1. Chapter 7 “Ionic Bonding”
Chemistry

2. Section Ions
OBJECTIVES:
Determine the number of valence electrons in an atom of a representative element.

3. Section Ions
OBJECTIVES:
Explain how the octet rule applies to atoms of metallic and nonmetallic elements.

4. Describe how cations form.
Section Ions
OBJECTIVES:
Describe how cations form.

5. Explain how anions form.
Section Ions
OBJECTIVES:
Explain how anions form.

6. Valence Electrons are…?
The electrons responsible for the chemical properties of atoms, and are those in the outer energy level.
Valence electrons - The s and p electrons in the outer energy level
the highest occupied energy level
Core electrons – are those in the energy levels below.

7. Keeping Track of Electrons
Atoms in the same column...
Have the same outer electron configuration.
Have the same valence electrons.
The number of valence electrons are easily determined. It is the group number for a representative element
Group 2: Be, Mg, Ca, etc.
have 2 valence electrons

8. Electron Dot diagrams are…
A way of showing & keeping track of valence electrons.
How to write them?
Write the symbol - it represents the nucleus and inner (core) electrons
Put one dot for each valence electron (8 maximum)
They don’t pair up until they have to (Hund’s rule) X

9. The Electron Dot diagram for Nitrogen
Nitrogen has 5 valence electrons to show.
First we write the symbol. N
Then add 1 electron at a time to each side.
Now they are forced to pair up.
We have now written the electron dot diagram for Nitrogen.

10. The Octet Rule
In Chapter 6, we learned that noble gases are unreactive in chemical reactions
In 1916, Gilbert Lewis used this fact to explain why atoms form certain kinds of ions and molecules
The Octet Rule: in forming compounds, atoms tend to achieve a noble gas configuration; 8 in the outer level is stable
Each noble gas (except He, which has 2) has 8 electrons in the outer level

11. Formation of Cations
Metals lose electrons to attain a noble gas configuration.
They make positive ions (cations)
If we look at the electron configuration, it makes sense to lose electrons:
Na 1s22s22p63s1 1 valence electron
Na1+ 1s22s22p6 This is a noble gas configuration with 8 electrons in the outer level.

12. Electron Dots For Cations
Metals will have few valence electrons (usually 3 or less); calcium has only 2 valence electrons Ca

13. Electron Dots For Cations
Metals will have few valence electrons
Metals will lose the valence electrons Ca

14. Electron Dots For Cations
Metals will have few valence electrons
Metals will lose the valence electrons
Forming positive ions
Ca2+
This is named the “calcium ion”.
NO DOTS are now shown for the cation.

15. Electron Dots For Cations
Let’s do Scandium, #21
The electron configuration is: 1s22s22p63s23p64s23d1
Thus, it can lose 2e- (making it 2+), or lose 3e- (making 3+)
Sc = Sc2+ Sc
= Sc3+
Scandium (II) ion
Scandium (III) ion

16. Electron Dots For Cations
Let’s do Silver, element #47
Predicted configuration is: 1s22s22p63s23p64s23d104p65s24d9
Actual configuration is: 1s22s22p63s23p64s23d104p65s14d10
Ag = Ag+ (can’t lose any more, charges of 3+ or greater are uncommon)

17. Electron Dots For Cations
Silver did the best job it could, but it did not achieve a true Noble Gas configuration
Instead, it is called a “pseudo-noble gas configuration”

18. Electron Configurations: Anions
Nonmetals gain electrons to attain noble gas configuration.
They make negative ions (anions)
S = 1s22s22p63s23p4 = 6 valence electrons
S2- = 1s22s22p63s23p6 = noble gas configuration.
Halide ions are ions from chlorine or other halogens that gain electrons

19. Electron Dots For Anions
Nonmetals will have many valence electrons (usually 5 or more)
They will gain electrons to fill outer shell. 3- P
(This is called the “phosphide ion”, and should show dots)

20. Stable Electron Configurations
All atoms react to try and achieve a noble gas configuration.
Noble gases have 2 s and 6 p electrons.
8 valence electrons = already stable!
This is the octet rule (8 in the outer level is particularly stable). Ar

21. Section 7.2 Ionic Bonds and Ionic Compounds
OBJECTIVES:
Explain the electrical charge of an ionic compound.

22. Section 7.2 Ionic Bonds and Ionic Compounds
OBJECTIVES:
Describe three properties of ionic compounds.

23. Ionic Bonding
Anions and cations are held together by opposite charges (+ and -)
Ionic compounds are called salts.
Simplest ratio of elements in an ionic compound is called the formula unit.
The bond is formed through the transfer of electrons (lose and gain)
Electrons are transferred to achieve noble gas configuration.

