Chapter 5 Compounds and Their Bonds

1. Chapter 5 Compounds and Their Bonds Covalent Bonds Naming and Writing Formulas of Covalent Compounds Bond Polarity

2. Ionic Bond H2, A Covalent Molecule • In hydrogen, two hydrogen atoms share their electrons to form a covalent bond. • Each hydrogen atom acquires a stable outer shell of two (2) electrons like helium (He). H+H H : H = HH = H2 hydrogen molecule

3. A covalent bond between two hydrogen atoms is shown in this picture. Fig 5.1 A covalent bond is the result of attractive and repulsive forces between atoms.

4. Spherical 1S orbital of two individual hydrogen atoms bends together and overlap to give an egg shaped region in the hydrogen molecule. The shared pair of electrons in a covalent bond is often represented as a line between atoms.

5. Bond length: The optimum distance between nuclei involved in a covalent bond. If the atoms are too far apart, the attractive forces are small and no bond exists. If the atoms are too close, the repulsive interaction between the nuclei is so strong that it pushes the atoms apart, Fig 5.2.

6. When two chlorine atoms approach each other, the unpaired 3p electrons are shared by both atoms in a covalent bond. Each chlorine atom in the Cl2 molecule now have 6 electrons in its own valence shell and sharing two giving each valence shell octet.

7. Diatomic Elements • As elements, the following share electrons to form diatomic, covalent molecules.

8. In addition to H2 and Cl2, five other elements always exist as diatomic molecule.

9. Learning Check What is the name of each of the following diatomic molecules? H2 hydrogen N2 nitrogen Cl2 _______________ O2 _______________ I2 _______________

10. Solution What are the names of each of the following diatomic molecules? H2 hydrogen N2 nitrogen Cl2chlorine O2oxygen I2iodine

11. Covalent Bonds in NH3 • The compound NH3 consists of a N atom and three H atoms.   N and 3 H  • By sharing electrons to form NH3, the electron dot structure is written as H Bonding pairs   H : N : H  Lone pair of electrons

12. Number of Covalent Bonds • Often, the number of covalent bonds formed by a nonmetal is equal to the number of electrons needed to complete the octet.

13. Dot Structures and Models of Some Covalent Compounds

14. Multiple Bonds • Sharing one pair of electrons is a single bond.X : X or X–X • In multiple bonds, two pairs of electrons are shared to form a double bond or three pairs of electrons are shared in a triple bond.X : : X or X =XX ::: X or X ≡X

15. Multiple Bonds in N2 • In nitrogen, octets are achieved by sharing three pairs of electrons. • When three pairs of electrons are shared, the multiple bond is called a triple bond. octets       N  + N  N:::N  triple bond

16. 5.4 Coordinate Covalent Bonds • Coordinate Covalent Bond: The covalent bond that forms when both electrons are donated by the same atom.

17. Fig 5.7 Electronegativities and the periodic table

18. 5.6 Drawing Lewis Structure • Draw skeletal structure of compound showing what atoms are bonded to each other. Put the unique element ( or least electronegative atom) in the center. • Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge. • Complete an octet for all atoms except hydrogen • If structure contains too many electrons, form double and triple bonds on central atom as needed.

19. Write the Lewis structure of nitrogen trifluoride (NF3). F N F F Step 1 – N is less electronegative than F, put N in center Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5) 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

20. Write the Lewis structure of the carbonate ion (CO32-). O C O 2 single bonds (2x2) = 4 1 double bond = 4 8 lone pairs (8x2) = 16 O Total = 24 Step 1 – C is less electronegative than O, put C in center Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4) -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 Step 5 - Too many electrons, form double bond and re-check # of e-

21. H H C O H C O H formal charge on an atom in a Lewis structure total number of valence electrons in the free atom total number of nonbonding electrons ( total number of bonding electrons ) 1 - - = 2 Two possible skeletal structures of formaldehyde (CH2O) An atom’s formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure. The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion.

22. H C O H C : 4 e- O : 6 e- 2H :2x1 e- 12 e- formal charge on an atom in a Lewis structure total number of valence electrons in the free atom total number of nonbonding electrons ( total number of bonding electrons ) 1 - - = 2 single bonds (2x2) = 4 2 1 double bond = 4 2 lone pairs (2x2) = 4 Total = 12 -1 +1 formal charge on C = 4 - 2- ½ x 6 = -1 formal charge on O = 6 - 2- ½ x 6 = +1

23. C – 4 e- H O – 6 e- C O H 2H – 2x1 e- 12 e- formal charge on an atom in a Lewis structure total number of valence electrons in the free atom total number of nonbonding electrons ( total number of bonding electrons ) 1 - - = 2 single bonds (2x2) = 4 2 1 double bond = 4 2 lone pairs (2x2) = 4 Total = 12 0 0 formal charge on C = 4 - 0- ½ x 8 = 0 formal charge on O = 6 - 4- ½ x 4 = 0

24. Which is the most likely Lewis structure for CH2O? H C O H H C O H -1 +1 0 0 Formal Charge and Lewis Structures • For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present. • Lewis structures with large formal charges are less plausible than those with small formal charges. • Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms.

