1. Chapter 6  Chemical Bonds
6.1 Ionic Bonding & Naming
6.2 Covalent Bonding & Naming
Metallic Bonds

2. 6.1  Ionic Bonding

3. Stable Electron Configurations
When the highest occupied energy level of an atom is filled with electrons, the atom is STABLE and NOT LIKELY TO REACT.
ex: the Noble gases (8 valence e-)
Chemical properties of an element depend on the number of valence e-

5. Representing Electron Configurations
Use electron dot diagram or a Lewis Structure to focus on the valence e-
Model of an atom in which each dot represents a valence e- (symbol in the center represents the nucleus and all the other e- in the atom)
ex:

7. Ionic Bonds
Elements that do not have complete sets of valence e- tend to react
By reacting, they achieve e- configurations similar to those of noble gases (stable configurations)
Achieve stable e- configurations through the transfer of e- between atoms
ex: NaCl

8. Formation of Ions
Ion: an atom that has a net positive or negative electric charge
Atoms that have the same number of protons as electrons have NO charge. They are considered neutral!
Ions form when an atom gains or loses an e- (number of protons is no longer equal to the number of e-)
Charge is represented by a plus or a minus sign
Atoms will transfer electrons to achieve a stable electron configuration. This is what creates ions.

10. Common Ions

11. Formation of Ions
Question: When is an atom considered to have a stable electron configuration?
Answer: When the highest occupied energy level is filled with 8 valance electrons!!!

12. Formation of Ions Cations are pawsitive
Cation= ion with a positive charge (has lost e-)
Cations are pawsitive
Anion= ion with a negative charge (has gained e-)

13. Some Humor… 
A sodium ion walks into a class. It says to the teacher, “I’ve lost an electron.”
The teachers says, “Are you sure?”
The sodium ion says, “I’m positive.”

14. Formation of Ionic Bonds
Particle with negative charge attracts particle with positive charge
Chemical Bond: force that holds atoms or ions together as a unit
Ionic Bond: force that holds cations and anions together
Forms when e- are transferred from one atom to another

15. Ionic Bond Formation

16. Ionization Energy
Ionization Energy- amount of energy used to remove an electron
It allows electrons to overcome the attraction of the protons in the nucleus 
Trend- ionization energy increases from left to right and decreases from top to bottom
It takes more energy to remove an electron from a nonmetal than a metal
Example: It takes more energy to remove an electron from Na than K.

17. Ionization Energy

19. Ionization Energy
Question: how does ionization energy relate to the reactivity of the alkali metals?
Answer: low ionization energy = easier to remove electrons.
The easier it is to remove an electron the more reactive an atom is.

20. DRAWING IONIC COMPOUNDS
FIRST, YOU HAVE TO THINK...!
HOW MANY ATOMS OF CHLORINE DOES SODIUM NEED TO REACT WITH TO BE STABLE??? 
THEN FOLLOW THE DIRECTIONS BELOW...
Draw electrons (dots)
Move e- to most stable configuration
Draw arrows to show new location
Write the new charge of the ions formed.
Draw ionic bond for sodium & chlorine
Draw ionic bond for magnesium and chlorine

21. Chemical Formulas
Chemical formula: a notation that shows what elements a compound contains and the ratio of the atoms or ions of those elements in the compound
Examples: H20, NaCl, C6H1206
Identify the number of each type of atom in the space below:
H20  How many hydrogen? Oxygen?
NaCl  How many sodium? Chlorine?
C6H1206  How many carbon? Hydrogen? Oxygen?

22. Properties of Ionic Compounds
Properties of ionic compounds can be explained by the strong attractions among ions within a crystal lattice:
High melting points
Poor conductor of electric current as a solid
Good conductor of electric current when melted
Will shatter

23. Example - NaCl

25. Naming Binary Ionic Compounds
Binary compound- a compound made of only two elements
Naming binary compounds is easy!!!  
The first part is the name of the cation (+) metal  with NO change to the name
Sodium  stays sodium
The second part is the name of the anion (-) nonmetal  use part of the nonmetal name with the suffix “IDE” at the end
Chlorine  becomes chloride
EXAMPLE  Sodium Chloride
Refer to Figure 16 on page 171 for a list of the names and charges of the 8 common anions. 

26. Writing Formulas Ionic Compounds
STEPS:
Place the symbol of the cation first!
Follow this symbol with the symbol for the anion
Use subscripts to show the ratio of the ions in the compound
Because all compounds are neutral, the total charges on the cations and anions must add up to zero!