24. Ionic Compounds
Also called SALTS
Made from: a CATION with an ANION (or literally from a metal combining with a nonmetal)

25. Ionic Bonding Na Cl
The metal (sodium) tends to lose its one electron from the outer level.
The nonmetal (chlorine) needs to gain one more to fill its outer level, and will accept the one electron that sodium is going to lose.

26. Ionic Bonding
Na+
Cl -
Note: Remember that NO DOTS are now shown for the cation!

27. Ionic Bond Negative charges are attracted to positive charges.
Negative anions are attracted to positive cations.
The result is an ionic bond.
A three-dimensional crystal lattice of anions and cations is formed.

28. Preserve Electroneutrality
When ions combine, electroneutrality must be preserved.
In the formation of magnesium chloride,
2 Cl- ions must balance a Mg2+ ion:
Mg Cl- → MgCl2

29. Ionic Bonding
Let’s do an example by combining calcium and phosphorus: Ca P
All the electrons must be accounted for, and each atom will have a noble gas configuration (which is stable).

30. Ionic Bonding Ca P

31. Ionic Bonding
Ca2+ P

32. Ionic Bonding
Ca2+ P Ca

33. Ionic Bonding
Ca2+
P 3- Ca

34. Ionic Bonding
Ca2+
P 3- Ca P

35. Ionic Bonding
Ca2+
P 3-
Ca2+ P

36. Ionic Bonding Ca
Ca2+
P 3-
Ca2+ P

37. Ionic Bonding Ca
Ca2+
P 3-
Ca2+ P

38. Ionic Bonding
Ca2+
Ca2+
P 3-
Ca2+
P 3-

39. = Ca3P2 Ionic Bonding Formula Unit
This is a chemical formula, which shows the kinds and numbers of atoms in the smallest representative particle of the substance.
For an ionic compound, the smallest representative particle is called a: Formula Unit

40. Properties of Ionic Compounds
Crystalline solids - a regular repeating arrangement of ions in the solid: Fig. 7.9, page 197
Ions are strongly bonded together.
Structure is rigid.
High melting points
Coordination number- number of ions of opposite charge surrounding it

41. NaCl CsCl TiO2 - Page 198 Coordination Numbers:
Both the sodium and chlorine have 6
NaCl
Both the cesium and chlorine have 8
CsCl
Each titanium has 6, and each oxygen has 3
TiO2

42. Do they Conduct?
Conducting electricity means allowing charges to move.
In a solid, the ions are locked in place.
Ionic solids are insulators.
When melted, the ions can move around.
Melted ionic compounds conduct.
NaCl: must get to about 800 ºC.
Dissolved in water, they also conduct (free to move in aqueous solutions)

43. - Page 198
The ions are free to move when they are molten (or in aqueous solution), and thus they are able to conduct the electric current.

44. Section 9.1 Naming Ions OBJECTIVES:
Model the valence electrons of metal atoms.

45. Section 9.1 Naming Ions
OBJECTIVES:
Identify the charges on monatomic ions by using the periodic table, and name the ions.

46. Section 9.1 Naming Ions OBJECTIVES:
Define a polyatomic ion and write the names and formulas of the most common polyatomic ions.

47. Identify the two common endings for the names of most polyatomic ions.
Section 9.1 Naming Ions
OBJECTIVES:
Identify the two common endings for the names of most polyatomic ions.

48. Atoms and Ions Atoms are electrically neutral.
Because there is the same number of protons (+) and electrons (-).
Ions are atoms, or groups of atoms, with a charge (positive or negative)
They have different numbers of protons and electrons.
Only electrons can move, and ions are made by gaining or losing electrons.

