Chapter 10: Molecular Structure and Bonding Theories
|
|
- Theodora Dalton
- 6 years ago
- Views:
Transcription
1 hapter 10: Molecular Structure and Bonding Theories 10.1 See Section The main premise of the VSEPR model is that the electron pairs within the valence shell of an atom repel each other and determine the molecular geometry of the molecule or ion of interest See Section ormaldehyde,, is an example of a trigonal planar molecule in which the carbon forms four bonds, one of which is part of a double bond to the oxygen atom. Phosgene, l, is another example of a trigonal planar molecule in which carbon forms four bonds. is an example of a molecule with a central atom that makes four bonds but has a linear-bonded atom lone-pair arrangement See Section 10.1 and igure 10.6 I I The VSEPR model states that the order of importance of repulsions within the valence shell of an atom is lp-lp > lp-bp > bp-bp. In addition, repulsions diminish as the angle between the electron pairs increases from 90º to 10º to 180º. ence, lone pairs always occupy equatorial rather than axial positions in trigonal bipyramidal arrangements because there are only two interactions at 90º for equatorial positions in trigonal bipyramidal positions compared to three interactions at 90º for axial positions. The structure on the right is favored and is observed experimentally See Section 10. and Examples 10.4, Any molecule with a totally symmetrical arrangement of atoms and lone pairs of electrons is nonpolar. ence, any of the electron-pair arrangements shown in igure 10.1 having like atoms attached to a central atom that is different has polar bonds but is nonpolar overall. or example, Bel (g) as is nonpolar because the bond dipoles cancel each other. This always occurs when the molecule is totally symmetrical. I l Be l 10.9 See Section Valence bond theory describes bonds as being formed by atoms sharing valence electrons in overlapping valence orbitals. These overlaps are caused by the attraction between the nuclear charge of one of the bonded atoms and the electron cloud of the other atom and vice versa. 133
2 10.11 See Section 10.3 and Example l [e] 3s 3p [e] s p The partially filled 3p orbital of l overlaps with the partially filled p orbital of to form the bond in l See Section 10.1, 10.3, igure 10.15, 10.16, and Example l Bl 3 B S (steric number for B = 3, trigonal planar electron-pair arrangement, sp l l hybrids for B. l Bl 4 l B l l S for B = 4, tetrahedral electron-pair arrangement, sp 3 hybrids for B See Section 10.3 and Table l ybrid orbitals are predicted by looking at Lewis structures, observing the steric number, predicting the electron-pair arrangement and selecting the corresponding hybrid orbitals for the central atom. owever, we can only determine the positions of attached atoms and are unable to determine the positions of lone pairs. ence, any prediction of use of hybrid atomic orbitals by l in l would be a matter of pure conjecture that could not be verified by experiment See Section 10.4 and igures 10., A sigma (σ) bond is a bond in which the shared pair of electrons is symmetric about the axis joining the two nuclei of the bonded atoms. A pi (π) bond is a bond that places electron density above and below the line joining the bonded atoms and can be formed by the sideways overlap of p orbitals. p z p z σ p y p y π 134
3 10.19 See Section 10.5 and igures 10.35, σ* 1s 1s 1s σ 1s 10.1 See Section 10.6 and igure According to molecular orbital theory, there is a weak interaction between the s orbital of Li and the p orbital of pointing toward Li resulting in a sigma bonding molecular orbital. This orbital is mainly centered around the fluorine atom, and the valence electron of Li is nearly completely transferred to the fluorine atom. According to the ionic bonding description, the valence electron of lithium is completely transferred from Li to giving Li + and. ence, the two descriptions differ slightly in terms of electron transfer from Li to See Section 10.1 and igure (a) S = 3, trigonal planar (b) S = 4, tetrahedral (c) S = 4, tetrahedral (d) S = 5, trigonal bipyramidal Remember that bonded-atom lone-pair arrangement takes into account both the outer atoms bound to the central atom and any lone pairs present on the central atom See Section 10.1, igures 10.1, 10.6, and Examples 10.1, 10.. (a) 4 (b) S (c) As 5 S S S = 4 tetrahedral S =, linear S = 5, trigonal bipyramidal (d) (e) As S = 3, trigonal planar S = 4, tetrahedral 135
