Chapter 9 Chemical Bonding CHM 1045 Bushra Javed

1. Chapter 9 Chemical Bonding CHM 1045 BushraJaved

2. Chemical Bond Concept • Recall that an atom has core and valence electrons. • Core electrons are found close to the nucleus. • Valence electrons are found in the most distant s , pand partially filled dsubshells. • It is valence electrons that are responsible for holding two or more atoms together in a chemical bond.

3. Octet Rule • The octet rule states that atoms bond in such a way so that each atom acquires eight electrons in its outer shell. • There are two ways in which an atom may achieve an octet. (a) by transfer of electrons from one atom to another (b) by sharing one or more pairs of electrons

4. Types of Bonds • Ionic bonds are formed from the complete transfer of electrons between atoms to form ionic compounds. • Covalent bonds are formed when two atoms share electrons to form molecular compounds

5. Ionic Bonds An ionic bond is formed by the attraction between positively charged anions and negatively charged anions. • This “electrostatic attraction” is similar to the attraction between opposite poles on two magnets.

6. Ionic Bonding between Na and Cl

7. Ionic Bonds • The ionic bonds formed from the combination of anions and cations are very strong and result in the formation of a rigid, crystalline structure. The structure for NaCl, ordinary table salt, is shown here.

8. Energies involved in Ionic Bonding • Ionic bonds have energies, which are released when ionic bonds are formed • In general, the energy of the ionic bond according to Coulomb's Law is given as: E = qanion . qcation r where qanion and qcation are charges on the anion and the cation, • r is the distance between them The greater the charges and the smaller the distance between them, the stronger is this ionic bond.

9. Ionic Radii • The radius of a cation is smaller than the radius of its starting atom. • The radius of an anion is larger than the radius of its starting atom.

10. Ionic Radii Example 1 Rank the following ions in order of decreasing ionic radius: S2–, O2–, F–, Na+, Mg2+. a) S2–, O2–, F–, Na+, Mg2+ b) O2–, F–, Na+, Mg2+, S2– c) Mg2+, Na+, F–, O2–, S2– d) Mg2+, S2–, Na+, F–, O2– e) O2–, S2–, F–, Na+, Mg2+

11. Energies involved in Ionic Bonding • The transfer of an electron from a sodium atom to a chlorine atom is not in itself energetically favorable; it requires 147 kJ/mol of energy. • However, 493 kJ of energy is released when these oppositely charged ions come together. • An additional 293 kJ of energy is released when the ion pairs solidify.

12. Lattice Energy • The energy necessary to separate ionic solid into gaseous ions: NaCl (s) → Na+(g) + Cl-(g) ΔH = U • This ‘lattice energy’ is the negative of the energy released when gaseous ions form an ionic solid. • For a given arrangement of ions, lattice energy increases as the charges on the ions increase and as their radii decrease.

13. Lattice Energy Example 2 When the cations Na+, K+, Rb+, Cs+ are combined with chloride ion in the gas phase to form ion pairs, which pair formation releases the greatest amount of energy? a) NaCl b) KCl c) RbCl d) CsCl

14. Born Haber Cycle: Energetics of Ionic Bond Formation

15. Energetics of Ionic Bond Formation Example 3 In the Born–Haber cycle for NaCl(s), which of the following processes corresponds to the electron affinity of Cl? a) Cl(g) → Cl+(g) + e– b) NaCl(s) → Na+(g) + Cl–(g) c) Cl2(g) → 2Cl(g) d) Cl–(g) → Cl(g) + e– e) Cl(g) + e– → Cl–(g)

16. Lattice Energy Example 4 Which of the following compounds would be expected to have the highest melting point? a) LiF b) LiCl c) CsF d) NaBr

17. Electron Dot Formulas • An electron dot formula of an element shows the symbol of the element surrounded by its valence electrons. • We use one dot for each valence electron. • Let’s draw EDF of phosphorus.

18. Formation of Cations • We can use electron dot formulas to look at the formation of cations. • Each of the metals in Period 3 form cations by losing 1, 2, or 3 electrons, respectively. Each metal atom becomes isoelectronic with the preceding noble gas, neon.

