Chapter 7. Chemical Bonding I: Basic Concepts

Transcription

1 Chapter 7. Chemical Bonding I: Basic Concepts Chemical bond: is an attractive force that holds 2 atoms together and forms as a result of interactions between electrons found in combining atoms We rarely deal with isolated atoms Chemical bonds are broken and formed in reactions Properties of substances often determined by bonds Bonding lowers the potential energy

2 How are Chemical Bonds Formed? We use the electronic structure of atoms to predict which of two types of chemical bonds are likely: 1)Ionic (transferring electrons) 2)Covalent (sharing electrons) Octet Rule: When atoms bond, they lose, gain, or share e - to attain filled outer level of eight e -

3 Ionic Bonds - Transfer of electrons from metal to nonmetal to form ions that come close together in sold ionic compound - Electrostatic attractions of closely packed, oppositely charged ions in a regular 3-D array - Formed when metal (that easily loses electrons) reacts with non-metal (that easily gain electrons). - No. e - lost by metal = No. e - gained by nonmetal

4 Electron Arrangements and Ion Charge Metals form cations by losing enough e - to form same configuration as previous noble gas Nonmetals form anions by gaining enough e - to from same configuration as next noble gas Completely fill outer s & p orbitals Atom Atoms Electron Config Ion Ions Electron Config Na [Ne]3s 1 Na +1 [Ne] Mg [Ne]3s 2 Mg +2 [Ne] Al [Ne]3s 2 3p 1 Al +3 [Ne] O [He]2s 2 2p 4 O -2 [Ne] F [He]2s 2 2p 5 F -1 [Ne]

5 Properties of Ionic Compounds All solids at room temperature are brittle, hard, & rigid Melting points greater than 300 C Liquid state conducts electricity, solid state does not If soluble in water then good electrical conductor Chemical formula is empirical formula - simply giving the ratio of ions based on charge balance (no separate molecules) Ions arranged in a pattern called a crystal lattice maximizes attractions between + and ions

6 Electrostatic forces and the reason ionic compounds crack

7 Electrical conductance and ion mobility Solid ionic compound Molten ionic compound Ionic compound dissolved in water

8 Covalent Bonds Atoms bond by sharing pairs of electrons Commonly found between nonmetal atoms e.g. Cl 2, O 2, N 2, H etc 2 Example H 2 : overlap of 1s orbitals 2H(g) H 2 (g) H = -432 kj Most atoms form covalent bonds by sharing enough electrons to satisfy octet rule

9 Bond Polarity Covalent bonding between unlike atoms results in unequal sharing of the electrons One end has larger electron density than other The result is bond polarity End with larger e - density gets partial - charge End that is e - deficient gets partial + charge δ+ H F δ Dipole Moment (µ)

10 Electronegativity Relative ability of bonded atom in a molecule to attract shared electrons Larger electronegativity: atom attracts more strongly Values 0.7 to 4.0 across period (left to right) on Periodic Table down group (top to bottom) on Periodic Table Larger difference = more polar bond negative end toward more electronegative atom

11 Electronegativity values for selected elements

12 Electronegativity and Bond Type Difference in electronegativity ( EN) provides a good estimate of bond type: 0 to 0.4 non-polar covalent polar covalent > 1.7 ionic

13 Comparison of Ionic vs Covalent Property NaCl CCl 4 State solid liquid Melting point ( C) Molar heat of fusion (kj/mol) Molar heat of vaporization (kj/mol) Electrical conductivity Good Poor

14 Covalent Bonds Single Bond: atoms share 2 e - (1 pair) eg., C-C bond order = 1 Double Bond: atoms share 4 e - (2 pair) eg., C=C bond order = 2 Triple Bond: atoms share 6 e - (3 pair) eg., C C bond order = 3

15 Covalent Bonds Bond Strength: Triple > Double > Single For bonds between same atoms: C N > C=N > C N Bond Length: Single > Double > Triple For bonds between same atoms: C N > C=N > C N

