1. An Experiment too Dangerous to do in class…

2. Unit 4 – Chemical Bonding

3. Two Kinds of Bonding
Ionic Bonding: Involves a transfer of electrons. One element loses electrons and the other gains electrons.
Covalent Bonding: (we’ll talk about this later on)

4. Ionic Bonds Forces of attraction bind oppositely charged ions together
Must be electrically neutral
Positive charges equal negative charges

5. Example – Sodium Chloride

6. What is the key to Ionic Bonding?
Octet Rule:
Atoms will gain or lose enough electrons to make them isoelectronic with a noble gas.
Except for He, there will be 8 electrons in their outermost s & p orbitals.

7. Ionic Bonds
Always a metal (positive charge or “cation”) + non-metal (negative charge or “anion”)
Look at valence electrons!

8. Electron Dot Structures
A way to show the number of valence electrons in an atom
Valence electrons must be labeled like this: X 4 6 3 1 2 7 5 8

9. The number of v.e. can be determined by the group1 2 13 14 15 16 17 18
v.e. 1 2 3 4 5 6 7 8
Dot strux • : •
• : •
• :
• • : : •
• •
: : Li • Be :
• • • : • N B C O F Ne

10. Valence electrons determine chemical reactivity
Elements in the same group behave the same!
Group # 1 2 13 14 15 16 17 18
Electron activity
Lose 1
Lose 2
Lose 3
Lose or gain 4
Gain 3
Gain 2
Gain 1 X
Oxidation # +1 +2 +3
+4/-4 -3 -2 -1 X

11. Ionic Bonds
Ionic bonds are not true “bonds”, which are shared electrons.
Cations and anions align to form a crystal lattice

12. To write the formula write the metal first, then the non metal…
Ionic Bonds
Use electron dot structures to predict the compounds that would form between Mg and Cl: Mg -1 : . • Cl +2
To write the formula write the metal first, then the non metal… : : . -1
MgCl2 Cl •

13. Try Another!
Use electron dot structures to predict the compounds that would form between Na and P:
Na P Na . . . : . . .
Na3P

14. Ionic Bonds
Ionic bonds are not true bonds, but ratios of cations to anions.
Formula unit shows lowest whole number ratio of ions in an ionic compound. Also called the Empirical Formula.
Examples:
NaCl – 1:1
MgCl2 – 1:2
AlBr3 – 1:3

15. Ionic Bonds When naming anions, change the suffix to “-ide” Examples:
Cl- : Chloride
Br- : Bromide
O-2 : Oxide
P-3 : Phosphide

16. How do you know when a bond is an Ionic Bond?
We can look at the differences in (EN) electronegativity of each atom involved in the bond.
If EN > 1.7, then it is ionic.

17. Properties of Ionic Compounds
They have a very high melting points.
They are solids and crystal shape at room temperature.
They conduct electric current when melted or dissolved in water (electrolytes)
Most of them are able dissolved in water

18. A) A two-dimensional cross section of a sodium chloride crystal is shown. (B) When struck by a hammer, the negatively-charged chloride ions are forced near each other. (C) The repulsive force causes the crystal to shatter.

20. (A) Distilled water does not conduct electricity
(A) Distilled water does not conduct electricity. (B) A solid ionic compound is also not very conductive. (C) A solution of an ionic compound dissolved in water conducts electricity well.

21. In an ionic solution, the A+ ions migrate toward the negative electrode, while the B- ions migrate toward the positive electrode.

28. Naming Compounds
A Love story…

29. Ionic Charges and the Octet Rule
All atoms want to have noble gas electron configuration (8 valence, except H)
To do this, they gain or lose electrons.
If the number of valence e- is smaller than 4, e- are lost, and a cation is formed.
If the number of valence e- is greater than 4, e- are gained, and an anion is formed.

30. Ionic charges of elements
For metals in groups 1, 2, and 13, refer to the group number to find the charge
Group 1 – What charge do they form?
1+ (Li+, Na+, K+)
Group 2 – What charge do they form?
2+ (Mg 2+, Ca2+)
Group 13 – What charge do they form?
3+ (Al 3+)

31. Ionic charges of elements
When looking at non-metals (anions), subtract to find the charge → group number minus 18
Group 17 – What charge do they form?
-1 (F1-, Cl1-, Br1-)
Group 16 – What charge do they form?
-2 (O2-, S2-, Se2-)

32. Forming Binary Ionic Compounds
Ionic compounds are electrically neutral
To accomplish this, we balance the charges with the number of ions.
The number of ions is represented with a subscript, below and to the right
H2O
This means I have two hydrogen ions.
I have one oxygen ion (the 1 is not written)

33. How do we do it? It’s the old “criss-cross”
Like a crossing route in football
NOT THIS!

34. Forming Binary Ionic Compounds
Let’s try to balance potassium and chlorine
Potassium Chlorine
K Cl1-
KCl

35. Forming Binary Ionic Compounds
Let’s try to balance calcium and bromine
Calcium Bromine
Ca Br1-
You need two bromine anions with a minus 1 charge to balance the one calcium with a +2 charge
CaBr2

36. Forming Binary ionic compounds
Let’s try aluminum and oxygen
Aluminum Oxygen
Al O2-
How many oxygen and aluminum atoms will you need to obtain a net ionic charge of zero?
Al2O3

37. Forming Binary Ionic Compounds
When balancing a compound, you must keep it to the simplest whole number ratio
Calcium Oxygen
Ca O2-
Ca2O2
Must simplify! For IONIC compounds, 2:2 ratio is the same as 1:1
CaO

38. Naming Binary Ionic Compounds
Simply say the first element, which will be a metal
Change the ending of the nonmetal to –ide
KCl→ Potassium Chloride
CaBr2 → Calcium Bromide
Al2O3 → Aluminum Oxide
CaO → Calcium Oxide

