Lecturenote 3 General Chemistry: Atoms &

1. Lecturenote 3 General Chemistry: Atoms & Dr.Sadeem M.Al-barody

2. 3 Chemical Bonding I: Basic Concepts 3.1 The Octet Rule Lewis Structures Multiple Bonds 3.2 Electronegativity and Polarity Electronegativity Dipole Moment, Partial Charges and Percent Ionic Character 3.3 Drawing Lewis Structures 3.4 Lewis Structures and Formal Charge 3.5 Resonance 3.6 Exceptions to the Octet Rule Incomplete Octets Odd Numbers of Electrons Expanded Octets

3. 3.1 • • • • • • • • • • • • • F• •F FF + • • • • • • • • • • • • • The Octet Rule According to the octet rule, atoms will lose, gain, or share electrons in order to achieve a noble gas electron configuration. Only valence electrons contribute to bonding. Each F counts both shared electrons to “feel” as though it has a Ne configuration. F 1s22s22p5 valence electrons

4. • • • • • • • • • • • • • FF FF • • • • • • • • • • • • • • • • • • • • FF • • • • • • • The Octet Rule Only two valence electrons participate in the formation of the F2 bond. Pairs of valence electrons not involved in bonding are called lone pairs. or

5. •• •• •• •• H O H H O H •• •• Lewis Structures A Lewis structureis a representation of covalent bonding. Shared electron pairs are shown either as dashes or as pairs of dots. Lone pairs are shown as pairs of dots on individual atoms. H with 2 e− H with 2 e− Shared electrons shown as dashes (bonds) O with 8 e−

6. •• •• •• •• •• •• •• •• •• O C O H O H •• •• •• •• •• •• H O H = = O C O •• •• •• Multiple Bonds In a single bond, atoms are held together by one electron pair. In a double bond,atoms share two pairs of electrons. C with 8 e− O with 8 e− O with 8 e− 2 shared pairs of electrons result in double bonds one shared pair of electrons results in a single bond

7. •• •• •• •• •• •• •• ≡ N N N N Multiple Bonds A triple bondoccurs when atoms are held together by three electron pairs. each N has 8 e− 3 shared pairs of electrons result in a triple bond

8. Multiple Bonds Bond length is defined as the distance between the nuclei of two covalently bonded atoms.

9. N≡N N=N N−N 110 pm 124 pm 147 pm Multiple Bonds Multiple bonds are shorter than single bonds. For a given pair of atoms: • triple bonds are shorter than double bonds • double bonds are shorter than single bonds < <

10. Multiple Bonds The shorter multiple bonds are also stronger than single bonds. We quantify bond strength by measuring the quantity of energy required to break it. H2(g) → H(g) + H(g) Bond energy = 436.4 kJ/mol We will examine this in more detail in Chapter 10.

11. 3.2 M:X Mδ+Xδ− M+X− Pure covalent bond Neutral atoms held together by equally shared electrons Polar covalent bond Partially charged atoms held together by unequally shared electrons Ionic bond Oppositely charged ions held together by electrostatic attraction Electronegativity and Polarity There are two extremes in the spectrum of bonding: • covalent bonds occur between atoms that share electrons • ionic bonds occur between a metal and a nonmetal and involve ions Bonds that fall between these extremes are polar. In polar covalent bonds, electrons are shared but not shared equally.

12. Electronegativity and Polarity Electron density maps show the distributions of charge. • Electrons spend a lot of time in red and very little time in blue. Electrons are shared equally nonpolar covalent Electrons are not shared equally and are more likely to be associated with F polar covalent Electrons are not shared but rather transferred from Na to F ionic

13. Electronegativity Electronegativity is the ability of an atom in a compound to draw electrons to itself.

14. Electronegativity and Polarity Electronegativity varies with atomic number.

15. Electronegativity There is no sharp distinction between nonpolar covalent and polar covalent or between polar covalent and ionic. The following rules help distinguish among them: • A bond between atoms whose electronegativites differ by less than 0.5 is general considered purely covalent or nonpolar. • A bond between atoms who’s electronegativies differ by the range of 0.5 to 2.0 is generally considered polar covalent. • A bond between atoms whose electronegativities differ by 2.0 or more is generally considered ionic.

