CHM 101 INTRODUCTORY CHEMISTRY II Lecturers: Dr. Abayomi, Dr. Oluseyi & Ms. Adetunde PHYSICAL CHEMISTRY - ONE

Electrovalent / Ionic bonding Lecture One CHEMICAL BONDING Electrovalent / Ionic bonding

Chemical Bonding The attractive chemical force which keeps the atoms in any molecule together is commonly described as a chemical bond. A chemical bond is the physical phenomenon of chemical substances being held together by attraction of atoms to each other through sharing as well as exchanging of electrons or electrostatic forces.

CHEMICAL BOND Two concepts - Valence Electrons - Octet Rule

VALENCE ELECTRONS Not all electrons in a given atom participate in bonding Only valence electrons are available for bonding (electrons in the outer most shell) For representative and noble-gas elements these electrons are always found in the s or p sub shells

VALENCE ELECTRONS Using electron configuration to determine the number of valence electrons C: 1s 2 2s 2 2p 2 O: 1s 2 2s 2 2p 4 Na: 1s 2 2s 2 2p 6 3s 1 Using electron-dot structure (Lewis symbol) to designate the number of valence electrons. (place first 4 dots separately on four sides and pair up as needed) ∙C∙∙C∙ :O∙:O∙ Na∙.....

Some atoms have too many electrons (one or two extra). These atoms like to give up their electrons.

Some atoms are really close to having a full shell. Those atoms go around looking for other atoms who want to give up an electron

Important facts about valence electrons Representative elements in the same group of the periodic table have the same number of valence electrons The number of valence electrons for representative elements is the same as the group number in the periodic table The maximum number of valence electrons for any given element is eight

OCTET RULE Electrons arranged with 8 valence electrons are more stable than all others The valence electron configuration of the noble gases are considered the most stable (all have 8 valence electrons; helium has 2) All noble gases have the outermost s and p subshell completely filled

OCTET RULE The noble gases are the most unreactive of all elements Atoms of many elements tend to acquire the 8 valence electron configuration through chemical reactions Atoms of elements tend to gain, lose, or share electrons to produce a noble-gas electron configuration This results in the formation of compounds This tendency is known as the OCTET RULE

When elements form compounds, they either lose, gain or share electrons so as to achieve stable electron configurations similar to the next higher or lower noble gas in the periodic table. This idea forms the basis of the electronic theory of bonding.

Loss of electron(s) by atoms Atoms of the first three elements in a period (Groups I, II & III) lose electrons from their outermost shell to form positively charged ions (Electropositive). Metals donate electrons to form positive ions These ions have an electron structure like Neon the previous noble gas.

Gain of electron(s) by atoms Elements at the end of the period, Groups VI & VII, gain electrons to form negatively charged ions. Nonmetals accept electrons to form negative ions These ions have the same electron structure like the next noble gas, Argon. Element Electron structureno. of e- inion formed electron structure of atomoutermost shell of ion S2,8,66S 2- 2,8 Cl2,8,77Cl - 2,8

Sharing of electrons by atoms Other elements across the period, Groups IV & V achieve electron structures similar to noble gases by sharing electrons with other atoms and not by electron transfer. (this will be discussed in the next lecture)

Electrovalent / Ionic bonding An ionic (or electrovalent) bond is formed when metals in Group I or II react with non metals in Groups VI & VII. When the reaction occurs, electrons are transferred from the metal to the non metal until the outer electron shells of the resulting ions are identical to those of a noble gas.

Na∙ + ∙Cl:[Na] + [:Cl:] - NaCl CaCl 2.. ∙Ca∙ +.. ∙Cl:.. [Ca] 2+ [:Cl:] -.. Formed by attraction between two oppositely charged ions as a result of the transfer of electron(s) from one atom to another atom(s) Formed between metal and nonmetal ions through electrostatic attraction

Energetics Removing an electron from Na(g) to form Na + (g) Na(g) → Na + (g) + e -  E = +496 kJ/mol Adding an electron to Cl(g) to form Cl - (g) Cl(g) + e - → Cl - (g)  E = -349 kJ/mol Attraction between the unlike charges draws ions together causing energy to be released Heat of formation of ionic substances is quite exothermic Na(s) + 1/2Cl 2 (g) → NaCl(s)  H f o = kJ IONIC BOND

Ionic compounds do not contain discrete molecules but ordered arrays of positive and negative ions (result of energy released) Formula unit that indicates combining ratio E.g. NaCl

The number of electrons lost by metal atoms equals the number gained by nonmetals so the compounds charge is NEUTRAL

Other examples of Ionic Compounds Formation of magnesium chloride,(MgCl 2 ) -The electronic configuration of magnesium (Atomic. No. 12) is 2,8,2. It has two electrons in its valence shell. -The electronic configuration of chlorine (Atomic. no. 17) is 2,8,7. It has seven valence electrons. In terms of the Lewis (electron dot) structures, one can write :

Formation of aluminium fluoride (AlF 3 ).

Properties of Ionic (or electrovalent) Bond An ionic bond is formed due to the coulombic attraction between the positively and negatively charged ions. An ionic bond is non-directional, i.e., the strength of interaction between two ions depend upon distance, but not on the direction. An ionic bond gets broken when the substance is dissolved in a polar solvent such as water, or when the substance is melted.

Factors Influencing the Formation of an Ionic Bond Low ionisation enthalpy of the metallic element which forms the cation. Large electron gain enthalpy (electron affinity) of the non-metallic element which forms the anion. Large lattice energy, i.e., the smaller size and higher charge of the ions.

Properties of ionic compounds High melting and boiling points Good conductors of electricity when molten but non - conductors when solid Readily soluble in water and in polar solvents but insoluble in non – polar solvents Crystals of ionic solids are hard and brittle