16 Gilbert N. LEWIS The Wayne Gretzky of chemical bonding. American chemist: Gilbert N. Lewis ( )

17 Lewis Structure (Electron Dot Symbols) - Shows how valence electrons are arranged among atoms in molecules & ions - Use symbol of element to represent nucleus & inner e - - Use dots around symbol to represent valence electrons - Reflects idea that stability of a compound relates to noble gas electron configuration - Octet Rule: Elements tend to acquire e - configuration like that of noble gases

18 Writing Lewis structures of molecules 1) Count total number of valence electrons from all atoms (add or subtract if an ion). 2) Choose the central atom. Hint: - Central atoms have the lowest electronegativity. Hydrogen is always terminal while carbon is central. 3) Attach atoms together with one pair of electrons & subtract this total. Hint: - Line is often used for bonding electrons 4) Arrange remaining electrons in pairs: Hint: For hydrogens, a max of 2 e - (duet) & all others atoms have 8 e - (octet) in total around them.

19 Lewis Symbols Al [Ne] 3s 2 3p 1 How many valence e - s? 3 Al F [He] 2s 2 2p 5 F

20 electron transfer IONIC BONDING e.g. NaCl The formation of ionic bonds is represented in terms of Lewis symbols. Na x + Cl [Na] + [ xcl ] Complete transference of electron to Cl - anion 20

21 COVALENT BONDING electron sharing Atoms go as far as possible toward completing their octets by sharing electron pairs Consider F 2 xx x x F x xx xx + F x x F x xx F 21

22 Lewis Structures Practice Write Lewis structures for the following: Atoms: Na, Cl, Ar Ionic compounds: Covalent compounds: NaCl, CaF 2 CH 4, C 2 H 4, C 2 H 2, CCl 2 F 2, N 2, CH 4 O, H 2 NOH

23 Building Lewis structures of molecules HCN as an example... Step 1.Count the total number of valence electrons H has 1 C has 4 N has 5 Total of 10 Step 2. Place one e - pair between each BONDED atom H C N We have 6 e - left All atoms must have an octet or duet Step 3. Add electrons to terminal atoms first The H OK it has its duet... Next... 23

24 Step 3. Building Lewis structures of molecules Add remaining electrons to terminal atoms first Add 6 electrons in pairs to give the N an octet. H C N Step 4. Add any electrons left over to central atom We have none left! Step 5. Check for an acceptable Lewis Structure Do all atoms have an octet??? IN THIS CASE 24

25 Building Lewis structures of molecules H C N No! Both C and N need an octet.. the C and N have to share more than one pair of e - bring e - pairs from outer N atom to form shared pairs to give C its octet!!!

26 Building Lewis structures of molecules Step 5. Check for an acceptable Lewis Structure bring electron pairs from outer N atom to form shared pairs to give C its octet!!! H C N Still no octet on C Do it again!!!! H C N H C N three electron pairs between the C and N 26

27 Building Lewis structures of molecules three electron pairs between the C and N H C N Lewis (electron dot) structure of HCN There is a triple bond.. Also written H C N Another possible structure is. 27

28 Another structure H N C Lets do the Lewis structure... Step 1. Count the total number of valence electrons C has 4N has 5 H has 1 Total of 10 Step 2. Place one e - pair between each atom H N C Step 3. Place remaining electrons on terminal atoms until their octet complete We get. 28

29 Another structure Step 3. Place remaining electrons on terminal atoms until their octet complete H N C Step 4. No electrons left. Step 5. Check for acceptable Lewis structure. The N does not have an octet... We bring electron pairs from outer C atom to form shared pairs to give N its octet!!! Again we need a triple bond. 29

30 Lewis structure of HNC H N C three electron pairs between the C and N this is called a triple bond.. Also written H N C How can we choose? H C N The octet rule is obeyed!!. FORMAL CHARGE.. 30