16. Worked Example 3.1 Classify the following bonds as nonpolar, polar, or ionic: (a) the bond in ClF, (b) the bond in CsBr, and (c) the carbon-carbon double bond in C2H4. StrategyElectronegativity values are: Cl (3.0), F (4.0), Cs (0.7), Br (2.8), C (2.5). Use this information to determine which bonds have identical, similar, and widely different electronegativities. Solution • The difference between the electronegativies of F and Cl is 4.0 – 3.0 = 1.0, making the bond in ClF polar. • In CsBr, the difference is 2.8 – 0.7 = 2.1, making the bond ionic. • In C2H4, the two atoms are identical. (Not only are they the same element, but each C atom is bonded to two H atoms.) The carbon-carbon double bond is C2H4 is nonpolar. Think About ItBy convention, the difference in electronegativity is always calculated by subtracting the smaller number from the larger one, so the result is always positive.

17. Dipole Moment, Partial Charges, and Percent Ionic Character •• •• H− F •• δ+ δ− •• •• H− F •• Regions where electrons spend little time Regions where electrons spend a lot of time An arrow is used to indicate the direction of electron shift in polar covalent molecules. The consequent charge separation can be represented as: • Deltas (δ) denote a partial positive or negative charge.

18. μ = Q x r 1 D = 3.336×10−30 C∙m Dipole Moment, Partial Charges, and Percent Ionic Character A quantitative measure of the polarity of a bond is its dipole moment (μ). • Q is the charge (C) . • r is the distance between the charges (m). • μ is always positive and expressed in debye units (D).

19. μ = Q x r Dipole Moment, Partial Charges, and Percent Ionic Character A quantitative measure of the polarity of a bond is its dipole moment (μ).

20. 3.336×10-30 C∙m 1×10-10 m 1 Å Worked Example 3.2 Burns caused by hydrofluoric acid [HF(aq)] are unlike any other acid burns and present unique medical complications. HF solutions typically penetrate the skin and damage internal tissues, including bone, often with minimal surface damage. Less concentrated solutions actually can cause greater injury than more concentrated ones by penetrating more deeply before causing injury, thus delaying the onset of symptoms and preventing timely treatment. Determine the magnitude of the partial positive and partial negative charges in the HF molecule. StrategySolve for Q. Convert the resulting charge in coulombs to units of electronic charge. According to Table 3.2, μ = 1.82 D and r = 0.92 Å for HF. The dipole moment must be converted from debye to C∙m and the distance between ions must be converted to meters. μ = 1.82 D = 6.07×10-30 C∙m r = 0.92 Å × = 9.2×10-11

21. μ r 6.07×10-30 C∙m 9.2×10-11 m 1 e- 1.6022×10-19 C •• •• H− F •• Worked Example 3.2 (cont.) Solution In coulombs: In units of electronic charge: Therefore, the partial charges in HF are +0.41 and -0.41 on H and F, respectively. = Q = = 6.598×10-20C 6.598×10-20 C × = 0.41 e- +0.41 -0.41 Think About ItCalculated partial charges should always be less than 1. If a “partial” charge were 1 or greater, it would indicate that at least one electron had been transferred from one atom to the other. Remember that polar bonds involve unequal sharing of electrons, not a complete transfer of electrons.

22. μ (observed) μ (calculated assuming discrete charges) percent ionic character = ×100% Dipole Moment, Partial Charges, and Percent Ionic Character Although the designations “covalent,” “polar covalent,” and “ionic” can be useful, sometimes chemists wish to describe and compare chemical bonds with more precision. Comparing the calculated dipole moment with the measured values gives us a quantitative way to describe the nature of a bond using the term percent ionic character.