4 10.7 See Section 10.1, igures 10.1, 10.6, and Examples 10.1, 10.. (a) Se (b) (c) 3 + Se S = 3, trigonal planar S =, linear S = 4, tetrahedral shape: bent shape: linear shape: trigonal pyramidal (d) I 5 (e) Sl 4 I S = 6, octahedral shape: square pyramidal S l l l l S = 5, trigonal bipyramidal shape: see - saw 10.9 See Section 10.1, igures 10.1, 10.6, and Examples 10.1, (a) Bl 3 l 3 (b) S 6 l l B l S l l l S = 3 S = 4 S = 4 S = 6 trigonal planar tetrahedral tetrahedral octahedral electron pair electron pair electron pair electron pair arrangement arrangement arrangement arrangement 10º bond angles 109º bond angles 109º bond angles 90º & 180º bond angles l 3 has smaller bond angles than Bl 3. S 6 has smaller bond angles than See Section 10.1, igures 10.1, 10.6, and Examples 10.1, 10.. (a) l + 4 (b) S I 4 + l l S I S = 4 S =4 S = 4 S = 6 tetrahedral electron tetrahedral electron tetrahedral electron octahedral electron pair arrangement pair arrangement pair arrangement pair arrangement 136
5 109º bond angles 109º bond angles 109º bond angles 90º & 180º bond angles l has slightly smaller bond angles due to the lone pair on the. I 4 has smaller bond angles than S See Section 10.1, igures 10.1, 10.6, and Examples 10., (a) 3 (b) Br (c) 3 Br Br 180 Þ180 Þ 10 Þ 10 Þ See Section 10.1, igures 10.1, 10.6, and Examples 10., (a) S (b) l 3 (c) S S 10 Þ expanded valence shell l expanded valence shell S 180 Þ See Section 10.1, igures 10.1, 10.6, and Examples 10., (a)( 3 ) 10 Þ (b) Þ 137
6 10.39 See Section 10.1, igures 10.1, 10.6, and Examples 10., (a) 3 l l l b) P P 180 Þ 180 Þ See Section 10.1, igures 10.1, 10.6, and Examples 10., (a) 4 (b) 3 6 (c) Xe Xe 10 Þ 180 Þ 10 Þ See Section 10.1, igures 10.1, 10.6, and Examples 10., (a) 3 S (b) S 10 Þ See Section 10. and Examples 10.4, (a) 4 (b) S (c) As 5 (d) As S S symmetrical symmetrical symmetrical unsymmetrical nonpolar nonpolar nonpolar polar See Section 10. and Examplse 10.4, (a) Se (b) (c) Sl 4 l Se l S l l 138
7 expanded valence shell expanded valence shell unsymmetrical unsymmetrical unsymmetrical polar polar polar See Section 10. and Example (a) (b) I (c) I I unsymmetrical symmetrical unsymmetrical polar nonpolar polar See Section 10. and Example (a) 3 (b) Br 4 (c) BeI Br Br Br Br unsymmetrical symmetrical symmetrical polar nonpolar nonpolar I Be I See Section 10. and Example (a) (b) symmetrical unsymmetrical nonpolar polar bond dipoles cancel bond dipoles do not cancel See Section (a) 10º, sp or sp 3 d (b) 90º, sp 3 d or sp 3 d (c) 180º, sp, sp 3 d or sp 3 d hybrids hybrids hybrids See Section 10.3 and Examples 10.7, (a) 4 (b) Sbl 6 (c) As 5 (d) Si 4 l l l Sb l l l S = 4 S = 6 S = 5 S = 4 tetrahedral octahedral trigonal bipyramidal tetrahedral sp 3 for sp 3 d for Sb sp 3 d for As sp 3 for Si As Si 139
8 (e) S = 4 tetrahedral sp 3 for See Section 10.3 and Examples 10.7, (a) (b) Snl (g) (c) I 3 (d) Se l Sn I I I Se l S = S = 3 S = 5 S = 3 linear trigonal planar trigonal bipyramidal trigonal planar sp for central sp for Sn sp 3 d for central I sp for Se See Section 10.3 and Examples 10.7, (a) (b) (c) S = 3 S = 4 S = 3 trigonal planar tetrahedral trigonal planar sp for sp 3 for sp for See Section 10.3 and Examples 10.7, (a) 3 + (b) 3 (c) l + l l S = 4 S = 4 S = 4 tetrahedral tetrahedral tetrahedral sp 3 for sp 3 for sp 3 for 140
9 10.65 See Section 10.3 and Examples 10.7, (a) S = 4, tetrahedral sp 3 for and [e] for. s p A sp 3 orbital containing one electron overlaps with a p orbital from containing one electron to form an - bond in. The lone pairs of electrons are in sp 3 orbitals of. (b) 3 S = 4, tetrahedral sp 3 for and for. 1s A sp 3 orbital from containing one electron overlaps with a 1s orbital of containing one electron to form a - bond in 3. The lone pair of electrons is in a sp 3 orbital of. (c) Bl 3 S = 3, trigonal planar l l B l sp for B and [e] 3s 3p for l. A sp orbital from B containing one electron overlaps with a 3p orbital of l containing one electron to form a B-l bond in Bl See Section 10.3 and Examples 10.7, Se 4 Se S = 5, trigonal bipyramidal sp 3 d for Se and [e] s p for. A sp 3 d orbital from Se containing one electron overlaps with a p orbital of containing one electron to form a Se- bond in Se 4. The lone pair of electrons is in a sp 3 d orbital of Se See Section 10.3 and Examples 10.7, (a) or : S = 4, tetrahedral, sp 3. l or : S = 4, tetrahedral, sp 3. l or : S = 4, tetrahedral, sp 3. (b) or both : S =, linear, sp. P or P: S = 4, tetrahedral, sp
10 10.71 See Section 10.4, igures (a) p z p z σ (b) p y p y π (c) sp z p z σ See Section 10.3, 10.4, and Examples 10.7, 10.8, Bond rbital verlaps Bond Type - sp 3-1s σ 1 - sp 3 -sp σ α sp-p z σ p y -p y π p x -p x π S 1 = 4, tetrahedral, sp 3 S =, linear, sp There are a total of five σ bonds and two π bonds in See Section 10.3, 10.4, and Examples 10.7, 10.8, π σ Bond rbital verlaps Bond Type - sp -1s σ - sp -1s σ = sp -sp σ p x -p x π S = 3, trigonal planar, S = 3, trigonal planar, sp Bond overlaps are similar to those shown for 4 in igure σ 14
