Chapter 6. The Chemical Bond

Transcription

1 Chapter 6 The Chemical Bond

2 Some questions Why do noble gases rarely bond to other elements? How does this relate to why the atoms of other elements do form bonds? Why do certain elements combine to form ionic and others molecular compounds? Why is the formula of water H 2 O and not H 3 O or HO 2 2

3 Bond formation and noble gases Noble gases rarely form compounds They have filled s and p outer subshells This is a total of eight electrons, referred to as an octet Eight electrons in the outer s and p orbitals is a particularly stable configuration The energy required to remove an electron from these full subshells is particularly high 3

4 Lewis dot symbols for elements Since only the valence electrons are involved in bonding we can concentrate on those Lewis dot symbols are used to represent the valence electrons of an atom 4

5 How atoms achieve an octet Metals can lose one to three electrons to form a cation with the electron configuration of the previous noble gas Nonmetals can gain one to three electrons to form an anion with the electron configuration of the next noble gas Atoms can also share electrons 5

6 Formation of ions Metals can lose electrons to form ions Na ([Ne]2s 1 ) Na + ([Ne]) + e - If a metal loses all of its outer electrons, it acquires the octet of the previous noble gas Nonmetals can gain electrons to form ions Cl ([Ne]2s 2 2p 5 + e - Cl ([Ne]2s 2 2p 6 Lewis dot structures of the atoms can be very helpful here 6

7 Forming ionic compounds Reaction of Na with Cl Na donates an electron to Cl Na + has the previous noble gas structure (Ne) Cl - has the next noble gas structure (Ar) 7

8 Binary ionic compounds In NaCl, each Na + is surrounded by six Cl -, and each Cl - is surrounded by six Na + Ionic lattice is a three-dimensional array of ions These electrostatic attractions are called ionic bonds 8

9 But remember! Atoms can also share electrons - when this happens, the electrons form covalent bonds 9

10 Chemical bonds (let s review) A chemical bond is the force that holds two or more atoms together Chemical bonds involve the electrons A bond results if a more stable electron configuration results The valence electrons are the electrons in the outer s and p subshells 10

11 The Covalent Bond Covalent bonds result from electron sharing between two atoms We use Lewis dot structures to show the order and arrangement of the atoms in a molecule and all of the valence electrons H. + H. H : H 11

12 The Octet Rule The s block and p block elements (often termed the representative elements) will form bonds such that there are eight electrons surrounding each atom (the octet rule) Obtaining this configuration is the driving force for bond formation for many compounds formed by the representative elements The exceptions are H, Li and Be, which tend to follow a duet rule (filling the ns subshell) 12

13 Types of covalent bonds Two nonmetals can share one, two or three electron pairs The bonds resulting from this sharing are referred to as single, double or triple bonds respectively Multiple bonds are frequently observed in compounds of 2nd period elements 13

14 Writing Lewis dot structures The octet rule and Lewis dot structures allow us to justify the formulas that we know We can also predict the formulas of new compounds 14

15 Lewis dot structures 1. Count the number of valence electrons for the atoms in the molecule, include charge. 2. Place the most electropositive atom in the center (the inner atom). Draw simple diagram; single atom in middle, one bond to each outer atom. 3. Subtract two electrons from the total number of valence electrons for each bond. 4. Satisfy octet rule for all outer atoms. 5. Subtract these atoms from total. 6. Place remaining pairs on center atom. 7. Check octet rule for center atom. 8. CELEBRATE! 15

16 Lewis Dot Diagrams Examples: CF 4 ; NH 3 ; NH 4+ ; SCl 2 ; PF 3 Other examples: NO 2- ; 16

17 Resonance structures In compounds with multiple bonds, sometimes you can draw structures which vary only by placement of the double bonds The structures are called resonance structures, and are an approximation of the true structure of the molecule Actually, the molecule is a superposition of all of the resonance structures 17

18 Nitrate ion Nitrate has three resonance structures Each is identical except for the placement of the double bond and associated lone pairs Experimentally, all N-O bonds are identical 18