27. EXAMPLE: What is the name of the compound that is formed when sulfur and sodium react? 
Identify the cation and anion. Na = Sodium, S = Sulfur
How many electrons does sodium lose to become stable? 1
What is the overall charge on sodium after it loses electrons? +1 so, Na+
How many electrons does sulfur accept to become stable? 2
What is the overall charge on sulfur after it gains these electrons? –2 so, S2-
To have a neutral compound of sodium and sulfur ions, what ratio of ions will create an overall charge of zero? TWO sodiums for every ONE sulfur
Use subscripts to show the appropriate ratio of atoms: Na2S

28. 6.2 Covalent Bonds

29. Review Question: do nonmetals have high or low ionization energies?
 Answer: high
 Question: so, if it is difficult to remove, or transfer, electrons between nonmetals, then how do they form bonds?
 Answer: they have to share their electrons!

30. Covalent Bond
Covalent Bonds- a chemical bond in which 2 atoms share a pair of valance electrons.
When two atoms share a pair of electrons, the bond is called a single bond.
Refer to Figure 9 on page 166 to see how to model covalent bonds.
Molecule- atoms joined by covalent bonds

33. Covalent Bonds Question: What keeps the atoms in molecules together?
Answer: The attraction between the shared electrons and the protons in the nucleus hold them together.

35. The oxygen atom in water and the nitrogen atom in ammonia are each surrounded by eight electrons as a result of sharing electrons with hydrogen atoms.

36. IONIC/COVALENT BONDS Ionic bonds Metal to nonmetal
Electrons are transferred
Covalent bonds
Nonmetal to nonmetal
Electrons are shared.

37. Naming Molecular (Covalent) Compounds
Rules to naming:
Most metallic element appears first in the name
These are further to the left on the table
If in same group – more metallic is found closer to the bottom
Name of second element is changed to end in the suffix – “ide” (as in carbon dioxide)
Example: two compounds contain nitrogen and oxygen  N2O4 & NO2
The names reflect the actual number of atoms – using Greek prefixes
Dinitrogen tetraoxide & mononitrogen dixoxide (however, prefix mono is often not used for first element  nitrogen dioxide)

38. Greek Prefixes

39. Writing Molecular (Covalent) Formulas
Rules to writing a molecular formula:
Write the symbols for the elements in the order they appear in the name
Prefixes indicate the number of atoms in each element in the molecule
Prefixes appear as subscripts in the formulas
If there is NO prefix for an element – there is only one atom
Write formula for diphosphorous tetrafluoride:
Di- indicates two phosphorous & tetra- indicates four fluorine
Formula would be P2F4

40. Now, for one last type of bond…and I don't mean James Bond....
Waaayyy more daring and exciting are the infamous...
METALLIC BONDS

41. Metallic Bonds
REMEMBER: metals achieve stable configurations by losing electrons
Question: In most cases, which class of elements receive these electrons?
 Answer: nonmetals 

42. Metallic Bonds
Question: What happens when there are no nonmetal atoms available to accept these electrons? 
A: There is a way for metal atoms to lose and gain electrons at the SAME TIME!

44. Metallic Bonding
In a metal, the valance electrons are free to move around the atom.  
This movement basically changes the metal atom into a cation surrounded by a pool of shared electrons.
A metallic bond is the attraction between a metal cation and the shared electrons that surround it.  
The cations in a metal form a lattice that is held in place by strong metallic bonds between the cations and the surrounding valance electrons.

46. Not all are created equal…
However, not all metals are created equally!
The metallic bonds in some metals are stronger than in other metals. 
The more valance electrons an atom can contribute to the shared pool, the stronger the metallic bonds will be.

48. So what happens if we add more valence electrons?

49. Let’s look at the periodic table a moment….
Question: how many valence electrons do the alkali metals contribute?
 Answer: 1
So, the bonds in alkali metals are relatively weak. 
Because the bonds are weak, alkali metals, like sodium and potassium, are soft enough to cut with a knife! (Plus, they have low melting points)!

50. Explaining Properties of Metals
So what does this mean for the other properties of metals???
The mobility of electrons within a metal lattice explains some of the properties of metals.
 Question: what are two of the most important properties of metals?
 Answer: malleability and conductivity

51. Explaining Properties of Metals
For conductivity, remember, it’s the ability to allow a flow of electric charge…
Metals have a built-in supply of charged particles that flow from one location to another (the pool of shared electrons )
An electric current can be carried through a metal by the free flow of the shared electrons! 

52. Explaining Properties of Metals
For malleability, remember it is the ability to be hammered without shattering… 
the lattice (structure) in a metal is quite flexible. If a metal is struck, the ions simply shift to new positions, but they are still connected by metallic bonds.
Since the electrons still surround the ions, the metallic bond between the ions and the electrons is not broken.  
This also explains why metals are ductile (can be drawn into wires without breaking)!

53. If a metal is struck, the ions simply shift to new positions, but they are still connected by metallic bonds