11 10.77 See Section 10.3, 10.4, and Examples 10.7, 10.8, (a) (b) l The single bonded carbon atoms are tetrahedral and sp 3. The double bonded carbon atoms are trigonal planar and sp. The is trigonal planar and sp. l (c) 1 The is tetrahedral and sp 3. The 1 is tetrahedral and sp 3. The is trigonal planar and sp. The of -- is tetrahedral and sp See Section 10.3, 10.4, and Examples 10.7, 10.8, sp for sp for See Section 10.3, 10.4, and Examples 10.7, 10.8, (a) is tetrahedral and sp 3. is tetrahedral and sp 3. (b) is tetrahedral and sp 3. and 3 are linear and sp. 143
12 10.83 See Section 10.3, 10.4, and Examples 10.7, 10.8, is linear and sp See Section 10.5, igure 10.37, and Example is trigonal planar and sp. is trigonal planar and sp. σ* 1s 1s 1s σ 1s e + + e e + or e + the electron configuration is (σ 1s ) there are no unpaired electrons, the bond order is 1 [ 0] = 1, and it is predicted to be stable See Section 10.5, igures 10.37, Table 10., and Example σ* s s s σ s Li Li Li or Li the electron configuration is (σ s ) there are no unpaired electrons, the bond order is 1 [ 0] = 1, and it is predicted to be stable See Section 10.5, igure 10.40, and Examples 10.11, (a) + has ( ) = 7 valence electrons: (σ s ) (σ* s ) (π p ) 3. The bond order is 1 [5 ] = 1.5, and there is one unpaired electron in a π p orbital. (b) has ( ) = 11 valence electrons: (σ s ) (σ* s ) (π p ) 4 (σ p ) (π* p ) 1. The bond order is 1 [8 3] =.5, and there is one unpaired electron in a π* p orbital. (c) Be has ( + + 1) = 5 valence electrons: (σ s ) (σ* s ) (π p ) 1. The bond order is 1 [3 ] = 0.5, and there is one unpaired electron in a π p orbital. 144
13 10.91 See Section 10.5, igure 10.40, and Examples 10.11, has (5 + 5) + 10 valence electrons: (σ s ) (σ * s) (π p ) 4 (σ p ). The bond order is σ 1 [8 ] = 3.0. has (5+5+1) = 11 valence electrons: (σ s ) (σ * s) (π p ) 4 (σ p ) (π * p) 1. The bond order is 1 [8 3] =.5. has a higher bond order than because the additional electron in orbital. occupies an antibonding See Section 10.5, igure 10.40, and Examples 10.11, (a) B has (3+3) = 6 valence electrons: (σ s ) (σ * p) (π p ). The bond order is 1 [4 ] = 1.0. B has (3+3+1) = 7 valence electrons: (σ s ) (σ * p) (π p ) 3. The bond order is 1 [5 ] = 1.5. B has a higher bond order and stronger bond than B because the additional electron in B occupies a bonding orbital. (b) has (4+4+1) = 9 valence electrons: (σ s ) (σ * s) (π p ) 4 (σ p ) 1. The bond order is 1 [7 ] =.5. + has (4+4 1) = 7 valence electrons: (σ s ) (σ * s) (π p ) 3. The bond order is 1 [5 ] = 1.5. has a higher bond order and stronger bond than +. orming from involves adding an electron to a bonding orbital, whereas forming + from involves removing an electron from a bonding orbital. (c) + has (6+6 ) = 10 valence electrons: (σ s ) (σ * s) (π p ) 4 (σ p ). The bond order is 1 [8 ] = 3.0. has (6+6) = 1 valence electrons: (σ s )(σ * s)(σ p ) (σ * p). The bond order is 1 [8 4] =.0. + has a higher bond order and stronger bond than because two electrons are removed from antibonding orbitals of to form
14 10.95 See Section 10.5 and Exercises 10.89, 10.91, (a) and + have 13 valence electrons. A bond order of 1 [8 5] = 1.5 indicates these species should be stable. (b),, and + have 10 valence electrons. A bond order of 1 species should be stable. [8 ] = 3.0 indicates these (c) Li, Be, and B + have 4 valence electrons. A bond order of 1 [ ] = 0 indicates these species would not be stable See Section 10.5 and igure (a) and have 10 valence electrons and (σ s ) (σ * s) (π p ) 4 (σ p ). (b) B and Be have 6 valence electrons and (σ s ) (σ * s) (π p ) See Section 10.5 and igure umber of umber of Species valence electrons Electron onfiguration Bond rder unpaired electrons (a) 9 (σ s ) (σ * s) (π p ) 4 (σ p ) 1 1 [7 ] =.5 1 in σ p (b) 11 (σ s ) (σ * s) (π p ) 4 (σ p ) (π * p) 1 1 [8 3] =.5 1 in π* p (c) BeB 6 (σ s ) (σ * s) (π p ) 1 [4 ] = 1.0 in π p (d) B + 6 (σ s ) (σ * p) (π p ) 1 [4 ] = 1.0 in π p See Section 10.5 and igure σ p p π p p σ p π p The bond order for is 1 [8 6]=1. s σ* s s σ s
15 See Section 10.6 and igure is isoelectronic with 3. The delocalized π molecular orbital is formed by the p orbitals of and that are perpendicular to the plane of the atoms. It looks like the delocalized π molecular orbital shown for 3 in igure See Section 10.1, 10.3, 10.4, igures 10.15, 10.17, and Example Each is trigonal planar with 10º bond angles and sp. The central is tetrahedral with 109º bond angles and sp See Section 10.1, igures 10.1, 10.6, and Examples 10.1, 10., l l + + l l < l = 180º < l = 109º See Section 10.3, 10.4, and Examples 10.8,