25. 5.7 Shape of Molecules VSEPR • The shape of a molecule is predicted from the geometry of the electron pairs around the central atom. • In the valence-shell electron-pair repulsion theory (VSEPR), the electron pairs are arranged as far apart as possible to give the least amount of repulsion of the negatively charged electrons.

26. Two Electron Pairs • In a molecule of BeCl2, there are two bonding pairs around the central atom Be. (Be is an exception to the octet rule.) • The arrangement of two electron pairs to minimize their repulsion is 180° or opposite each other. • The shape of the molecule is linear.

27. Two Electron Pairs with Double Bonds • The electron-dot structure for CO2 consists of two double bonds to the central atom C. • Because the electrons in a double bond are held together, a double bond is counted as a single unit. • Repulsion is minimized when the double bonds are placed opposite each other at 180° to give a linear shape.

28. Three Electron Pairs • In BF3, there are 3 electron pairs around the central atom B. (B is an exception to the octet rule.) • Repulsion is minimized by placing three electron pairs in a plane at angles of 120°, which is a trigonal planar arrangement. • The shape with three bonded atoms is trigonal planar.

29. Two Bonding Pairs and A Nonbonding Pair • In SO2, there are 3 electron units around the central atom S. • Two electron units are bonded to atoms and one electron pair is a nonbonding pair. • Repulsion is minimized by placing three electron pairs in a plane at angles of 120°, which is trigonal planar. • The shape with two bonded atoms is bent.

30. Four Electron Pairs • In CH4, there are 4 electron pairs around the central atom C. • Repulsion is minimized by placing four electron pairs at angles of 109°, which is a tetrahedral arrangement. • The shape with four bonded atoms is called tetrahedral.

31. Three Bonding Atoms and One Nonbonding Pair • In NH3, there are 4 electron pairs around the N. • Three pairs are bonded to atoms and one is a nonbonding pair. • Repulsion is minimized by placing four electron pairs at angles of 109°, which is a tetrahedral arrangement. • The shape with three bonded atoms is pyramidal.

32. Two Bonding Atoms and Two Lone Pairs • In H2O, there are 4 electron pairs around O. • Two pairs are bonded to atoms and two are nonbonding pairs. • Repulsion is minimized by placing four electron pairs at angles of 109° called a tetrahedral arrangement. • The shape with two bonded atoms is called bent.

33. Some Steps Using VSEPR to Predict Shape • Draw the electron dot structure. • Count the charged clouds around the central atom. • Arrange the charged clouds to minimize repulsion. • Determine the shape using the number of bonded atoms in the electron arrangement.

34. Summary of Electron Arrangements and Shapes Number of atoms bonded to the central atom

35. The shape depends on the number of charged clouds surrounding the atom as summarized in Table 5.1

36. Learning Check Use VSEPR theory to determine the shape of the following molecules or ions. 1) tetrahedral 2) pyramidal 3) bent A. PF3 B. H2S C. CCl4 D. PO43-

37. Solution Use VSEPR theory to determine the shape of the following molecules or ions. 1) tetrahedral 2) pyramidal 3) bent A. PF3 2) pyramidal B. H2S 3) bent C. CCl4 1) tetrahedral D. PO43- 1) tetrahedral

38. Comparing Nonpolar and Polar Covalent Bonds

39. Electronegativity • Electronegativity is the attraction of an atom for shared electrons. • The nonmetals have high electronegativity values with fluorine as the highest. • The metals have low electronegativity values.

40. Fig 5.7 Electronegativities and the periodic table

42. Nonpolar Covalent Bonds • The atoms in a nonpolar covalent bond have electronegativity differences of 0.4 or less. • Examples: Atoms Electronegativity Type of Difference Bond N-N3.0 - 3.0 = 0.0 Nonpolar covalentCl-Br3.0 - 2.8 = 0.2 Nonpolar covalentH-Si 2.1 - 1.8 = 0.3 Nonpolar covalent

43. Polar Covalent Bonds • The atoms in a polar covalent bond have electronegativity differences of 0.5 to 1.9. • Examples: Atoms Electronegativity Type of Difference BondO-Cl3.5 - 3.0 = 0.5 Polar covalentCl-C3.0 - 2.5 = 0.5 Polar covalentO-S 3.5 - 2.5= 1.0 Polar covalent

44. Comparing Nonpolar and Polar Covalent Bonds

45. Polar, Nonpolar and Ionic Bond

46. Ionic Bonds • The atoms in an ionic bond have electronegativity differences of 2.0 or more. • Examples: Atoms Electronegativity Type of Difference BondCl-K3.0 – 0.8 = 2.2 IonicN-Na3.0 – 0.9 = 2.1 Ionic

47. Predicting Bond Type

48. Learning Check Identify the type of bond between the following as 1) nonpolar covalent 2) polar covalent 3) ionic A. K-N B. N-O C. Cl-Cl

49. Solution A. K-N 3) ionic B. N-O 2) polar covalent C. Cl-Cl 1) nonpolar covalent

50. 5.9 Polar Molecules • Entire molecule can be polar if electrons are attracted more strongly to one part of the molecule than to another. • Molecule’s polarity is due to the sum of all individual bond polarities and lone-pair contribution in the molecule.