49. F1- O2- An Anion is… A negative ion. Has gained electrons.
Nonmetals can gain electrons.
Charge is written as a superscript on the right.
Has gained one electron (-ide is new ending = fluoride)
F1-
O2-
Gained two electrons (oxide)

50. K+ Ca2+ A Cation is… A positive ion. Formed by losing electrons.
More protons than electrons.
Metals can lose electrons K+
Has lost one electron (no name change for positive ions)
Ca2+
Has lost two electrons

51. Predicting Ionic Charges
Group 1A:
Lose 1 electron to form 1+ ions H+
Li+
Na+ K+
Rb+

52. Predicting Ionic Charges
Group 2:
Loses 2 electrons to form 2+ ions
Be2+
Mg2+
Ca2+
Sr2+
Ba2+

53. Predicting Ionic Charges
Group 13:
Loses 3
electrons to form
3+ ions
B3+
Al3+
Ga3+

54. Predicting Ionic Charges
Do they lose 4 electrons or gain 4 electrons?
Neither! Group 14 elements rarely form ions (they tend to share)
Group 14:

55. Predicting Ionic Charges
Nitride
Group 15:
Gains 3
electrons to form
3- ions
P3-
Phosphide
As3-
Arsenide

56. Predicting Ionic Charges
Oxide
Gains 2
electrons to form
2- ions
Group 16:
S2-
Sulfide
Se2-
Selenide

57. Predicting Ionic Charges
Gains 1 electron to form
1- ions
Group 17: F-
Fluoride
Br-
Bromide
Cl-
Chloride I-
Iodide

58. Predicting Ionic Charges
Stable noble gases do not form ions!
Group 18:

59. Predicting Ionic Charges
Many transition elements
have more than one possible oxidation state.
Note the use of Roman numerals to show charges
Iron (II) = Fe2+
Iron (III) = Fe3+

60. Naming cations
Two methods can clarify when more than one charge is possible:
Stock system – uses roman numerals in parenthesis to indicate the numerical value
Classical method – uses root word with suffixes (-ous, -ic)
Does not give true value

61. Naming cations We will use the Stock system.
Cation - if the charge is always the same (like in the main group of metals) just write the name of the metal.
Transition metals can have more than one type of charge.
Indicate their charge as a roman numeral in parenthesis after the name of the metal (Table 9.2, p.255)

62. Predicting Ionic Charges
Some of the post-transition elements also
have more than one possible oxidation state.
Tin (II) = Sn2+
Lead (II) = Pb2+
Tin (IV) = Sn4+
Lead (IV) = Pb 4+

63. Predicting Ionic Charges
Some transition elements have only one possible oxidation state, such as these three:
Silver = Ag+
Zinc = Zn2+
Cadmium = Cd2+

64. Do not need to use roman numerals for these:
Exceptions:
Some of the transition metals have only one ionic charge:
Do not need to use roman numerals for these:
Silver is always 1+ (Ag+)
Cadmium and Zinc are always 2+ (Cd2+ and Zn2+)

65. Practice by naming these:
Ca2+
Al3+
Fe3+
Fe2+
Pb2+
Li+

66. Write symbols for these:
Potassium ion
Magnesium ion
Copper (II) ion
Chromium (VI) ion
Barium ion
Mercury (II) ion

67. Anions are always the same charge
Naming Anions
Anions are always the same charge
Change the monatomic element ending to – ide
F- a Fluorine atom will become a Fluoride ion.

68. Practice by naming these:
Cl-
N3-
Br-
O2-
Ga3+

69. Write symbols for these:
Sulfide ion
Iodide ion
Phosphide ion
Strontium ion

70. Polyatomic ions are…
Groups of atoms that stay together and have an overall charge, and one name.
Usually end in –ate or -ite
Acetate: C2H3O2-
Nitrate: NO3-
Nitrite: NO2-
Permanganate: MnO4-
Hydroxide: OH- and Cyanide: CN-?

71. Know Table 9.3 on page 257 Phosphate: PO43- Sulfate: SO42-
Phosphite: PO33-
Ammonium: NH4+
Sulfate: SO42-
Sulfite: SO32-
Carbonate: CO32-
Chromate: CrO42-
Dichromate: Cr2O72-
(One of the few positive polyatomic ions)
If the polyatomic ion begins with H, then combine the word hydrogen with the other polyatomic ion present: H CO → HCO hydrogen + carbonate → hydrogen carbonate ion

72. Section 9.2 Naming and Writing Formulas for Ionic Compounds
OBJECTIVES:
Apply the rules for naming and writing formulas for binary ionic compounds.

73. Section 9.2 Naming and Writing Formulas for Ionic Compounds
OBJECTIVES:
Apply the rules for naming and writing formulas for compounds containing polyatomic ions.