16 (a) There are thirty-nine sigma bonds and six pi bonds shown in the Lewis structure. The pi bonding in the six-membered carbon ring is actually delocalized pi bonding. (b) The S for each carbon atom forming a double bond with oxygen is three, the electron-pair arrangement is trigonal planar and the hybridization for each of these carbon atoms is sp. (c) The S for each atom is four, the electron-pair arrangement is tetrahedral and the hybridization for each atom is sp See Section 10.1, 10., 10.3, 10.4, and Examples 10.7, 10.8, l The has a S of 3, a trigonal planar electron-pair arrangement and uses sp hybrids. l The molecule is unsymmetrical and therefore polar See Sections 10.1, 10.3, 10.4, and Examples 10.1, Bond 1: S = 3, trigonal planar and therefore 10º. Bond : S = 4, tetrahedral and therefore 109º. Bond 3: S = 4, tetrahedral and therefore 109º. Bond 4: S = 3, trigonal planar and therefore 10º. Bond 5: S = 4, tetrahedral and therefore 109º
17 See Section 10.5 and igure σ p p π p p σ p π p s σ* s s σ s - - σ p p π p p σ p π p s σ* s s σ s + + The bond order for = 1 [8 4] =. The bond order for + = 1 [8 ] = 3. ence, + has a higher bond order than. 149
18 See Section 10.1, 10., 10.3, (a) Total valence electrons = [ 5() + 1() + 6()] = 4. ourteen electrons remain after assigning five single bonds, and sixteen unshared electrons are needed to give each atom a noble gas configuration ( for each and 6 for each ). ence, two electrons (16-14) must be used to form one additional bond. ormal charge considerations indicate the additional bond is formed by and. S of 4 for on the left indicates bond angles of approximately 109º, and S of 3 for on right indicates bond angles of 10º. The on the left uses sp 3 hybrids, and the on the right uses sp hybrids. The molecule is unsymmetrical and polar. (b) 3 Total valence electrons = [1 1() + 3 5()] = 16 Ten electrons remain after assigning three single bonds, and fourteen unshared electrons are needed to give each atom a noble gas configuration (4 for left, 4 for center and 6 for right ). ence, four electrons (14-10) must be used to form two additional bonds. S of 3 for on left indicates bond angles of approximately 10º and S of for in the center indicates bond angles of 180º. The on the left uses sp hybrids and the in the center uses sp hybrids. The molecule is unsymmetrical and polar. ote: is plausible with sp 3 for left and sp for central See Sections 10.1, 10., 10.3, (a) Total valence electrons =[1 6() + 1 4() + 1 5() + 1(charge)] = 16. Twelve electrons remain after assigning two single bonds, and sixteen unshared electrons are needed to give each atom a noble gas configuration (6 for, 4 for, and 6 for ). ence, four electrons (16-1) must be used to form two additional bonds. This gives the following resonance possibilties with lowest formal charges: A B hybrids: sp sp sp Structure B is likely to be the most important because it has the lowest formal charges and places the negative formal charge on rather than, on the more electronegative atom. S of for indicates bond angle of 180º. (b) 3 Total valence electrons =[1 5() + 3 6() + 1(charge)] = 4. Eighteen electrons remain after assigning three single bonds, and twenty unshared 150
19 Electrons are needed to give each atom a noble gas configuration. ence, two electrons (0-18) must be used to form one additional bond. This gives the following resonance possibilities: A B hybrids: sp sp sp These structures are equivalent and equally important. S of 3 for indicates bond angles of 10º See Section 10.1, 10., 10.3, Vitamin-A contains ten sp hybridized carbon atoms. These are the ten carbon atoms that are involved in double bonds. Vitamin-A also contains ten sp 3 hybridized carbon atoms. These are the four carbon atoms that are part of the ring and not part of double bonds, the six that are part of 3 groups and the one that is bonded to oxygen See Section 3.3, 10.1, 10.4, and Example 3.1. Assume the sample has a mass of g and therefore contains g, 9.15 g, 9.15 g, and 36.3 g. 1 mol 4.54 mol? mol = g = 4.54 mol relative mol = =.00 mol 1.01 g.7? mol = 9.15 g 1 mol g? mol = 36.3 g 1 mol 16.0 g 9.08 mol = 9.08 mol relative mol =.7.7 mol =.7 mol relative mol =.7 = 4.00 mol = 1.00 mol 151
20 The simplest formula is 4 ), and the simplest formula molar mass is 44.0 g/mol. molar mass compound 44.0 g / mol n = = = 1, so the molecular formula is also molar mass g / mol 4. The two possible Lewis structures for 4 are: A B 10 Þ 10 Þ 10 Þ ybrids:, sp 3 ;, sp, sp ;, sp ;, sp See Section 10.1, 10., 10.3, S Total valence electron = [ 7() + 6(s)] = 6 S S Twenty electrons remaining after assigning three single bonds and twenty unshared electrons are needed to give each atom a noble gas electron configuration, no unshared electrons are leftover. S S S = 4 for each sulfur atom, therefore hybridization for both S atoms is sp 3. S S Twenty electrons remaining after assigning three single bond and eighteen unshared electrons are needed to give the fluorine atoms and the sulfur atom bonded to the other sulfur atom a noble gas electron configuration, two unshared electrons are assigned to the sulfur atom bonded to fluorine atoms. S S An alternate representation closer to reality is: S S This will indicate sp 3 hybridization in S atom bonded to fluorine atoms, the other S atom is using atomic orbitals, p orbitals to form the covalent bond. Since sulfur can violate the octet rule, a double bond between the two sulfur atoms can be formed to eliminate the positive and negative formal charges on both S atoms. S S This still leaves the central sulfur atom with sp 3 hybridization; the double bond is formed by overlap of a p orbital on the terminal sulfur and a d orbital on the central sulfur atom. 15
