 4.2

31. Section 4.2 A. Naming Covalent Compounds 3
1. The names of the elements are written in the order in which they appear in the formula.
2. A prefix indicates the number of each kind of atom
mono- 1 hexa- 6
di- 2 hepta- 7
tri- 3 octa- 8
tetra- 4 nona- 9
penta- 5 deca- 10
Section 4.2

32. 3. If only one atom of a particular kind is present in the molecule, the prefix mono- is usually omitted from the first element.
Example: CO is carbon monoxide
4. The stem of the name of the last element is used with the suffix -ide
Section 4.2
5. The final vowel in a prefix is often dropped before a vowel in the stem name.

33. Section 4.2 Let's Practice Name the following covalent compounds
1. SiO2
2. N2O5
3. CCl4
4. IF7
Section 4.2

34. Section 4.2 B. Writing Formulas of Covalent Compounds
Use the prefixes in the names to determine the subscripts for the elements.
Examples:
nitrogen trichloride:
NCl3
diphosphorus pentoxide
P2O5
Some common names that are used: 2
Section 4.2
H2O water
NH ammonia
C2H5OH ethanol
C6H12O6 glucose

35. 4.3 Properties of Ionic and Covalent Compounds
Physical State
Ionic compounds are solids at room temperature
Covalent compounds are solids, liquids and gases
Melting and Boiling Points
melting point - the temperature at which a solid is converted to a liquid
boiling point - the temperature at which a liquid is converted to a gas

36. 4.3 Properties Melting and Boiling Points
Ionic compounds have much higher melting points and boiling points than covalent compounds.
Due to the large amount of energy required to break the electrostatic attractions between ions.
Ionic compounds typically melt at several hundred degrees Celsius.
Structure of Compounds in the Solid State
Ionic compounds are crystalline
Covalent compounds are crystalline or amorphous - have no regular structure.
4.3 Properties

37. 4.3 Properties Solutions of Ionic and Covalent Compounds
Ionic compounds often dissolve in water, when they do they dissociate - form positive and negative ions in solution.
Electrolytes -ions present in solution allow the solution to conduct electricity.
Covalent solids usually do not dissociate and do not conduct electricity - nonelectrolytes
4.3 Properties

38. 4.4 Drawing Lewis Structures on Molecules and Polyatomic Ions
Lewis Structure Guidelines
1. Use chemical symbols for the various elements to write the skeletal structure of the compound.
the least electronegative atom will be placed in the central position,
hydrogen and fluorine occupy terminal positions,
carbon often forms chains of carbon-carbon covalent bonds. 5

39. 4.4 Drawing Lewis Structures
2. Determine the total number of valence electrons associated with each atom in the compound.
for polyatomic cations, subtract one electron for every positive charge;
for polyatomic anions, add one electron for every negative charge.
3. Connect the central atom to each of the surrounding atoms using electron pairs. Then give each atom an octet.
Remember, hydrogen needs only two electrons
4.4 Drawing Lewis Structures

40. 4.4 Drawing Lewis Structures
Count the number of electrons you have and compare to the number you used.
If they are the same, you are finished.
If you used more electrons than you have add a bond for every two too many you used. Then give every atom an octet.
If you used less electrons than you have….(see later when discuss exceptions to the octet rule)
5. Check that all atoms have the octet rule satisfied and that the total number of valance electrons are used.
4.4 Drawing Lewis Structures

41. 4.4 Drawing Lewis Structures
Let's Practive Lewis Structures
Using the guidelines presented, write Lewis structures for the following:
1. H2O
2. NH3
3. CO2
4. NH4+
5. CO32-
6. N2
4.4 Drawing Lewis Structures

42. 4.4 Drawing Lewis Structures
Lewis Structure, Stability, Multiple Bonds, and Bond Energies 6
Single bond - one pair of electrons are shared between two atoms
Double bond - two pairs of electrons are shared between two atoms
Triple bond - three pairs of electrons are shared between two atoms
4.4 Drawing Lewis Structures

43. 4.4 Drawing Lewis Structures
Bond energy - the amount of energy required to break a bond holding two atoms together.
triple bond > double bond > single bond
Bond length - the distance separating the nuclei of two adjacent atoms.
single bond > double bond > triple bond
4.4 Drawing Lewis Structures

44. 4.4 Drawing Lewis Structures
Lewis Structures and Resonance
Write the Lewis structure at CO32- on your paper.
If you look at the people around you they probably put the double bond in different places.
Who is right?
You all are.
4.4 Drawing Lewis Structures