31 Formal Charge Helps predict most reasonable arrangement of atoms i.e. helps in writing Lewis structure. Hypothetical charge atom would have if the bonding electrons were shared equally. Difference between no. of valence electrons (V) on the free atom & no. assigned to atom in molecule. Formal charge = V (L + ½S) V= No. valence electrons in the free atom L = No. of nonbonding electrons S = No. of shared electrons

32 Formal Charges Formal charge = V (L + ½S) Formal charges help identify most likely arrangement if more than one is possible All formal charges of zero is best or one with lowest number of nonzero formal charges (closest to zero) (a) H N C (b) H C N FC (H) = 1-(0+1) = 0 FC (H) = 1-(0+1) = 0 FC (N) = 5-(0+4) = +1 FC (C) = 4-(0+4) = 0 FC (C) = 4-(2+3) = -1 FC (N) = 5-(2+3) = 0

33 Using formal charges determine which of these CO 2 structures are the most stable? O C O O C O FC (O-) = 6-(6+1) = -1 FC (O=) = 6-(4+2) = 0 FC (C) = 4-(0+4) = 0 FC (C) = 4-(0+4) = 0 FC ( O) = 6-(2+3) = +1 FC (=O) = 6-(4+2) = O C O O C O

34 Lewis Structures Practice (a) Assign formal charges for the following molecules: S O O O N O (b) Write Lewis structures for the following compounds: PCl 3, NO 2-, PO 4 3-, CO 3 2-

35 QUESTION Oxygen difluoride is a powerful oxidizing and fluorinating agent. Select its Lewis structure (a) F O F (b) F O F (c) F O F (d) F O F (e) None of these 35

36 Resonance Structures More than one Lewis structure that differs only in position of e - s Lone pairs & multiple bonds in different positions Actual molecule is combination (or blending) of all resonance forms (delocalized) Actual structure is average of resonance structures O S O.... O S O

37 RESONANCE We use a double headed arrow between the structures.. O O O N N N O O O O O O The electrons involved are said to be DELOCALIZED over the structure. The blended structure is a RESONANCE HYBRID 37

38 QUESTION Which of the following molecules exhibit resonance? 1 CO 2 2 ClO 3-3 O 3 4 Cl 2 CO 5 F 2 O 38

39 Exceptions to the Lewis Structure Rules Some covalent molecules have central atoms that do not have noble gas configuration because: 1) Odd No. of valence electrons (unpaired e - ) 2) Central atom has fewer e - than needed for a noble gas configuration 3) Central atom has more e - than needed for a noble gas configuration

40 1. Odd No. of valence electrons (unpaired e-) Molecules that contains one or more unpaired electrons are paramagnetic or radicals. Elements in 2 nd period have only 4 orbitals in valence shell and can t have > 8 e - around them. Try to write the Lewis structure for NO N O

41 2. Electron-deficient Few molecules contain central atoms that don t have filled valence shell (incomplete octet) B & Be are often octet deficient with outer atoms of hydrogen or other atoms that do not readily form multiple bonds (BeH 2 ) These are very reactive

42 3. Molecules with Extra Electrons Elements in 3 rd & higher periods (n > 2) have more than 4 valence orbitals & can share more than 4 pairs of electrons with other atoms Expanded octet: Empty d orbitals available & able to expand to 10, 12 or more e - Examples include PCl 5, SF 6, ICl 5

43 Elements in rows 3 and following can exceed the octet rule: SF 6 F F F S F F F F F F S F F F SF 6 looks like this. F F F S F F How do we get Lewis Structure??? F 43

44 Chapter 8: Shapes of Molecules 3-D arrangement of atoms in molecule Structure is described by bond angle & bond distance Bond angle: angle between any 2 bonds that includes common atom (degrees) Bond distance: distance between nuclei of 2 bonded atoms (Å or pm) Look at formaldehyde (H 2 CO)