21 10.17 See Sections 3., 3.3, 5.4, 10.1, 10.3, To determine the value of x, the mol per mol are determined and then the mol per mol.? mol in = 1.30 g 1 mol mol = mol 6.0 g 1 mol? mol in = 1.30 g 1 mol 6.0 g? mol in 1. L = mol. Known Quantities: P = 1.01 atm V = 1. L Solving PV = nrt for n gives n = PV in 1. L mol 1 mol = mol T = = 300 K RT. n = (1.01 atm) (1. L) L g atm = mol mol g K (300 K) = mol mol = mol 1 mol? mol mol mol.00 mol = = mol mol 1 mol ence, there are 4.00 mol per.00 mol, and the compound is 4. The Lewis structure for 4 is: The carbon-carbon double bond is formed by a sp -sp σ overlap and a p-p π overlap. See igures 10.3, 10.4, 10.5 for illustration of these overlaps. 153
Chemical Bonding II: Molecular Geometry and Hybridization of Atomic Orbitals
Chemical Bonding II: and ybridization of Atomic rbitals Chapter 10 Valence shell electron pair repulsion (VSEPR) model: Predict the geometry of the molecule from the electrostatic repulsions between the
More informationChapter 9. Chemical Bonding II: Molecular Geometry and Bonding Theories
Chapter 9 Chemical Bonding II: Molecular Geometry and Bonding Theories Topics Molecular Geometry Molecular Geometry and Polarity Valence Bond Theory Hybridization of Atomic Orbitals Hybridization in Molecules
More informationChapter 9 Molecular Geometry and Bonding Theories
Lecture Presentation Chapter 9 Geometry James F. Kirby Quinnipiac University Hamden, CT Shapes Lewis Structures show bonding and lone pairs, but do not denote shape. However, we use Lewis Structures to
More informationOrganic Chemistry. Review Information for Unit 1. VSEPR Hybrid Orbitals Polar Molecules
rganic hemistry Review Information for Unit 1 VSEPR ybrid rbitals Polar Molecules VSEPR The valence shell electron pair repulsion model (VSEPR) can be used to predict the geometry around a particular atom
More informationChapters 9&10 Structure and Bonding Theories
Chapters 9&10 Structure and Bonding Theories Ionic Radii Ions, just like atoms, follow a periodic trend in their radii. The metal ions in a given period are smaller than the non-metal ions in the same
More informationMolecular Geometry and Bonding Theories. Chapter 9
Molecular Geometry and Bonding Theories Chapter 9 Molecular Shapes CCl 4 Lewis structures give atomic connectivity; The shape of a molecule is determined by its bond angles VSEPR Model Valence Shell Electron
More information: Bond Order = 1.5 CHAPTER 5. Practice Questions
CAPTER 5 Practice Questions 5.1 5.3 S 5.5 Ethane is symmetrical, so does not have a dipole moment. owever, ethanol has a polar group at one end and so has a dipole moment. 5.7 xygen has the valence electron
More informationCHAPTER 9 COVALENT BONDING: ORBITALS. Questions
APTER 9 VALET BDIG: RBITALS Questions 11. In hybrid orbital theory, some or all of the valence atomic orbitals of the central atom in a molecule are mixed together to form hybrid orbitals; these hybrid
More informationChapter 9. Molecular Geometry and Bonding Theories
Chapter 9. Molecular Geometry and Bonding Theories 9.1 Molecular Shapes Lewis structures give atomic connectivity: they tell us which atoms are physically connected to which atoms. The shape of a molecule
More informationCh. 9- Molecular Geometry and Bonding Theories
Ch. 9- Molecular Geometry and Bonding Theories 9.0 Introduction A. Lewis structures do not show one of the most important aspects of molecules- their overall shapes B. The shape and size of molecules-
More informationReview questions CHAPTER 5. Practice exercises 5.1 F F 5.3
CHAPTER 5 Practice exercises 5.1 S 5.3 5.5 Ethane is symmetrical, so does not have a dipole moment. However, ethanol has a polar H group at one end and so has a dipole moment. 5.7 xygen has the valence
More informationChapter 9: Molecular Geometry and Bonding Theories
Chapter 9: Molecular Geometry and Bonding Theories 9.1 Molecular Geometries -Bond angles: angles made by the lines joining the nuclei of the atoms in a molecule -Bond angles determine overall shape of
More informationChapter 10: Chemical Bonding II: Molecular Shapes; VSEPR, Valence Bond and Molecular Orbital Theories
C h e m i s t r y 1 A : C h a p t e r 1 0 P a g e 1 Chapter 10: Chemical Bonding II: Molecular Shapes; VSEPR, Valence Bond and Molecular Orbital Theories Homework: Read Chapter 10: Work out sample/practice
More informationWhat Do Molecules Look Like?