45. 4.4 Drawing Lewis Structures
Experimental evidence shows all bonds are the same length. Shorter than single bond but longer than a double bond.
None of the three exist but the actual structure is an average or hybrid of these three Lewis structures.
Resonance - two or more Lewis structures that contribute to the real structure.
4.4 Drawing Lewis Structures

46. 4.4 Drawing Lewis Structures
Lewis Structures and Exceptions to the Octet Rule
1. Incomplete Octet - less then eight electrons around an atom other than H.
Let’s look at BF3
2. Odd Electron - if there is an odd number of valence electrons it isn’t possible to give every atom eight electrons.
Let’s look at NO
3. Expanded Octet - elements in 3rd period and below may have 10 and 12 electrons around it.
4.4 Drawing Lewis Structures

47. 4.4 Drawing Lewis Structures
Expanded octet is the most common exception.
Write the Lewis structures of SF6 and SF4
4.4 Drawing Lewis Structures

48. 4.4 Drawing Lewis Structures
Lewis Structures and Molecular Geometry: VSEPR Theory 7
VSEPR theory - valance shell electron pair repulsion theory.
This is used to predict the shape of the molecules.
Electrons around the central atom (both bonding and nonbonding pairs) arrange themselves so they can be as far away from each other as possible.
4.4 Drawing Lewis Structures

49. 4.4 Drawing Lewis Structures
We will do several examples in order to see how the geometry is determined.
Let’s do BeH2 on the overhead.
4.4 Drawing Lewis Structures

50. 4.4 Drawing Lewis Structures
Let’s do BF3 on the overhead.
4.4 Drawing Lewis Structures
Let’s do CH4 on the overhead

51. 4.4 Drawing Lewis Structures
Let’s do NH3 on the overhead
Trigonal Pyramidal
4.4 Drawing Lewis Structures
Let’s do H2O on the overhead
Bent

52. 4.4 Drawing Lewis Structures
The basic procedure to follow to determine the shape:
1. Write the Lewis structure.
2. Count the number of bonded atoms and non-bonded electrons around the central atom.
2 legs - linear
3 legs - trigonal planer
4 legs - tetrahedron
3. Look at the atoms and name the shape.
These would include: linear, trigonal planer, bent, trigonal pyramid, tetrahedon.
4.4 Drawing Lewis Structures

53. 4.4 Drawing Lewis Structures
Determine the Geometry
of the Following
PCl3
SO2
4.4 Drawing Lewis Structures

54. 4.4 Drawing Lewis Structures
Lewis Structures and Polarity
Polar molecules when placed in an electric field will align themselves in the field.
Molecules that are polar behave as a dipole (two “poles”)
One end is positively charged the other is negatively charged.
Nonpolar molecules will not align themselves in an electric field.
4.4 Drawing Lewis Structures

55. 4.4 Drawing Lewis Structures
To determine if a molecule is polar or has a dipole:
Write the Lewis structure.
Draw the Geometry.
Use the following symbol to denote the polarity of each bond.
4.4 Drawing Lewis Structures
This end of the arrow points to the negative end of the bond.
The more electronegative atom attracts the electrons more strongly towards it.
This end of the arrow points to the positive end of the bond.
The less electronegative atom.

56. 4.4 Drawing Lewis Structures
Let's do a few examples.
Determine the polarity of the following molecules:
1. O2
2. HF
3. CH4
4. H2O
4.4 Drawing Lewis Structures

57. 4.5 Properties Based on Molecular Geometry8 9
Intramolecular forces - attractive forces within molecules. (Chemical bonds)
Intermolecular forces - attractive forces between molecules.
We will look at how intermolecular forces affect:
I. Solubility.
II. Boiling points and melting points.

58. 4.5 Properties Based on Molecular Geometry
I. Solubility - the maximum amount of solute that dissolves in a given amount of solvent at a specific temperature.
“Like dissolves like”
polar molecules are most soluble in polar solvents
nonpolar molecules are most soluble in nonpolar solvents.
Will ammonia (NH3) dissolve in water?
Yes, both molecules are polar.
4.5 Properties Based on Molecular Geometry

59. 4.5 Properties Based on Molecular Geometry

60. 4.5 Properties Based on Molecular Geometry
What do you know about oil and water?
“They don’t mix”
Why?
Because water is polar and oil in nonpolar.
4.5 Properties Based on Molecular Geometry

61. 4.5 Properties Based on Molecular Geometry
II. Boiling Points and Melting Points
The greater the intermolecular force the higher the melting point and boiling point.
Two factors to consider:
Larger molecules have higher m.p. and b.p. than smaller molecules.
Polar molecules have higher m.p. and b.p. than nonpolar molecules (of similar molecular mass.)
4.5 Properties Based on Molecular Geometry

62. The End Chapter 4