74. Writing Ionic Compound Formulas
Example: Barium nitrate (note the 2 word name)
1. Write the formulas for the cation and anion, including CHARGES!
( )
Ba2+
NO3- 2
2. Check to see if charges are balanced.
Now balanced.
Not balanced!
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross method to balance subscripts.
= Ba(NO3)2

75. Writing Ionic Compound Formulas
Example: Ammonium sulfate (note the 2 word name)
( )
1. Write the formulas for the cation and anion, including CHARGES!
NH4+
SO42- 2
Now balanced.
2. Check to see if charges are balanced.
Not balanced!
= (NH4)2SO4
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross method to balance the subscripts.

76. Writing Ionic Compound Formulas
Example: Iron (III) chloride (note the 2 word name)
1. Write the formulas for the cation and anion, including CHARGES!
Fe3+
Cl- 3
2. Check to see if charges are balanced.
Not balanced!
Now balanced.
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross method to balance the subscripts.
= FeCl3

77. Writing Ionic Compound Formulas
Example: Aluminum sulfide (note the 2 word name)
1. Write the formulas for the cation and anion, including CHARGES!
Al3+
S2- 2 3
2. Check to see if charges are balanced.
Not balanced!
Now balanced.
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross method to balance the subscripts.
= Al2S3

78. Writing Ionic Compound Formulas
Example: Magnesium carbonate (note the 2 word name)
1. Write the formulas for the cation and anion, including CHARGES!
Mg2+
CO32-
2. Check to see if charges are balanced.
They are balanced!
= MgCO3

79. Writing Ionic Compound Formulas
Example: Zinc hydroxide (note the 2 word name)
1. Write the formulas for the cation and anion, including CHARGES!
( )
Zn2+
OH- 2
2. Check to see if charges are balanced.
Not balanced!
Now balanced.
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross method to balance the subscripts.
= Zn(OH)2

80. Writing Ionic Compound Formulas
Example: Aluminum phosphate (note the 2 word name)
1. Write the formulas for the cation and anion, including CHARGES!
Al3+
PO43-
2. Check to see if charges are balanced.
They ARE balanced!
= AlPO4

81. Naming Ionic Compounds
1. Name the cation first, then anion
2. Monoatomic cation = name of the element
Ca2+ = calcium ion
3. Monoatomic anion = root ide
Cl- = chloride
CaCl2 = calcium chloride

82. Naming Ionic Compounds
(Metals with multiple oxidation states)
some metals can form more than one charge (usually the transition metals)
use a Roman numeral in their name:
PbCl2 – use the anion to find the charge on the cation (chloride is always 1-)
Pb2+ is the lead (II) cation
PbCl2 = lead (II) chloride

83. Things to look for:
If cations have ( ), the number in parenthesis is their charge.
If anions end in -ide they are probably off the periodic table (Monoatomic)
If anion ends in -ate or –ite, then it is polyatomic

84. Practice by writing the formula or name as required…
Iron (II) Phosphate
Stannous Fluoride
Potassium Sulfide
Ammonium Chromate
MgSO4
FeCl3

85. Practice by writing the formula for the following:
Magnesium hydroxide
Iron (III) hydroxide
Zinc hydroxide

86. Hydrates
Some compounds contain H2O in their struc-ture. These compounds are called hydrates.
The H2O can usually be removed if heated.
A dot separates water: e.g. CuSO4•5H2O is copper(II) sulfate pentahydrate.
A greek prefix indicates the # of H2O groups.
sodium sulfate decahydrate
nickel(II) sulfate hexahydrate
Na2CO3•H2O
BaCl2•2H2O
Na2SO4•10H2O
NiSO4•6H2O
sodium carbonate monohydrate
barium chloride dihydrate

87. Prefixes 1 mono 2 di 3 tri 4 tetra 5 penta 6 hexa 7 hepta 8 octa
9 nona
10 deca

88. Hydrates Examples: I. Give the name of the following: 1. CuSO4  5H2O
2. MgCl2  6H2O
3. Na2SO4  10H2O
II. Write the formula for:
1. zinc chloride hexahydrate
2. calcium phosphate dihydrate
3. copper (I) chloride pentahydrate

89. Summary of Naming and Formula Writing
For naming, follow the flowchart- Figure 9.20, page 277
For writing formulas, follow the flowchart from Figure 9.22, page 278

90. Helpful to remember...
1. In an ionic compound, the net ionic charge is zero (criss-cross method)
2. An -ide ending generally indicates a binary compound
3. An -ite or -ate ending means there is a polyatomic ion that has oxygen
4. Prefixes generally mean molecular; they show the number of each atom

91. Helpful to remember...
5. A Roman numeral after the name of a cation is the ionic charge of the cation