1. Unit 4 Notes
Chemistry Mr. Nelson
2009

2. Chemical Bonds Three basic types of bonds Ionic Covalent Metallic
Electrostatic attraction between ions
Covalent
Sharing of electrons
Metallic
Metal atoms bonded to several other atoms

3. Why do atoms bond? Why DON’T some atoms bond?
The noble gases – why?
Why do other atoms bond, then?
They are more chemically stable when bonded

4. How do atoms bond? The octet rule
The octet rule, or rule of eight, says that an atom will strive for a full s and p subshell
Atoms will either lose or gain electrons to get 8 in the outer shell
NOTE: when an atom loses or gains electrons, it’s nucleus remains the same – only the outer electron shell has changed!!!

5. Bonding and energy changes
Energy is the ability to do work
Stability is a measure of inability to do work
So, the lower the energy, the more stable something is!
When atoms bond, the process favors stability (lower energy). Things will never go from a stable to an unstable state on their own!

6. Electrons, bonding, and IONS
When they do this, they get a CHARGE, because protons (+) and electrons (-) are no longer equal. They are now IONS
Positive and negative IONS come together and balance each other out in IONIC BONDS.

7. Cations and Anions
Remember:
+ + +
An “antion”
A “plussy cat”

8. Ionic Bonding
Sodium wants to GIVE an electron, Chlorine wants to GET an electron.

9. Ionic Bonding
The low ionization energy of sodium and the high electronegativity of chlorine is one reason this works so well.

10. Energetics of Ionic Bonding
As we saw in the last chapter, it takes 495 kJ/mol to remove electrons from sodium.

11. Energetics of Ionic Bonding
But these numbers don’t explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so exothermic!

12. Energetics of Ionic Bonding
There must be a third piece to the puzzle.
What is as yet unaccounted for is the electrostatic attraction between the newly-formed sodium cation and chloride anion.

13. Lattice Energy This third piece of the puzzle is the lattice energy:
The energy required to completely separate a mole of a solid ionic compound into its gaseous ions.

14. Lattice Energy
Lattice energy, then, increases with the charge on the ions.
It also increases with decreasing size of ions.

15. Energetics of Ionic Bonding
By accounting for all three energies (ionization energy, eletronegativity, and lattice energy), we can get a good idea of the energetics involved in such a process.

16. Energetics of Ionic Bonding
These phenomena also helps explain the “octet rule.”
Metals, for instance, tend to stop losing electrons once they attain a noble gas configuration because energy would be expended that cannot be overcome by lattice energies.

17. Naming ions Monatomic ions = Polyatomic ions = Naming monatomic ions
One atom ions
Polyatomic ions =
Many atom ions
Naming monatomic ions
To name positive ions, just add the word “ion”
To name negative ions, drop the last part of the word, and add “-ide ion”

18. Naming monatomic ions Rubidium loses an electron to become Rb+
Rubidium ion
Calcium loses two electrons to become Ca2+
Calcium ion
Chlorine gains an electron to become Cl-
Chloride ion
Oxygen gains two electrons to become O2-
Oxide ion
Nitrogen loses three electrons to become N3-
Nitride ion

19. Compounds made of two monatomic ions
These are called BINARY COMPOUNDS
You always put the positive part first and the negative part last: Na+ + Cl-  NaCl
Names = name of the positive ion + name of the negative ion: Sodium Chloride

20. Examples: Name the following

21. Write the formulas of the following:

22. Back to ions: Writing Ionic Formulas
The nomenclature (naming system):
Write the symbols for the ions side by side. Write the cation first.
Al3+ O2-
Find the lowest common multiple that will make the charges on each ion cancel out
Al O2-
Check the subscripts for the lowest whole number ratio of ions. Then write the formula. Al2O3

23. d-block naming Write the electron configuration for Iron.
Predict the oxidation number

24. d-block The d-block (yo) has its own rules
Metals in the d-block have variable charges
All d-block metals must get a ROMAN NUMERAL to indicate the charge
EXAMPLE: copper (I) chloride is made of Cu1+ and Cl-
EXAMPLE: copper (II) chloride is made of Cu2+ and Cl-
Don’t use roman numerals if you don’t have to

25. Examples
Write the formulas for
Tin(II) iodide
Cobalt(III) chloride

26. Working backward
If you are given the formula you need to calculate the charge of the d-block metal.
Assume the anion did not change its charge (they are very consistent)
Example: FeO, to write the name we need the charge of iron.

27. A few more examples
PbS2
MnBr3
Cu3P2

28. More d-block (old school)
Roman numerals are the ‘new’ way.
The ‘old’ way is based on Latin names
Two endings ic
ous
ic is for the highest charge
ous is for the lower charge
Example
Ferric = Iron(III)
Ferrous = Iron(II)

29. Old School The only three I expect you to know are: Tin (Sn)
Stannic = Tin(IV)
Stannous = Tin(II)
Copper (Cu)
Cupric = Copper(II)
Cuprous = Copper(I)

30. Polyatomic ions
When two or more ions are clumped together it is a polyatomic ions.
They usually end with –ates or -ites

31. Nomenclature of Oxyanions
They are not standard!
Example Sulfate vs Phosphate
Nomenclature examples
Perchlorate
Chlorate Nitrate
Chlorite Nitrite
hypochlorite

32. Writing formulas for compounds with polyatomic ions
Polyatomic ions should ALWAYS be treated like BOY BAND. Don’t ever break it up!
If you need more than one polyatomic ion to balance a charge, use PARENTHESES ( )

33. Polyatomic ions Naming compounds that contain polyatomic ions:
The steps are the same:
the name of the first ion + the name of the second:
NH4+ = ammonium ion (polyatomic)
Cl- = chloride ion (monatomic)
NH4Cl = ammonium chloride

34. Write the formulas for:
Example
Write the formulas for:
potassium perchlorate
tin(IV) sulfate
Iron(II) chromate
ammonium sulfate

35. Ionic vs. Metallic bonds
In an IONIC BOND, the electrons of one atom (that wants to lose electrons) are donated to the electrons of another atom (that wants to gain electrons). The charges on each ion balance each other out and equal ZERO.
In a METALLIC BOND, all the atoms are the same (all copper, for example) and the electrons don’t belong to any one atom. They move around a lot – that’s why electricity is conducted.

36. Metallic Bonds A “sea” of mobile outer electrons.
Low ionization energies means the atoms don’t hold electrons well.