Ionic Bonding

Types of Bonding• There are 2 bond types : ionic and

covalent-3rd. type of bonding is metallic!• In ionic bonding one atom has a stronger

attraction for electrons than the other, and “steals” an electron from a second atom

• In covalent bonding the attraction for electrons is similar for two atoms. They share their electrons to obtain an octet.

• MgO (ionic), CaCl2 (ionic),SO2(covalent),

PbCl2(ionic), CCl4(covalent), CH4 (covalent)

Ionic bondingIonic bonding• Ionic bonding involves 3 steps (3 energies)• 1) loss of an electron(s) by one element, 2)

gain of electron(s) by a second element, 3) attraction between positive and negative

Cl–

Na Cl

e–1) 2)

Na+

• Ionic bonding (stealing/transfer of electrons) can be represented in three different ways

• Li + Cl [Li]+[Cl]–

3p+

4n02e-1e-

17p+18n0

7e- 8e- 2e-

1e-

3p+

4n02e- 17p+

18n08e-8e-2e

Li Cl Cl[Li]+

Mg + O [Mg]2+[O]2–

12p+

12n02e- 8e- 2e-

[ O ]2–[Mg]2+

6e- 2e-

8n0

8p+

8e- 2e-

8n0

8p+ 12p+

12n02e- 8e-

OMg

Properties of Ionic Compounds

• Crystalline solids at room temperature • Have higher melting points and boiling points

compared to covalent compounds • Conduct electrical current in molten or

solution state • Are extremely polar bonds • Most are soluble in water but not soluble in

non-polar solvents

Ionic Bonds

• The nature of the bond• The sodium ions and chloride ions are held

together by the strong electrostatic attractions between the positive and negative charges.

• Coordination number: is the number of ions of opposite charge that surrounds the ion in a crystal figure.

Coordination numbers cont.

• Each Sodium atom is surrounded by 6 Chlorine ions, therefore Na+ has a coordination number of 6 and each Chlorine ion is surrounded by 6 sodium ion and also has a coordination number of 6. See page 198 in book for better detail!

Metallic Bonding• The metals are the most numerous of the

elements. About 80 of the 100 or so elements are metals. You know from your own experience something about how metallic atoms bond together. You know that metals have substance and are not easily torn apart. They are ductile and malleable. That means they can be drawn into shapes, like the wire for a paper clip, and their shape can be changed. They conduct heat and electricity. They can be mixed to form alloys. How is it that metallic bonding allows metals to do all these things?

• The nature of metals and metallic atoms is that they have loosely held electrons that can be taken away fairly easily. Let's use this idea to create a model of metallic bonding to help us explain these properties. I will use potassium as an example. Its valence electron can be represented by a dot. The valence electron is only loosely held and can move to the next atom fairly easily. Each atom has a valence electron nearby but who knows which one belongs to which atom. It doesn't matter as long as there is one nearby.

• To emphasize that the valence electron is very loosely held, we can separate it from the rest of the atom and write it as "K+ and e-" rather than "K·". Packed in a cluster they look like this. These electrons are more or less free to move from one atom to another. Chemists often describe metals as consisting of metal ions floating in a sea of electrons.

• The mutual attraction between all these positive and negative charges bonds them all together. Atom to electron to atom to electron and so forth. We have an array of atoms bonded to one another, that is, a network. The network in this paper clip (which of course is not made out of potassium), has a vast number of atoms that are all bonded to one another. This paper clip has on the order of 1022 atoms. The network of metallic bonding holds that entire chunk of metal together. Each metallic bond gives strength and the network extends that strength over the entire chunk of metal.

How can this particular model of metallic bonding be used to explain the properties of metals (such as electrical conductivity, malleability, and thermal conductivity)? The key is in the loosely held electrons spread around and between all the metal atoms, or metal ions. These electrons can move easily from one place to another, allowing for good electrical conductivity.

• To a limited extent, the atoms can also move from one place to another and still remain in contact with and bonded to the other atoms and electrons around them. Take the beads to represent the atoms. If these are shifted in position, the atoms still remain pretty much in contact with one another. Although the external shape of the metal is changing, the internal pattern is pretty much the same. Thus the shape of metals can be changed.

Alloys• If all of the atoms in a piece of metal are the

same, we have a pure element. But what about metals like brass and steel that are not pure elements? What happens if we add different elements, different kinds of atoms, heat them until they melt and mix them? When we do that, we get a mixture of atoms that we call an alloy. The ratio of atoms of one kind to another is not fixed. The composition of the mixture will vary depending on how many of each kind of atom (or how much of each kind of element) happens to be available.

• If the atoms are about the same size, we get a substitution alloy. One kind of atom just takes the place of another kind of atom. The network pattern remains more or less the same.

• If the added atoms are much smaller than the atoms in the network, like the carbon atoms added to iron to make steel, they can fit into the holes between the layers of atoms in the network. When this happens we call it an interstitial alloy.

The presence of different kinds of atoms alters the bonding and changes the properties of the metal. For example, steel is stronger than iron; bronze and brass are both stronger than copper.

• In metallic bonding, metal atoms form a close-packed, regular arrangement. The atoms lose their outer-shell electrons to become positive ions. The outer electrons become a ‘sea’ of mobile electrons surrounding a lattice of positive ions. The lattice is held together by the strong attractive forces between the mobile electrons and the positive ions.

Covalent bondingCovalent bonding

• Thus far we have looked at when atoms bond due to the transfer of electrons

• An ionic bond forms when an atom has a greater attraction for e–s than a second atom

• However, if two atoms have approximately the same pull on electrons, they share the electrons (forming a “covalent” bond)