Chemical Bonds

1. Chemical Bonds Chapter 8

2. A Chemical Bond is a link between atoms. • An Ionic Bond is the electrical attraction between the opposite charges of cations and anions. • A Lewis symbol consists of the chemical symbol of an element and a dot for each of its valence electrons.Example He: • The formation of ionic bonds is represented in terms of Lewis symbols by the loss or gain of electrons until both species have reached an octet of electrons.

3. The tendency to form cations two units lower in charge than expected from the group number is called the inert pair effect.Example consider the group 13 elements Al and Indium. Al forms Al+3 but Indium forms In+1 and In+3. Group 14 has Pb that forms Pb+2 oxide when heated and tin forms tin (IV)oxide when heated.

4. Lattice enthalpy • A measure of the attraction between ions is the lattice enthalpy, the enthalpy change per mole of formula units when a solid is broken up into a gas of widely separated ions. All lattice enthalpies are positive. Heat equal to the lattice enthalpy is released when the solid lattice forms from gaseous ions.

5. Lattice enthalpies and Born-Haber Cycle • The lattice enthalpy for a particular ionic compound is defined as ∆H for the process . • This cannot be measured directly as it is not possible to get this reaction to happen on its own, without many other reactions happening around it. • Separate ions cannot be brought together in this way. However, we can use other pieces of experimental data to calculate this result. This is known as a Born-Haber Cycle. • In a Born-Haber cycle, we imagine that we break apart the elements into atoms, ionize the atoms, combine the gaseous ions to form the ionic solid, then form the elements again from the ionic solid. Only the lattice enthalpy, the enthalpy of the step in which the ionic solid is formed from the gaseous ions, is unknown. • The overall energy change for a complete Born Haber cycle is 0.

6. Start with the elements in the proportions in which they appear in the compound and atomize them. Write the corresponding enthalpies of formation of the gas phase atoms next to the upward pointing arrows. • Form gaseous cations from the metal atoms. This step requires the ionization energy of the metal and the sum of the first and higher ionization energies. The arrow points upwards.

7. Form gaseous anions from the nonmetal atoms. The enthalpy change of this step is called the electron gain enthalpy ∆Heg°. It is the negative of the electron affinity. If the electron affinity is positive, the electron gain enthalpy is negative and the corresponding arrow points downwards as the energy is released. If the electron affinity is negative then the electron gain enthalpy is positive and the arrow points upward. • Let the gas of ions form a solid compound. This step is the reverse of the formation of ions from the solid so its enthalpy change is the negative of the lattice enthalpy, -∆Hl. Denote it by an arrow pointing downward, since the formation of the solid is exothermic. This is the unknown value in the cycle. • Complete the cycle with the arrow from the compound to the element; the enthalpy change in this step is the negative of the enthalpy of the formation of the compound from its elements, ∆Hf°. The arrow points up if the ∆Hf° is negative, down if it is positive.

8. Finally calculate ∆Hl, from the fact that the sum of all the enthalpy changes for the complete cycle is 0.

9. Covalent Bonds • A covalent bond is a pair of electrons shared between two atoms.

10. Octet rule and Lewis structure • In covalent bonds, atoms share electrons to reach a noble gas configuration. Lewis called this the octet rule.The valence of an element is the number of covalent bonds an atom of the element forms.Consider molecular hydrogen, H2. Each atom completes its helium like duplet by sharing its electron with the other:

11. Class Practice • Write the Lewis structure for the compound HBr and state how many lone pairs each atom in the compound possesses?

12. Lewis structure for polyatomic species • To write the Lewis structure for polyatomic species we count the valence electron from all the atoms in the molecule. For example for methane there are 8 valence electrons. • The next step is to arrange the dots representing the electrons so that the carbon atom has an octet and each hydrogen atom has a duplet. • A single shared pair of electrons is called a single bond. • Atoms can share two or more electron pairs. • Two shared electron pairs form a double bond, and three shared electron pairs form a triple bond.

13. Bonds

14. Home work • Page 358 • 8.48,8.52 • Write a Lewis structure for the amide ion NH2− .

15. Resonance • In some Lewis structure the multiple bonds can be written in several equivalent locations. Consider the nitrate ion,NO3−.

16. The three Lewis structures shown differ only in the position of the double bond. All the three structures are valid. The bonds in the nitrate ion have a character intermediate between a pure single bond and a pure double bond.We present this as a blend of all three Lewis structure.

17. Formal charge • The formal charge gives an indication of the extent to which atoms have gained or lost electrons in the process of covalent bond formation. Structures with lowest formal charges are likely to have the lowest energy. • Formal charge=number of valence electron in the free atom−(number of electrons present as lone pairs−½(number of electrons shared in bonds) • =V−(L+½S)

18. Class practice • Write three plausible structures with different atomic arrangements for the cyanate ion, NCO− , and suggest which one is likely to be the most plausible structure.

19. Radical and Biradical • A radical is a species with an unpaired electron; a biradical has two unpaired electron

20. Lewis acids and bases • When a coordinate covalent bond forms, one species provides a lone pair and the other species accepts it.The species that provides the lone pair is called as Lewis base and the species that accepts it is called as Lewis acid. • To introduce this new class of reactions, lets investigate the molecular structure of the colorless gas boron trifluoride,BF3. The Lewis structure indicates that the boron atom has an incomplete octet: its valence shell consists of only six electrons. The molecule could complete its octet by sharing more electrons with fluorine, but fluorine has such a high ionization energy that this arrangement is not likely.

21. This boron octet can be completed if another atom or ion with a lone paired electrons forms a bond by providing the needed pair of electrons. Example BF4−( tetrafluroborate anion) forms when boron trifluoride is passed over a metal fluoride. Now all the fluorine atoms have their normal valence of 1and the boron atom has an octet.

22. Ionic versus Covalent bonds. • Ionic and covalent are terms used to describe two extremes of chemical bonds. When describing bonds with non metals covalent bonds are good models and when a metal is involved we say that it is an ionic bond. A covalent bond acquires some ionic character if one atom has a greater electron withdrawing power than the other atom. This electron withdrawing power is called as electronegativity.When a chemical bond forms between two atoms the atom with a higher electronegativity pulls the atom with a lower electronegativity.

23. Electronegativity is a measure of the electron pulling power of an atom on an electron pair in a molecule. Compounds composed of elements with a large difference in electronegativity (≥2) tend to have significant ionic character in their bonding.

24. Class Practice • In which of the following compounds do the bonds have greater ionic character; NH₃ or NO₂? • Indicate which atom in each compound has the partial negative charge.

25. Home work • Page359 • 8.84,8.88 • Page 357 • 8.36,8.38