1. The Ties That Bind Chemical Bonding and Interactions

2. 1. Stable Electron Configurations 2. Electron-Dot (Lewis) Structures 1. Drawing, Rules for Drawing 2. The Octet Rule 3. Some Exceptions to the Rule 3. Ionic Bonding 1. Naming ionic compounds 2. Drawing 4. Covalent Bonding 1. Naming covalent compounds 2. Drawing 3. Electronegativity and Polar Covalent Compounds 5. Molecular Shapes and the VSEPR Theory 6. Intermolecular Forces of Attraction 1. H-bonds, Dipole-Dipole, Ion-Dipole, London Dispersion Forces

3. It’s all about stability… 1. Fact: Noble gases are inert (undergo few, if any, chemical reactions) 2. Theory: The inertness of these gases is a result of their electronic structure 3. If elements could alter their electron structures like those of the noble gases, they would be less reactive.

4. Electron Configurations of Cations and Anions Na [Ne]3s 1 Na + [Ne] Ca [Ar]4s 2 Ca 2+ [Ar] Al [Ne]3s 2 3p 1 Al 3+ [Ne] Atoms lose electrons so that cation has a noble-gas outer electron configuration. H 1s 1 H - 1s 2 or [He] F 1s 2 2s 2 2p 5 F - 1s 2 2s 2 2p 6 or [Ne] O 1s 2 2s 2 2p 4 O 2- 1s 2 2s 2 2p 6 or [Ne] N 1s 2 2s 2 2p 3 N 3- 1s 2 2s 2 2p 6 or [Ne] Atoms gain electrons so that anion has a noble-gas outer electron configuration. Of Representative Elements

5. +1+2+3 -2-3 Cations and Anions Of Representative Elements

6. Na + : [Ne]Al 3+ : [Ne]F - : 1s 2 2s 2 2p 6 or [Ne] O 2- : 1s 2 2s 2 2p 6 or [Ne]N 3- : 1s 2 2s 2 2p 6 or [Ne] Na +, Al 3+, F -, O 2-, and N 3- are all isoelectronic with Ne What neutral atom is isoelectronic with H - ? H - : 1s 2 same electron configuration as He

7. 9.1 Valence electrons are the outer shell electrons of an atom. The valence electrons are the electrons that particpate in chemical bonding. 1A 1ns 1 2A 2ns 2 3A 3ns 2 np 1 4A 4ns 2 np 2 5A 5ns 2 np 3 6A 6ns 2 np 4 7A 7ns 2 np 5 Group# of valence e - e - configuration

8. (Electron Dot Structures) *We generally place the electrons one four sides of a square around the element symbol. *Octet rule: we know that s 2 p 6 is a noble gas configuration. We assume that an atom is stable when surrounded by 8 electrons (4 electron pairs).

9. A violent bond: Sodium and Chlorine Sodium is a soft, reactive metal. Chlorine is a greenish-yellow gas which is toxic (WWI poison) Sodium chloride is used for our food.

10. Ionic Bonding Consider the reaction between sodium and chlorine: Na(s) + ½Cl 2 (g)  NaCl(s)

11. Chemistry In Action: Sodium Chloride Mining SaltSolar Evaporation for Salt

12. 9.2 Li + F Li + F - The Ionic Bond 1s 2 2s 1 1s 2 2s 2 2p 5 1s 2 1s 2 2s 2 2p 6 [He][Ne] Li Li + + e - e - + FF - F - Li + + Li + F - IONIC BONDING -bonding due to electrostatic attraction arising from an exchange of electrons.

13. ionic compounds consist of a combination of cations and an anions the formula is always the same as the empirical formula the sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero The ionic compound NaCl 2.6

14. Ionic Bonding *NaCl forms a very regular structure in which each Na + ion is surrounded by 6 Cl  ions. Similarly, each Cl  ion is surrounded by six Na + ions. There is a regular arrangement of Na + and Cl  in 3D. Note that the ions are packed as closely as possible. Note that it is not easy to find a molecular formula to describe the ionic lattice.

15. 9.3 Lattice energy (E) increases as Q increases and/or as r decreases. cmpd lattice energy MgF 2 MgO LiF LiCl 2957 3938 1036 853 Q= +2,-1 Q= +2,-2 r F < r Cl SIDEBAR: Electrostatic (Lattice) Energy E = k Q+Q-Q+Q- r Q + is the charge on the cation Q - is the charge on the anion r is the distance between the ions Lattice energy (E) is the energy required to completely separate one mole of a solid ionic compound into gaseous ions.

16. 9.3

17. An ion is an atom, or group of atoms, that has a net positive or negative charge. cation – ion with a positive charge If a neutral atom loses one or more electrons it becomes a cation. anion – ion with a negative charge If a neutral atom gains one or more electrons it becomes an anion. Na 11 protons 11 electrons Na + 11 protons 10 electrons Cl 17 protons 17 electrons Cl - 17 protons 18 electrons 2.5

18. A monatomic ion contains only one atom A polyatomic ion contains more than one atom 2.5 Na +, Cl -, Ca 2+, O 2-, Al 3+, N 3- OH -, CN -, NH 4 +, NO 3 -