19. Formation of Anions • We can also use electron dot formulas to look at the formation of anions. • The nonmetals in Period 3 gain 1, 2, or 3 electrons, respectively, to form anions. Each nonmetal ion is isoelectronic with the following noble gas, argon.

20. Electron Configuration of Ions Example 5 The ground-state electron configuration of the Mg 2+ ion is a) 1s22s22p6. b) 1s22s22p63s2. c) 1s22s22p63s23p2. d) 1s22s22p3.

21. Electron Configuration of Ions Example 6 What is the electron configuration for Cr2+? a) [Ar]4s1 3d5 b) [Ar]4s2 3d2 c) [Ar]3d3 d) [Ar]4s2 3d7 e) [Ar]3d4

22. Electron Configuration of Ions Example 7 Which of the following species is isoelectronic with Xe? a) Kr b) Rb+ c) Tl3+ d) Se2– e) Ba2+

23. Electron Configuration of Ions Example 8 Which pair of species is isoelectronic? a) Ne and Ar b) Na+ and K+ c) K+ and Cl– d) Li+ and Ne

24. Covalent Bonds • Covalent bonds are formed when two nonmetal atoms share electrons and the shared electrons in the covalent bond belong to both atoms. • In any covalent bond, the attractive energy between the nuclei and electrons exceeds the repulsive energy arising from nuclear-nuclear and electron-electron interactions • For a covalent bond to be formed, e.g. between two hydrogen atoms, the orbital overlap of the 1s orbitals gives a molecular orbital with high electron density between the nuclei.

25. Covalent Bonding Example 9 Which of the following is the best explanation for a covalent bond? a) electrons simultaneously attracted by more than one nucleus b) the overlapping of two electron-filled orbitals having different energies c) the overlapping of unoccupied orbitals of two or more atoms d) a positive ion attracting negative ions

26. Covalent Bonding

27. Bond Length • When a covalent bond is formed, the valence shells of the two atoms overlap with each other. • In HCl, the hydrogen 1s energy sublevel overlaps with the chlorine 3p energy sublevel. The mixing of sublevels draws the atoms closer together. • The distance between the two atoms is smaller than the sum of their atomic radii and is the bond length.

28. Bond Dissociation Energy • Energy is released when two atoms form a covalent bond: • H(g) + Cl(g)  HCl(g) + heat • Conversely, energy is needed to break a covalent bond. • The energy required to break a covalent bond is referred to as the bond energy. • To pull apart a H2 molecule into individual atoms,the bond dissociation energy is 104 kcal/mol

29. Types of Covalent Bonds Polar covalent bond: • A polar covalent bond is a covalent bond in which the bonding electrons between the atoms are shared unequally due to the electronegativity difference. Non polar covalent bonds: • A non polar covalent bond is a covalent bond in which the bonding electrons are shared equally because the two atoms are alike.

30. Electronegativity (EN or X) • Electronegativity - describes the ability of an atom to attract a shared pair of electrons • Scale devised by L. Pauling varies from 0 - 4. • Can be related to ionization energy and electron affinity of atoms, X = ½(IE + EA)

31. Trend in Electronegativities

32. Electronegativity EN decreases from the top to the bottom in a group and increases from left to right along a period. • ΔEN = 2 or > 2 ionic bond • ΔEN = 0 non polar covalent bond • ΔEN > 0 and < 2 polar covalent bond

33. Electronegativity Example 10 Which of the following atoms is the most electronegative? a) B b) Na c) N d) Cs e) Al

34. Electronegativity Example 11 The larger the difference in electronegativity between two bonded atoms, 1. the more ionic the bond. 2. the more covalent the bond. 3. the more polar the bond. a) 1 only b) 2 only c) 3 only d) 1 and 3 only e) 2 and 3 only

35. Valence Shell: • The outermost shell of an atom containing the valence electrons. • Core electrons: • The inaccessible electrons. Core electrons do not participate in the chemical reactions. • Bonding electrons: • The valence electrons which are shared (in bonding orbitals) between two nuclei. • Bonding pair: • Bond where two electrons are shared between two atoms. • Nonbonding electrons: • The valence electrons that are not involved in bonding & are localized in individual atoms. • Hydrogen is most stable with the duet (helium electron configuration).

