Ionic Bonding Isn t it Ionic? Don t you Think? Chemical Bonds Chemical bonds result from changes in the locations of the valence electrons of atoms. Because electronic structures are described by their positions on the periodic table, the bond-forming ability of an atom is predictable. Three basic types of bonds: Ionic Electrostatic attraction between ions Covalent Sharing of electrons Metallic Metal atoms bonded to several other atoms Gilbert Newton Lewis (1875 1946) From the electron configurations of atoms, valence electrons can be separated from the core electrons. Valence electrons are the electrons responsible for the formation of ions and covalent bonds. Even before the electronic structure of atoms was known, G.N. Lewis devised a very simple way of keeping track of valence electrons, known as electron dot symbols or Lewis Symbols. In Lewis Symbols, Valence electrons for atoms are shown using dots around the atomic symbol of the atom. X, X, X, X, X, etc.. The Octet Rule Most atoms seek the same electron configuration as the closest noble gas, which is very stable. Low Electron Affinity and High Ionization Energy H wants a He arrangement, Duet Rule P wants an Ar arrangement, Octet Rule To attain a noble gas arrangement, two generalizations apply: 1. A non-metal will take electrons from a metal to fill its valence shell, leaving the metal with a full valence shell in the next lowest main energy level.. Two non-metals will share electrons to fill their valence shells. Ionic Bonds are created by the physical attraction between two oppositely charged particles. Cation (+) Anion (-) Metals w/ relatively low Ionization energy Non-metals w/ relatively high electron affinity + - attraction 1
Ion Nomenclature Cations are named using the original name of the metal. Na + Sodium Anions are named using the nonmetal element name followed by an ide suffix. Cl - Chloride Ionic Compound Nomenclature Name cation first, with appropriate charge for transition metals in parenthesis, followed by the anion name with an ide suffix. Examples: K O Cu O CuO potassium oxide copper (I) oxide copper (II) oxide As we saw in the last chapter, it takes 495 kj/mol to remove electrons from sodium. We get 349 kj/mol back by giving electrons to chlorine. It generally requires more energy to remove an electron from a metal than is released by a gain of an electron by a non-metal So, the question is: Why is ΔH o f for ionic solids generally very exothermic? 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.
High Energy Arrangement Li Cl 1s s 1 3s 3p 5 + s 1 Almost [Ar] s Li + Cl 3P s 1s 3p 3s 1s 3s Low Energy Arrangement Li + Cl - s 3P 1s 3s The principle reason for ionic stability is due to the attraction between oppositely charged ions The physical attraction between ions in an ionic compound releases energy as the ions are drawn together. Once the ions reach the lowest energy possible, distance is minimum given electron repulsion, a crystal lattice is formed The strength of the ion attraction is described by the lattice energy between the ions. Lattice energy (aka lattice enthalpy) = The energy required to separate one mole of a solid ionic compound into its gaseous ions. Q1Q LE k r Where: Lattice Energy Q1Q LE k r k = proportionality constant dependent on structure of solid and on electron configuration of the ions. (8.99x10 9 Jm/C ) Q 1 and Q = charges on the ions r = the shortest distance between the centers of the cation and anion 3
Lattice Energy Lattice energy, then, increases with the charge on the ions. It also increases with decreasing size of ions. 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. Lattice energies can not be measured directly. However, by knowing the ionization energy and electron affinity of the ions involved, along with some basic thermodynamic information, it can be calculated. (see page 308) Max Born (188-1970) Born Haber Cycle Fritz Haber (1868-1934) In the Born-Haber cycle, we imagine that we break apart the elements into atoms (called enthalpies of atomization, described by ΔH o f of an element in the gas phase), ionize the atoms (cation ionization energies, anion electron affinity), combine the gaseous ions to form the ionic solid (Assumed lattice energy), then form the elements again from the ionic solid (negative enthalpy of formation). In the Born-Haber Cycle, only the lattice enthalpy, the enthalpy of the step where the ionic solid is formed from the gaseous ions is unknown. The sum of the enthalpy changes for a complete B-H Cycle is 0, because the enthalpy of the system must be the same at both the start and the finish. Therefore, using known enthalpy changes, the lattice energy can be found applying Hess s Law. Steps in a Born-Haber cycle 1. Start with elements in proportional amounts and atomize them using heats of formation [endothermic]. Form gaseous cations from the metal atoms using ionization energies [endothermic] 4
3. Form gaseous anions by using electron-gain energy (negative value of electron affinity) [usually exothermic]. 3. 4. Let the gas of ions form the solid compound (negative value of lattice energy) [exothermic] 5. Complete the cycle by forming elements from the ionic compound (negative value of formation enthalpy) [endothermic] 1. Begin Here 5. 4. 1.Using values from our text, let us calculate the lattice energy for sodium chloride together..using the Born-Haber cycle for Potassium chloride seen here, calculate its lattice enthalpy. 3.Draw the Born-Haber Cycle for magnesium oxide and label each enthalpy change by name and whether it is endo- or exothermic. Ionic characteristics The electrostatic forces holding ions in their rigid 3-D arrangement also accounts for their characteristics: Crystalline solid High melting point (mp) Brittle Cleavable 5