What Do Molecules Look Like? The Lewis Dot Structure approach provides some insight into molecular structure in terms of bonding, but what about 3D geometry? Recall that we have two types of electron pairs:
More informationChemical Bonding II. Molecular Geometry Valence Bond Theory Phys./Chem. Properties Quantum Mechanics Sigma & Pi bonds Hybridization MO theory
Chemical Bonding II Molecular Geometry Valence Bond Theory Phys./Chem. Properties Quantum Mechanics Sigma & Pi bonds ybridization MO theory 1 Molecular Geometry 3-D arrangement of atoms 2 VSEPR Valence-shell
More informationChapter 10. Structure Determines Properties! Molecular Geometry. Chemical Bonding II
Chapter 10 Chemical Bonding II Structure Determines Properties! Properties of molecular substances depend on the structure of the molecule The structure includes many factors, including: the skeletal arrangement
More information; (c) [Li] [: O :] [Li]. 5a. The electrostatic potential map that corresponds to IF is the one with the most red in it. ... C C H
hapter 10 Answers ractice Examples 1a Mg 1b n, Ge, [: Br :], K, : e: + 2 : : +, [Tl ] +, 2 : : [] 2a (a) [a] [ ] [a] ; (b) [Mg] [: :] [Mg] [: :] [Mg] 2+ 3 2+ 3 2+ 2+ 2b (a) [: I :] [a] [: I :] 2+ 2 ; (b)
More information11/14/2014. Chemical Bonding. Richard Philips Feynman, Nobel Laureate in Physics ( )
Chemical Bonding Lewis Theory Valence Bond VSEPR Molecular rbital Theory 1 "...he [his father] knew the difference between knowing the name of something and knowing something" Richard Philips eynman, Nobel
More informationMolecular Shape and Molecular Polarity. Molecular Shape and Molecular Polarity. Molecular Shape and Molecular Polarity
Molecular Shape and Molecular Polarity When there is a difference in electronegativity between two atoms, then the bond between them is polar. It is possible for a molecule to contain polar bonds, but
More informationDrawing Good Lewis Structures. Molecular Shape
3//05 Drawing Good Lewis Structures. # valence e in atoms (± charge) must = # e in structure ; always. determine connectivity: least EN usually central; avoid small rings; always terminal ( e ); work out
More informationChapter 8. Molecular Shapes. Valence Shell Electron Pair Repulsion Theory (VSEPR) What Determines the Shape of a Molecule?
PowerPoint to accompany Molecular Shapes Chapter 8 Molecular Geometry and Bonding Theories Figure 8.2 The shape of a molecule plays an important role in its reactivity. By noting the number of bonding
More informationCarbon-based molecules are held together by covalent bonds between atoms
hapter 1: hemical bonding and structure in organic compounds arbon-based molecules are held together by covalent bonds between atoms omposition: Mainly nonmetals; especially,, O, N, S, P and the halogens
More informationSHAPES OF MOLECULES (VSEPR MODEL)
1 SAPES MLEULES (VSEPR MDEL) Valence Shell Electron-Pair Repulsion model - Electron pairs surrounding atom spread out as to minimize repulsion. - Electron pairs can be bonding pairs (including multiple
More informationCHEM 110 Exam 2 - Practice Test 1 - Solutions
CHEM 110 Exam 2 - Practice Test 1 - Solutions 1D 1 has a triple bond. 2 has a double bond. 3 and 4 have single bonds. The stronger the bond, the shorter the length. 2A A 1:1 ratio means there must be the
More informationChapter 11 Answers. Practice Examples
hapter Answers Practice Examples a. There are three half-filled p orbitals on, and one half-filled 5p orbital on I. Each halffilled p orbital from will overlap with one half-filled 5p orbital of an I.
More informationChapter 9: Molecular Geometries and Bonding Theories Learning Outcomes: Predict the three-dimensional shapes of molecules using the VSEPR model.