36. Exceptions to the Octet Rule • Not all the molecules follow the octet rule. Many elements beyond the • 2nd row do not comply with this rule. • Single bond • When one pair of electrons is shared by two atoms. • Notation. Two dots or one line is used. • Multiple bonds (double bonds, triple bonds) • A covalent bond in which two pairs of electrons are shared by two atoms is called a double bond while a covalent bond in which the three pairs of electrons are shared between the two atoms is called a triple bond.

37. CONSTRUCTING LEWIS STRUCTURES: • STEP 1 • If the substance is ionic, treat each ion separately. • A compound is ionic if it contains a polyatomic anion and a metal or an ammonium cation. • If there is a metal atom but no polyatomic anion, apply electronegativity differences to identify whether or not the compound is ionic. • STEP 2 • Count the total number of valence electrons of all elements in a substance. • Forthe anions, add electrons and for cations subtract electrons.

38. STEP 3 • Assemble the bonding framework. Account for the electrons used in the framework. Distribute the rest in Step 4. • While drawing the framework keep in mind the following points. • The order in which atoms are listed in the formula indicates the bonding pattern. • Hydrogen atoms are always outer atoms. • Outer atoms other than hydrogen are the most electronegative. • Atoms enclosed in parentheses are bonded together and the entire group can be bound to another atom.

39. STEP 4 Place three nonbonding pairs of electrons on each outer atom except H. An outer atom other than hydrogen is more stable when it is associated with an octet of electrons.

40. STEP 5 • Assign the remaining valence electrons to inner atoms. • If the molecule has more than one inner atom, place nonbonding pairs around the most electronegative atom until it has an octet. • If there are still unassigned electrons, do the same for the next most electronegative atom. Continue in this manner until all the electrons have been assigned. • Place any remaining electrons on an inner atom that has n>2. Atoms with n>2 have valence dorbitals that allow them to accommodate more than 8 electrons. • Compounds with P, S and Cl may have between 8 & 12 electrons associated with these atoms.

41. STEP 6 • Satisfy electron configurations of the inner atoms. • In case an inner atom lacks an octet, a more stable distribution can be achieved by moving some of the electrons from outer atoms to make double or triple bonds to the inner atoms. • For inner atoms beyond the row 2, find the formal charge. If it leaves a + ve charge on the atom, shift electrons to form double bonds, even if this gives the inner atom more than eight electrons. Amore stable structure is achieved this way.

42. STEP 7 • Identify equivalent Lewis structures. • Draw resonance structures when there is more than one way to shift electron pairs. •  A formal charge calculation may be necessary if more than one Lewis dot formulas are possible. • A structure with Formal charge gives the approximate distribution of electrons. • Formal charge = valence electrons on free atom – ½ (number of bonding electrons in a bond) – (number of lone pair electrons) • To gets an accurate Lewis structure, minimize formal charges in the Lewis structure.

43. Execeptions to Octet Rule Example 12 In which of the following molecules is the octet rule violated? a) PF3 b) SiF4 c) OF2 d) ClF3 e) ClF

44. Execeptions to Octet Rule Example 13 Which of the following species represents an exception to the octet rule? a) CH3OH b) CCl4 c) PH3 d) BF3

45. Execeptions to Octet Rule Example 14 Which of the following species represents an exception to the octet rule? • a) CH3OH • b) CCl4 • c) PH3 • d) BF3

46. Execeptions to Octet Rule Example 15 Which of the following species represents an exception to the octet rule? a) CO2 b) SF4 c) SiO2 d) PCl3

47. Resonance in organic structures when a single Lewis structure cannot explain the properties of a molecule.  In such cases we draw more than one structure to represent the molecule and take the original structure of the molecule as the weighed average of these structures.

48. Resonance

49. Resonance Example 16 The concept of resonance describes molecular structures a) that have several different geometric arrangements. b) that have no suitable single Lewis formula. c) that have electrons resonating. d) that are formed from hybridized orbitals

50. Resonance Example 17 All the following statements about resonance are true except a) A single Lewis formula does not provide an adequate representation of the bonding. b) Resonance describes the oscillation and vibration of electrons. c) Resonance describes a more stable situation than does any one contributing resonance formula. d) Resonance describes the bonding as intermediate between the contributing resonance formulas. e) The contributing resonance formulas differ only in the arrangement of the electrons.