1. Chemical Bonds
Chemical bonds are strong electrostatic forces holding atoms or ions together, which are formed by the rearrangement (transfer or sharing) of outermost electrons
Atoms tend to form chemical bonds in such a way as to achieve the electronic configurations of the nearest noble gases (The Octet Rule )

2. Chemical Bonds: The Octet Rule: Octet Rule:
1. Atoms form bonds to produce the electron configuration of a noble gas, He (2), Ne(10),Ar(18),Kr(36),Xe(56) and Rn(86).
2.This is because the valence shell electronic configuration of noble gases is particularly stable.
3. 1s 2 for He(2) , and ns2np6 for Ne(10),Ar(18),Kr(36),Xe(56) and Rn(86).
4. For atoms it is of interest to achieve a valence shell configuration of 8 electrons corresponding to that of the nearest noble gas.
5. Atoms close to helium achieve a valence shell configuration of 2 electrons.
6. Atoms can form either ionic bonds(by transfer of electrons) or covalent bonds(by sharing of electrons) to satisfy the octet rule.

3. Lewis Dot Formulas for Representative Elements

4. Three types of chemical bonds
1. Ionic bond (electrovalent bond)
Formed by transfer of electrons

5. Three types of chemical bonds
Covalent bond
Formed by sharing of electrons

6. Three types of chemical bonds
3. Metallic bond
Electrostatic attraction between metallic cations and delocalized electrons (electrons that have no fixed positions)

7. What properties do most metallic compounds share?
Because the electrons in a metal can move freely, most metals are good conductors of electric current.
When you turn on a lamp, valence electrons move through the copper wire that connects the light bulb to the electrical outlet.
The valence electrons in the copper atoms are free to move because they are not connected to any one atom.

8. What properties do most metallic compounds share?
Due to their free-moving electrons, metals have two properties that allow them to be reshaped.
Malleability is the ability to be hammered into sheets.
Ductility is the ability to be formed into long, thin wires.

9. The Ionic Bond
An ionic bond is a chemical bond formed by the electrostatic attraction between positive and negative ions
This type of bond involves the transfer of electrons from one atom (usually a metal from Group I or II) to another (usually a nonmetal from Group VII or the top of Group VI)
Metal with low Ionization Energy (IE) absorbs energy and loses one or two electrons (IE is positive)
Nonmetal with high Electron Affinity (EA) loses energy gaining one or two electrons (EA is negative)
The number of electrons lost or gained by an atom is determined by its need to be “isoelectronic” or the “same as” a noble gas electron configuration

10. The Ionic Bond
An ionic bond is the electrostatic force that holds ions together in an ionic compound. Consider, for example, the reaction between lithium and fluorine to form lithium fluoride

11. Bonding in Ionic Compounds
Such noble gas configurations and the corresponding ions are particularly stable
The atom that loses the electron becomes a positively charged ion, i.e., a cation (positive)
The atom that gains the electron becomes a negatively charged ion, i.e., an anion (negative)

13. The Ionic Bond
Because electrostatic attraction is not directional in the same way as is covalent bonding, there are many more possible structural types. However, in the solid state, all ionic structures are based on infinite lattices of cations and anions.
NaCl
CsCl
Zinc Blende F F Ca
Fluorite
Wurtzite

14. Lattice Energy
Elements are likely to form ionic compounds based on ionization energy and electron affinity, but how do we evaluate the stability of an ionic compound?
Ionization energy and electron affinity are defined for processes occurring in the gas phase.
The solid state is a very different environment because each cation in a solid is surrounded by a specific number of anions, and vice versa.
A quantitative measure of the stability of any ionic solid is its lattice energy, defined as the energy required to completely separate one mole of a solid ionic compound into gaseous ions.

15. 1. theoretical calculation using an ionic model, Or
Lattice Energy
Direct determination of lattice energy by experiment is very difficult, but it can be obtained from
1. theoretical calculation using an ionic model, Or
2. experimental results indirectly with the use of a Born-Haber cycle.

16. E.g.: Na+(g) + Cl-(g)  NaCl(s) Uo = -788 kJ/mol
Lattice Energy
The energy that holds the arrangement of ions together is called the lattice energy, Uo, and this may be determined experimentally or calculated.
Uo is a measure of the energy released as the gas phase ions are assembled into a crystalline lattice. A lattice energy must always be exothermic.
E.g.: Na+(g) + Cl-(g)  NaCl(s) Uo = -788 kJ/mol
Lattice energies are determined experimentally using a Born-Haber cycle such as this one for NaCl. This approach is based on Hess’ law and can be used to determine the unknown lattice energy from known thermodynamic values.
NaCl(s)
Na(s)  Na(g)  Na+(g)
½ Cl2(g)  Cl(g)  Cl-(g)
DH°ea
DH°d
DH°ie
DH°sub
DH°f
Lattice Energy, Uo

17. DH°f = DH°sub + DH°ie + 1/2 DH°d + DH°ea + Uo
Born-Haber cycle
NaCl(s)
Na(s)  Na(g)  Na+(g)
½ Cl2(g)  Cl(g)  Cl-(g)
DH°ea
DH°d
DH°ie
DH°sub
DH°f
Lattice Energy, Uo
DH°f = DH°sub + DH°ie + 1/2 DH°d + DH°ea + Uo
-411 = /2 (242) + (-349) + Uo
Uo = -788 kJ/mol
You must use the correct stoichiometry and signs to obtain the correct lattice energy.

18. Born-Haber Cycle
A series of hypothetical steps and their enthalpy changes needed to convert elements to an ionic compound and devised to calculate the lattice energy.
Using Hess’s law as a means to calculate the formation of ionic compounds

19. Born-Haber Cycle Steps
Elements (standard state) converted into gaseous atoms
Losing or gaining electrons to form cations and anions
Combining gaseous anions and cations to form a solid ionic compound