45 VSEPR Theory Valence shell electron-pair repulsion (VSEPR) theory is used to predict molecular geometry by examining no. of bonds & unshared electron pairs KEY: Most stable arrangement is one where valence electrons around central atom are as far away from each other as possible Minimizes repulsions

46 Predicting Molecular Geometry Shape around central atom(s) can be predicted by assuming that areas of electrons (bonding & nonbonding) on central atom will repel each other. Each bond counts as 1 area of electrons. single, double or triple all count as 1 area Each lone pair counts as 1 area of electrons Even though lone pairs are not attached to other atoms, they occupy space around central atom Look at CO 2

47 Bonds (shared e - s) and lone pairs of e - s are as far away from each other as is possible O C O Electron-group geometry is linear (AX 2 ) where A is the central atom and X are the terminal atoms H H C H H The electron-group repulsions force the groups as far apart as possible Tetrahedral AX 4

48 H H N H NH 3 has 4 electron groups and has a tetrahedral electrongroup geometry but it s actual molecular shape (molecular geometry) is not tetrahedral but rather trigonal-pyramidal VSEPR notation - AX 3 E where E is the lone pair of electrons Angle of Electron-group geometry deals with the distribution of the electron groups Molecular geometry deals with the molecular shape of the molecule

49 Electron Group Geometry & Molecular Geometry The electron-domain geometry is often not the shape of the molecule, however. The molecular geometry is defined by the positions of only the atoms in the molecules, not the nonbonding pairs.

50 Table 8.1

51 Electron-Pair Geometry Trigonal bipyramidal (5 regions) Equatorial or axial AX 5 (trigonal bipyramidal 90, 120 ) AX 4 E (seesaw 90, 120 ), AX 3 E 2 (T-shaped 90 ), AX 2 E 3 (linear 180 ) hosphorus%20pentafluoride.jpg

52 Octahedral (6 regions) AX 6 (octahedral, 90 ) - AX 5 E (square pyramidal 90 ) - AX 4 E 2 (square planar 90 ) m/boron%20pentafluoride% 20cloud%20pred.jpg

53 VSEPR All regions of high electron density are not same Certain high density electron areas want more room. Lone pair e - s generally spread out more than bonding e - s, affects bond angle structure looks a little different than expected. Order of repulsive forces: lone pair-lone pair > lone pair-bond pair > bond pair-bond pair

54 Why is the top structure incorrect?

55 Electron-pair geometry vs molecular geometry Electron-pair geometry: Includes all electron pairs (this is what we just looked at) Molecular geometry: Includes only placement of atoms in molecule Same when there are no unshared electron pairs around central atom Look at methane (CH 4 ) vs ammonia (NH 3 )

56 Molecular Geometry Trigonal Bipyramidal: 2 distinct positions (i) axial: smaller (ii) equatorial: larger Unshared pairs always occupy equatorial positions Linear 2 areas of electrons around central atom, both bonding Or two atom molecule is a trivial case

57 Predicting a VSEPR Structure 1. Draw Lewis structure. 2. Count number of regions of high e - density (unshared pairs and bonds) around central atom. 3. Identify electron-pair geometry. 4. If more than one arrangement is possible, choose one that minimizes unshared pair repulsions.

58 Multiple Covalent Bonds Predict the electron geometry and molecular geometry of SO 2 S S O O O O

59 Molecules With More Than One Central Atom What is the electron-group geometry of methyl isocyanate, CH 3 NCO? Ans: Draw the best fit Lewis structure Valence e - s: C = 8 N = 5 O = 6 H = 3 Total = 22 H H C H H N C O H C H N C O

60 H N C H 180 o C O 120 o H 109 o

61 Practice Predict electron-pair & molecular geometries for following: SiCl 4, H 3 O +, SF 4, XeF 4, CH 3 OH, NH 2 CH 2 COOH