23. Dipole Moment, Partial Charges, and Percent Ionic Character The figure below demonstrates the relationship between percent ionic character and the electronegativity difference in a heteronuclear diatomic molecule.

24. Worked Example 3.3 Using data from Table 3.2, calculate the percent ionic character of the bond in HI. StrategyCalculate the dipole moment in HI assuming that the charges on H and I are +1 and -1, respectively; bond length in HI is 1.61 Å , then calculate the percent ionic character. The magnitude of the charges must be expressed in coulombs SetupFrom Table 3.2, the bond length in HI is 1.61 Å (1.61×10-10 m) and the measured dipole moment of HI is 0.44 D. (1 e- = 1.6022×10-19 C); the bond length (r) must be expressed as meters (1 Å = 1×10-10 m); and the calculated dipole moment should be expressed as debyes (1 D = 3.336×10-30 C∙m).

25. 1 D 3.336×10-30 C∙m 0.44 D 7.73 D Worked Example 3.3 (cont.) Solution The dipole we would expect if the magnitude of the charges were 1.6022×10-19 C is Converting to debyes gives The percent ionic character of the H–I bond is ×100% = 5.7%. = 2.58×10-29C∙m μ = Q ×r = (1.6022×10-19 C) × (1.61×10-10 m) 2.58×10-29 C∙m × = 7.73 D Think About ItA purely covalent bond (in a homonuclear diatomic molecule such as H2) would have 0 percent ionic character. In theory, a purely ionic bond would be expected to have 100 percent ionic character, although no such bond is known.

26. 3.3 Drawing Lewis Structures Follow these steps when drawing Lewis structure for molecules and polyatomic ions. • Draw the skeletal structure of the compound. The least electronegative atom is usually the central atom. Draw a single covalent bond between the central atom and each of the surrounding atoms. • Count the total number of valence electrons present; add electrons for negative charges and subtract electrons for positive charges. • For each bond in the skeletal structure, subtract two electrons from the total valence electrons. • Use the remaining electrons to complete octets of the terminal atoms by placing pairs of electrons on each atom. Complete the octets of the most electronegative atom first. • Place any remaining electrons in pairs on the central atom. • If the central atom has fewer than eight electrons, move one or more pairs from the terminal atoms to form multiple bonds between the central atom and terminal atoms.

27. Drawing Lewis Structures Draw the Lewis Structure for ClO3−. Solution: Step 1 Draw the skeletal structure of the compound. The least electronegative atom is usually the central atom. Draw a single covalent bond between the central atom and each of the surrounding atoms. chlorine’s electronegativity = 3.0 oxygen’s electronegativity = 3.5

28. Drawing Lewis Structures Draw the Lewis Structure for ClO3−. Solution: Step 2 Count the total number of valence electrons present; add electrons for negative charges and subtract electrons for positive charges. Cl: 1 x 7 = 7 e− O: 3 x 6 = 18 e− Charge: 1 e− Total: 26 valence e−

29. Drawing Lewis Structures

30. Drawing Lewis Structures Draw the Lewis Structure for ClO3−. Solution: Step 4 Use the remaining electrons to complete octets of the terminal atoms by placing pairs of electrons on each atom. Complete the octets of the most electronegative atom first. 6 electrons used in single bonds 20 electrons must be placed in the structure starting with terminal atoms The structure now has a total of 24 valence electrons of the available 26.

31. Drawing Lewis Structures Draw the Lewis Structure for ClO3−. Solution: Step 5 Place any remaining electrons on the central atom. The structure now has a total of 24 valence electrons of the available 26. 2 e− remain to be added to the structure. Lewis structures for polyatomic ions are enclosed by square brackets.