Chapter 9: Molecular Geometries and Bonding Theories Learning Outcomes: Predict the three-dimensional shapes of molecules using the VSEPR model. Determine whether a molecule is polar or nonpolar based
More informationShapes of Molecules and Hybridization
Shapes of Molecules and Hybridization A. Molecular Geometry Lewis structures provide us with the number and types of bonds around a central atom, as well as any NB electron pairs. They do not tell us the
More informationChapter 9. Molecular Geometries and Bonding Theories. Lecture Presentation. John D. Bookstaver St. Charles Community College Cottleville, MO
Lecture Presentation Chapter 9 Theories John D. Bookstaver St. Charles Community College Cottleville, MO Shapes The shape of a molecule plays an important role in its reactivity. By noting the number of
More informationChemistry: The Central Science. Chapter 9: Molecular Geometry and Bonding Theory
Chemistry: The Central Science Chapter 9: Molecular Geometry and Bonding Theory The shape and size of a molecule of a particular substance, together with the strength and polarity of its bonds, largely
More informationValence Bond Theory - Description
Bonding and Molecular Structure - PART 2 - Valence Bond Theory and Hybridization 1. Understand and be able to describe the Valence Bond Theory description of covalent bond formation. 2. Understand and
More informationHelpful Hints Lewis Structures Octet Rule For Lewis structures of covalent compounds least electronegative
Helpful Hints Lewis Structures Octet Rule Lewis structures are a basic representation of how atoms are arranged in compounds based on bond formation by the valence electrons. A Lewis dot symbol of an atom
More informationChapter Molecules are 3D. Shapes and Bonds. Chapter 9 1. Chemical Bonding and Molecular Structure
Chapter 9 Chemical Bonding and Molecular Structure 1 Shape 9.1 Molecules are 3D Angle Linear 180 Planar triangular (trigonal planar) 120 Tetrahedral 109.5 2 Shapes and Bonds Imagine a molecule where the
More informationLecture outline: Section 9. theory 2. Valence bond theory 3. Molecular orbital theory. S. Ensign, Chem. 1210
Lecture outline: Section 9 Molecular l geometry and bonding theories 1. Valence shell electron pair repulsion theory 2. Valence bond theory 3. Molecular orbital theory 1 Ionic bonding Covalent bonding
More informationGeneral and Inorganic Chemistry I.
General and Inorganic Chemistry I. Lecture 1 István Szalai Eötvös University István Szalai (Eötvös University) Lecture 1 1 / 29 Outline István Szalai (Eötvös University) Lecture 1 2 / 29 Lewis Formulas
More informationCOVALENT BONDING CHEMICAL BONDING I: LEWIS MODEL. Chapter 7
Chapter 7 P a g e 1 COVALENT BONDING Covalent Bonds Covalent bonds occur between two or more nonmetals. The two atoms share electrons between them, composing a molecule. Covalently bonded compounds are
More informationCovalent Bonding. Chapter 8. Diatomic elements. Covalent bonding. Molecular compounds. 1 and 7
hapter 8 ovalent bonding ovalent Bonding A metal and a nonmetal transfer An ionic bond Two metals just mix and don t react An alloy What do two nonmetals do? Neither one will give away an electron So they
More informationLewis Structure. Lewis Structures & VSEPR. Octet & Duet Rules. Steps for drawing Lewis Structures
Lewis Structure Lewis Structures & VSEPR Lewis Structures shows how the are arranged among the atoms of a molecule There are rules for Lewis Structures that are based on the formation of a Atoms want to
More informationCovalent Compounds: Bonding Theories and Molecular Structure
CHM 123 Chapter 8 Covalent Compounds: Bonding Theories and Molecular Structure 8.1 Molecular shapes and VSEPR theory VSEPR theory proposes that the geometric arrangement of terminal atoms, or groups of
More informationChapter 9. Molecular Geometries and Bonding Theories. Lecture Presentation. John D. Bookstaver St. Charles Community College Cottleville, MO
Lecture Presentation Chapter 9 Theories John D. Bookstaver St. Charles Community College Cottleville, MO Shapes The shape of a molecule plays an important role in its reactivity. By noting the number of
More informationCHAPTER 8 BONDING: GENERAL CONCEPTS Ionic solids are held together by strong electrostatic forces that are omnidirectional.