39. DO NOW Lithium Bromide Calcium Chloride Magnesium Oxide Sodium Sulfide
Strontium Nitride
LiBr
CaCl2
MgO
Na2S
Sr3N2

40. Naming Ionic compounds
For cations that have only one possible charge the name of the metal is used.
Examples: Group I metals (charge 1+), Group II metals (charge 2+), Aluminium (charge 3+), Zinc (charge 2+), Silver (charge 1+)

41. What about those transition metals?
Some transition metals have multiple oxidation numbers
Fe can lose 2 or 3 electrons
Fe2+ is called Iron (II)
Fe3+ is called Iron (III)
Iron (II) oxide would be: FeO
Iron (III) oxide would be: Fe2O3

42. Stock System
Use a Roman numeral to refer to the oxidation state of the cation, NOT how many cations are present!
Other transition metals with more than one oxidation number: Pb, Hg, Mn, Co, Cr, Cu, Ni, Sn.
MUST use a Roman numeral for any metal that has more than one oxidation #, any time you use it.
NEVER use a Roman numeral when there is only one oxidation #.
When in doubt, check your periodic table!

43. Naming Binary Ionic Compounds
For many transition metals, you must use Roman numerals when naming to give oxidation state of the cation (NOT how many cations are present).
Do a REVERSE CRISS-CROSS to find the oxidation number!
SnCl2 : Sn Cl +2 -1 2

44. Naming Binary Ionic Compounds
Tin (II) Chloride (tin two chloride)
SnCl4 is Tin (IV) Chloride (tin four chloride)
Cr2O3 is Chromium (III) Oxide (chromium three oxide)

45. Naming Binary Ionic Compounds
Be careful with oxygen!
Do a reverse criss-cross
Cu O
You might think this is Copper (I) Oxide…
But is oxygen -1?
What is the name of this compound?

46. Naming Binary Ionic Compounds
CuO is Copper (II) oxide
Cu2O is Copper (I) oxide
Remember to check if the cations and anions have the oxidation numbers that they should…

47. Try these… KBr MgI2 CaS ZnO PbF4 PbS MnO2 Potassium Bromide
Magnesium Iodide
Calcium Sulfide
Zinc Oxide
Lead (IV) Fluoride
Lead (II) Sulfide
Manganese (IV) Oxide

48. The old-fashioned way Old naming system:
Fe2+ is called ferrous (lower oxidation #)
Fe3+ is called ferric (higher oxidation #)
Name Co+2, Co+3, Pb+2, Pb+4
This system lead to confusion. How many oxidation states does Mn have?

49. DO NOW Calcium Chloride Hydrogen sulfide Iron (III) nitride
Chromium (III) oxide
Cs2Se
Cu3P
CaCl2
H2S
FeN
Cr2O3
Cesium Selenide
Copper (I) phosphide

51. Polyatomic Ions Polyatomic ions are ions composed of two or more atoms
Most end in “-ate” or “-ite”
Examples:
(NO3)- = Nitrate
(SO4)2- = Sulfate
(ClO2)- = Chlorite

52. Polyatomic Ions
Look on your polyatomic ion sheet to find names and charges of polyatomic ions. Learn them quickly!
Treat polyatomic ions like you would treat any other ion when writing the formula for the compound.

53. Polyatomic Ions
Whenever you need more than a single polyatomic ion to balance the formula, use parentheses.
If you don’t need more than a single polyatomic ion, don’t use parentheses!
When naming, simply state the name of the polyatomic ion used.

54. Write the formula for potassium sulfate. Potassium Sulfate K1+ (SO4)2-
Polyatomic Ions
Write the formula for potassium sulfate.
Potassium Sulfate
K (SO4)2-
K2SO4

55. Write the formula for magnesium hydroxide. Magnesium Hydroxide
Polyatomic Ions
Write the formula for magnesium hydroxide.
Magnesium Hydroxide
Mg (OH)1-
Mg(OH)2

56. Try these: NaNO3 Sodium Nitrate K2CO3 Potassium Carbonate (NH4)2S
Fe(OH)3
Pb3(PO4)2
Sodium Nitrate
Potassium Carbonate
Ammonium Sulfide
Iron (III) Hydroxide
Lead (II) Phosphate

57. Polyatomic Ions with Varying numbers of oxygen
Some Ions have very similar names. Eg. Nitrate, nitrite, etc…
Start with these “normal” number of oxygen
NO3- = Nitrate
ClO3- = Chlorate
SO42- = Sulfate
PO43- = Phosphate

58. Polyatomic Ions with Varying numbers of oxygen
If the ion has one more number of oxygen than normal, the name becomes
per-(stem)-ate
Example: NO3- is nitrate
NO4- is pernitrate
Notice – Charge remains the same

59. Polyatomic Ions with Varying numbers of oxygen
If the ion has one less oxygen than “normal”, the name becomes (stem)-ite
Example: SO42- is sulfate
SO32- is sulfite
Notice: Charge remains the same

60. Polyatomic Ions with Varying numbers of oxygen
If the ion has two less oxygen than “normal”, it becomes hypo-(stem)-ite
Example: ClO3- is chlorate
ClO- is hypochlorite
Notice: Charge is the same

61. Polyatomic Ions with Varying numbers of oxygen
NO4- = Pernitrate
NO3- = Nitrate
NO2- = Nitrite
NO- = Hyponitrite
This is all different from N3- = Nitride

62. Try these… NaNO3 Mg(ClO2)2 K2SO5 Fe3(PO4)2 Pb(ClO)2 Sodium Nitrate
Magnesium Chlorite
Potassium Persulfate
Iron (II) Phosphate
Lead (II) Hypochlorite