32. Drawing Lewis Structures Draw the Lewis Structure for ClO3−. Solution: Step 6 If the central atom has fewer than eight electrons, move one or more pairs from the terminal atoms to form multiple bonds between the central atom and terminal atoms. All atoms already have an octet in this structure.

33. Drawing Lewis Structures

34. Worked Example 3.4 Draw the Lewis structure for carbon disulfide (CS2). Setup Step 1: C and S have identical electronegativities. We will draw the skeletal structure with the unique atom, C, at the center. Step 2: The total number of valence electrons is 16: 6 from each S atom and 4 from the C atom [2(6) + 4 = 16]. Step 3: Subtract 4 electrons to account for the bonds in the skeletal structure, leaving us 12 electrons to distribute. Step 4: Distribute the 12 remaining electrons as 3 lone pairs on each S atom.

35. Worked Example 3.4 (cont.) Draw the Lewis structure for carbon disulfide (CS2). Setup Step 5: There are no electrons remaining after step 4, so step 5 does not apply. Step 6: To complete carbon’s octet, use one lone pair from each S atom to make a double bond to the C atom. Solution Think About ItCounting the total number of valence electrons should be relatively simple to do, but it is often done hastily and is therefore a potential source of error in this type of problem. Remember that the number of valence electrons for each element is equal to the group number of that element.

36. 3.4 Formal charge = valence electrons – associated electrons Lewis Structures and Formal Charge Formal charge can be used to determine the most plausible Lewis Structure when more than one possibility exists for a compound. To determine associated electrons: 1)All the atom’s nonbonding electrons are associated with the atom. 2)Half the atom’s bonding electrons are associated with the atom.

37. •• •• •• •• 4 unshared + 4 shared = 6 e− = − O O O •• •• 2 2 unshared + 6 shared = 5 e− 6 unshared + 2 shared = 7 e− 2 2 Lewis Structures and Formal Charge Determine the formal charges on each oxygen atom in the ozone molecule (O3). Valence e− 6 6 6 e− associated with atom 6 5 7 0 +1 −1 Difference (formal charge)

38. Worked Example 3.5 The widespread use of fertilizers has resulted in the contamination of some groundwater with nitrates, which are potentially harmful. Nitrate toxicity is due primarily to its conversion in the body to nitrite (NO2-), which interferes with the ability of hemoglobin to transport oxygen. Determine the formal charges on each atom in the nitrate ion (NO3-). StrategyFollow the six steps to draw the Lewis structure of (NO3-). For each atom, subtract the associated electrons from the valence electrons. Setup The N atom has five valence electrons and four associated electrons (one from each single bond and two from the double bond). Each singly bonded O atom has six valence electrons and seven associated electrons (six in three lone pairs and one from the single bond). The doubly bonded O atom has six valence electrons and six associated electrons (four in two lone pairs and two from the double bond.)

39. Worked Example 3.5 (cont.) Solution The formal charges are as follows: +1 (N atom), -1 (singly bonded O atoms), and 0 (doubly bonded O atom). Think About ItThe sum of formal charges (+1) + (-1) + (-1) + (0) = -1 is equal to the overall charge on the nitrate ion.

40. •• •• •• •• •• = = O C O ≡ − O C O •• •• •• Lewis Structures and Formal Charge When there is more than one possible structure, the best arrangement is determined by the following guidelines: 1)A Lewis structure in which all formal charges are zero is preferred. 2)Small formal charges are preferred to large formal charges. 3)Formal charges should be consistent with electronegativities. Formal charge 0 0 0 +1 0 −1 better structure (based on formal charge)

41. Formal charge = valence electrons – associated electrons Lewis Structures and Formal Charge Based on formal charge, identify the best and the worst structures for the isocyanate ion below: Solution: Step 1 Assign formal charges on each atom using the formula Ve− 4 5 6 4 5 6 4 5 6 6 4 6 7 4 5 Ae− 5 4 7 FC −2 +1 0 −1 +1 −1 −3 +1 +1