19. Test Yourself: 1. MgO 2. MgBr 2 3. The combination of the ion of magnesium and the ion of nitrogen

20. Formula of Ionic Compounds Al 2 O 3 2.6 2 x +3 = +63 x -2 = -6 Al 3+ O 2- CaBr 2 1 x +2 = +22 x -1 = -2 Ca 2+ Br - Na 2 CO 3 1 x +2 = +21 x -2 = -2 Na + CO 3 2-

21. 2.6

22. 2.7

23. Chemical Nomenclature Ionic Compounds  often a metal + nonmetal  anion (nonmetal), add “ide” to element name BaCl 2 barium chloride K2OK2O potassium oxide Mg(OH) 2 magnesium hydroxide KNO 3 potassium nitrate 2.7

24. Transition metal ionic compounds  indicate charge on metal with Roman numerals FeCl 2 2 Cl - -2 so Fe is +2 iron(II) chloride FeCl 3 3 Cl - -3 so Fe is +3 iron(III) chloride Cr 2 S 3 3 S -2 -6 so Cr is +3 (6/2)chromium(III) sulfide 2.7

25. Test Yourself 1. Write the formula and draw the ionic structure 1. Magnesium (II) bromide 2. Iron (III) Oxide 3. Sodium Azide 4. Ammonium Chloride 5. Magnesium Nitrate 6. Ammonium Phosphate 2. Write the name 1. FeO 2. Fe 2 O 3 3. LiBr 4. KCrO 4 5. K 2 Cr 2 O 7

26. 2.7

27. Covalent Bonding

28. A covalent bond is a chemical bond in which two or more electrons are shared by two atoms. Why should two atoms share electrons? FF + 7e - FF 8e - F F F F Lewis structure of F 2 lone pairs single covalent bond 9.4

29. 8e - H H O ++ O HH O HHor 2e - Lewis structure of water Double bond – two atoms share two pairs of electrons single covalent bonds O C O or O C O 8e - double bonds Triple bond – two atoms share three pairs of electrons N N 8e - N N triple bond or 9.4

30. Bond Type Bond Length (pm) C-CC-C 154 CCCC 133 CCCC 120 C-NC-N 143 CNCN 138 CNCN 116 Lengths of Covalent Bonds Bond Lengths Triple bond < Double Bond < Single Bond 9.4

32. H F F H Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms electron rich region electron poor region e - riche - poor ++ -- 9.5

33. Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond. Electron Affinity - measurable, Cl is highest Electronegativity - relative, F is highest X (g) + e - X - (g) 9.5

35. Bond Polarity and Electronegativity Electronegativity

36. Drawing Lewis Structures 1. Add the valence electrons. 2. Identify the central atom (usually the one with the highest molecular mass, lowest electronegativity, or closest to the center of the periodic table). 3. Place the central atom in the center of the molecule and add all other atoms around it. 4.Place one bond (two electrons) between each pair of atoms. 5.Complete the octet for the central atom. 6.Complete the octets for all other atoms. Use double bonds if necessary. 7. Place remaining electrons on the central atom.

37. Write the Lewis structure of nitrogen trifluoride (NF 3 ). Step 1 – N is less electronegative than F, put N in center FNF F Step 2 – Count valence electrons N - 5 (2s 2 2p 3 ) and F - 7 (2s 2 2p 5 ) 5 + (3 x 7) = 26 valence electrons Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms. Step 4 - Check, are # of e - in structure equal to number of valence e - ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons 9.6

38. Write the Lewis structure of the carbonate ion (CO 3 2- ). Step 1 – C is less electronegative than O, put C in center OCO O Step 2 – Count valence electrons C - 4 (2s 2 2p 2 ) and O - 6 (2s 2 2p 4 ) -2 charge – 2e - 4 + (3 x 6) + 2 = 24 valence electrons Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms. Step 4 - Check, are # of e - in structure equal to number of valence e - ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons 9.6 Step 5 - Too many electrons, form double bond and re-check # of e - 2 single bonds (2x2) = 4 1 double bond = 4 8 lone pairs (8x2) = 16 Total = 24

39. Trial 1. Write the Lewis stucture for 1. H 2 O 2. CO 2 3. NH 3 4. CH 4 5. N 2 6. BH 3

40. Exceptions to the Octet Rule The Incomplete Octet HHBe Be – 2e - 2H – 2x1e - 4e - BeH 2 BF 3 B – 3e - 3F – 3x7e - 24e - FBF F 3 single bonds (3x2) = 6 9 lone pairs (9x2) = 18 Total = 24 9.9

41. Exceptions to the Octet Rule Odd-Electron Molecules N – 5e - O – 6e - 11e - NO N O The Expanded Octet (central atom with principal quantum number n > 2) SF 6 S – 6e - 6F – 42e - 48e - S F F F F F F 6 single bonds (6x2) = 12 18 lone pairs (18x2) = 36 Total = 48 9.9

42. Strengths of Covalent Bonds Bond Enthalpies and Bond Length

43. Covalent share e - Polar Covalent partial transfer of e - Ionic transfer e - Increasing difference in electronegativity Classification of bonds by difference in electronegativity DifferenceBond Type 0Covalent  2 Ionic 0 < and <2 Polar Covalent 9.5

44. Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H 2 S; and the NN bond in H 2 NNH 2. Cs – 0.7Cl – 3.03.0 – 0.7 = 2.3Ionic H – 2.1S – 2.52.5 – 2.1 = 0.4Polar Covalent N – 3.0 3.0 – 3.0 = 0Covalent 9.5