19 More examples Lewis dot diagrams NO 2 -- A bit more complicated: C 2 H 6 ; (hydrocarbons in general =C 3 H 8 ); halogens!) H 2 O 2 ; H 2 CO, BF 3 (no multiple bonds for Start to anticipate base/acid chemistry: NH 3, OH -, etc 19

20 Electronegativity the ability of an atom to attract electrons in a bond to itself Differences in electronegativities of atoms that are bonded together results in a partial transfer of electron charge to the more electronegative atom. The bond is therefore a polar covalent bond The polar bond has a negative end and a positive end (a so-called dipole; which we indicate with a δ with the appropriate sign added) 20

21 Nonpolar, polar and ionic bonds 21

22 Polarity of bonds Bonds that involve atoms of differing electronegativities have a concentration of negative charge at the more electronegative atom, and a deficiency of charge at the less electronegative atom This unequal distribution of negative charge creates a dipole, where one end of the bond is slightly negative and the other is slightly positive 22

23 Geometry of Simple Molecules Electron pairs will repel each other, and will govern the structure of the molecule, all other things being equal Electron pairs will arrange themselves to be as far apart as possible Note that molecular geometry is described by the bonded atoms and does not include the lone pairs 23

24 Electron pair repulsion Degree of repulsion depends on the electron pair types; in order of decreasing repulsion lone pair-lone pair lone pair-bonding pair bonding pair-bonding pair We will also treat all of the electrons that bond together two atoms as one electron group regardless of whether the bond is single, double or triple 24

25 Parent structures We will consider three parent structures to begin The central or inner atom is designated A, outer atoms are designated X, and lone pairs are designated E The parent structures are based on the number of electron pairs that surround the central atom 25

26 Parent structures Parent structure Name Bond angles AX 4 tetrahedron AX 3 trigonal planar 120 AX 2 linear

27 Parent structure derivatives Each parent structure can give rise to a family of derivatives, simply by replacing bonding pairs with lone pairs For AX 4, there are two derivatives AX 3 E - the trigonal pyramid (NH 3 ) AX 2 E 2 - bent (H 2 O) Replacement of the bonding pairs with the lone pair(s) compresses the bond angles 27

28 Tetrahedral - AX 4 Each molecule has four electron pairs around them Replacement of a bonding pair with a lone pair yields the AX 3 E (NH 3 ), AX 2 E 2 (H 2 O) 28

29 Trigonal - AX 3 The examples shown here have three electron pairs around the central atom Note that when one of the bonding pairs is replaced by a lone pair, the bond angle is smaller Structure is designated by AX 2 E, where E is a lone pair electron deficient 29

30 Linear examples - AX 2 The bonding pairs in the following molecules arrange themselves to be as far apart as possible These examples have two electron pairs 30

31 Expanded Valence Shells Valence Shell Electron Pair Repulsion (VSEPR) 3 rd Row and beyond, valence shell MAY have more than octet up to 12 electrons, 6 electron pairs. Use VSEPR chart from website: /chem101/lecture supplements/vsepr1.p DF 31

32 Examples SCl 4 ; KrCl 2 ; BrF 4 IF 5 ; SF 6 ; XeBr 4 32

33 Polarity of bonds can lead to polar molecules 33

34 Polarity of molecules If the forces are equal but opposite, the polarity of the bonds cancel Such molecules have polar bonds, but are themselves nonpolar (e. g. CO 2 ) 34

35 Nonpolar molecules 35

36 Polar molecules Polar molecules result from an arrangement of polar bonds such that the entire molecule has a dipole The polar bonds can be arranged to cancel the polarity (CO 2 ) The best way to predict this in complex cases is vector algebra, but it is important to learn to recognize this in clear cut structures 36

37 Polar molecules If the two bonds are not equally polar, then a net dipole exists (a) If the two bonds are equally or not equally polar, yet are not opposite, a net molecular dipole exists 37

38 Linear example The Be-Cl and Be-H are both polar, but not to the same extent Even though BeClH is linear, the polarity of the bonds are not equal so the molecule is polar 38

39 Noncanceling bonds Water is bent, so the dipolar O-H bonds cannot cancel each other Same for SO 2 or any bent molecule 39

40 That s all folks! On to Chapter 10 40