42. Lewis Structures and Formal Charge Based on formal charge, identify the best and the worst structures for the isocyanate ion below: Solution: Step 2 Determine the best and worst structure FC −2 +1 0 −1 +1 −1 −3 +1 +1 Best structure: Small formal charges Formal charges more consistent with electronegativities Worst structure: Large formal charges Formal charges inconsistent with electronegativities

43. Worked Example 3.6 Formaldehyde (CH2O), which can be used to preserve biological specimens, is commonly sold as a 37 percent aqueous solution. Use formal charges to determine which skeletal arrangements of atoms shown here is the best choice for the Lewis structure of CH2O. StrategyThe complete Lewis structures for the skeletons shown are:

44. Worked Example 3.6 (cont.) StrategyThe complete Lewis structures for the skeletons shown are: In the structure on the left, the formal charges are as follows: Both H atoms: 1 valence e-– 1 associated e- (from single bond) = 0 C atom: 4 valence e-– 5 associated e- (two in the lone pair, one from the single bond, and two from the double bond) = -1 O atom: 6 valence e-– 5 associated e- (two from the lone pair, one from the single bond, and two from the double bond) = +1 Formal charges 0 -1 +1 0

45. Worked Example 3.6 (cont.) StrategyThe complete Lewis structures for the skeletons shown are: In the structure on the right, the formal charges are as follows: Both H atoms: 1 valence e-– 1 associated e- (from single bond) = 0 C atom: 4 valence e-– 4 associated e- (one from each single bond, and two from the double bond) = 0 O atom: 6 valence e-– 6 associated e- (four from the two lone pairs and two from the double bond) = 0 Formal charges all zero

46. Worked Example 3.6 (cont.) Solution Of the two possible arrangements, the structure on the left has an O atom with a positive formal charge, which is inconsistent with oxygen’s high electronegativity. Therefore, the structure on the right, in which both H atoms are attached directly to the C atoms and all atoms have a formal charge of zero, is the better choice for the Lewis structure of CH2O. Think About ItFor a molecule, formal charges of zero are preferred. When there are nonzero formal charges, they should be consistent with the electronegativities of the atoms in the molecules. A positive formal charge on oxygen, for example, is inconsistent with oxygen’s high electronegativity.

47. 3.5 •• •• •• •• •• •• •• •• = − O O O − = O O O •• •• •• •• Resonance A resonance structureis one of two or more Lewis structures for a single molecule that cannot by represented accurately by only one Lewis structure. Resonance structures are a human invention. Resonance structures differ only in the positions of their electrons.

48. Worked Example 3.7 High oil and gasoline prices have renewed interest in alternative methods of producing energy, including the “clean” burning of coal. Part of what makes “dirty” coal dirty is its high sulfur content. Burning dirty coal produces sulfur dioxide (SO2), among other pollutants. Sulfur dioxide is oxidized in the atmosphere to form sulfur trioxide (SO3), which subsequently combines with water to produce sulfuric acid – a major component of acid rain. Draw all possible resonance structures of sulfur trioxide. StrategyFollowing the steps for drawing Lewis structures, we determine that a correct Lewis structure for SO3 contains two sulfur-oxygen single bonds and one sulfur-oxygen double bond. But the double bond can be put in any one of three positions in the molecule.

49. Worked Example 3.7 (cont.) Solution Think About ItAlways make sure that resonance structures differ only in the position of the electrons, not in the positions of the atoms.

50. 3.6 H Be H Exceptions to the Octet Rule Exceptions to the octet rule fall into three categories: 1) The central atom has fewer than eight electrons due to a shortage of electrons. • Elements in group 3A also tend to form compounds surrounded by fewer than eight electrons. • Boron, for example, reacts with halogens to form compounds of the general formula BX3 having six electrons around the boron atom. only 4 total valence electrons in the system