CAPTER 8 BDIG: GEERAL CCEPTS 1 CAPTER 8 BDIG: GEERAL CCEPTS Questions 15. a. This diagram represents a polar covalent bond as in. In a polar covalent bond, there is an electron rich region (indicated by
More informationCh 6 Chemical Bonding
Ch 6 Chemical Bonding What you should learn in this section (objectives): Define chemical bond Explain why most atoms form chemical bonds Describe ionic and covalent bonding Explain why most chemical bonding
More informationEx. 1) F F bond in F = 0 < % covalent, no transfer of electrons
#60 Notes Unit 8: Bonding Ch. Bonding I. Bond Character Bonds are usually combinations of ionic and covalent character. The electronegativity difference is used to determine a bond s character. Electronegativity
More informationChapter 10. VSEPR Model: Geometries
Chapter 10 Molecular Geometry VSEPR Model: Geometries Valence Shell Electron Pair Repulsion Theory Electron pairs repel and get as far apart as possible Example: Water Four electron pairs Farthest apart
More informationMolecular Orbitals. Chapter 9. Sigma bonding orbitals. Sigma bonding orbitals. Pi bonding orbitals. Sigma and pi bonds
Molecular Orbitals Chapter 9 Orbitals and Covalent Bond The overlap of atomic orbitals from separate atoms makes molecular orbitals Each molecular orbital has room for two electrons Two types of MO Sigma
More informationLewis Structure and Electron Dot Models
Lewis Structure and Electron Dot Models The Lewis Structure is a method of displaying the electrons present in any given atom or compound. Steps: 1. Make a skeleton structure 2. Count all e- available
More informationMolecular Geometry and Bonding Theories. Molecular Shapes. Molecular Shapes. Chapter 9 Part 2 November 16 th, 2004
Molecular Geometry and Bonding Theories Chapter 9 Part 2 November 16 th, 2004 8 Molecular Shapes When considering the geometry about the central atom, we consider all electrons (lone pairs and bonding
More informationChapter 13: Phenomena
Chapter 13: Phenomena Phenomena: Scientists measured the bond angles of some common molecules. In the pictures below each line represents a bond that contains 2 electrons. If multiple lines are drawn together
More informationChapter 9. and Bonding Theories
Chemistry, The Central Science, 11th edition Theodore L. Brown, H. Eugene LeMay, Jr., and Bruce E. Bursten Chapter 9 Theories John D. Bookstaver St. Charles Community College Cottleville, MO Shapes The
More informationChapter 9. Molecular Geometry and Bonding Theories
9.1 Molecular Shapes Read Sec. 9.1 and 9.2, then complete the Sample and Practice Exercises in these sections. Sample Exercise 9.1 (p. 347) Use the VSEPR model to predict the molecular geometries of a)
More informationUnit Six --- Ionic and Covalent Bonds
Unit Six --- Ionic and Covalent Bonds Electron Configuration in Ionic Bonding Ionic Bonds Bonding in Metals Valence Electrons Electrons in the highest occupied energy level of an element s atoms Examples
More informationMolecular Geometry and intermolecular forces. Unit 4 Chapter 9 and 11.2
1 Molecular Geometry and intermolecular forces Unit 4 Chapter 9 and 11.2 2 Unit 4.1 Chapter 9.1-9.3 3 Review of bonding Ionic compound (metal/nonmetal) creates a lattice Formula doesn t tell the exact
More informationChapter 9 Molecular Geometry. Lewis Theory-VSEPR Valence Bond Theory Molecular Orbital Theory
Chapter 9 Molecular Geometry Lewis Theory-VSEPR Valence Bond Theory Molecular Orbital Theory Sulfanilamide Lewis Structures and the Real 3D-Shape of Molecules Lewis Theory of Molecular Shape and Polarity
More informationPeriodic Trends. Homework: Lewis Theory. Elements of his theory:
Periodic Trends There are various trends on the periodic table that need to be understood to explain chemical bonding. These include: Atomic/Ionic Radius Ionization Energy Electronegativity Electron Affinity
More informationChapter 7 Chemical Bonding and Molecular Structure
Chapter 7 Chemical Bonding and Molecular Structure Three Types of Chemical Bonding (1) Ionic: formed by electron transfer (2) Covalent: formed by electron sharing (3) Metallic: attraction between metal
More informationHybridization of Orbitals
Hybridization of Orbitals Structure & Properties of Matter 1 Atomic Orbitals and Bonding Previously: Electron configurations Lewis structures Bonding Shapes of molecules Now: How do atoms form covalent
More informationChapter 9. and Bonding Theories. Molecular Shapes. What Determines the Shape of a Molecule? 3/8/2013
Chemistry, The Central Science, 10th edition Theodore L. Brown, H. Eugene LeMay, Jr., and Bruce E. Bursten Chapter 9 Theories John D. Bookstaver St. Charles Community College St. Peters, MO 2006, Prentice-Hall,
More informationMolecular Structure. Valence Bond Theory Overlap of atomic orbitals is a covalent bond that joins atoms together to form a molecule
Molecular Structure Topics 3-D structure shape (location of atoms in space) Molecular Geometry Valence Bond Theory Hybrid Orbitals Multiple Bonds VSEPR (Valence Shell Electron Pair Repulsion) Valence Bond
More informationChemical Bonding II: Molecular Geometry and Hybridization of Atomic Orbitals
Chemical Bonding II: Molecular Geometry and Hybridization of Atomic Orbitals 1 Chemical Bonding II Molecular Geometry (10.1) Dipole Moments (10.2) Valence Bond Theory (10.3) Hybridization of Atomic Orbitals
More informationProblems and questions How is a molecule or polyatomic ion held together? Why are atoms distributed at strange angles? Why are molecules not flat?
1 Cocaine 2 Problems and questions ow is a molecule or polyatomic ion held together? Why are atoms distributed at strange angles? Why are molecules not flat? Can we predict the structure? ow is structure
More information17/11/2010. Lewis structures
Reading assignment: 8.5-8.8 As you read ask yourself: How can I use Lewis structures to account for bonding in covalent molecules? What are the differences between single, double and triple bonds in terms
More informationIntroduction to VSEPR Theory 1
1 Class 8: Introduction to VSEPR Theory Sec 10.2 VSEPR Theory: The Five Basic Shapes Two Electron Groups: Linear Geometry Three Electron Groups: Trigonal Planar Geometry Four Electron Groups: Tetrahedral
More informationChapter 10. VSEPR Model: Geometries
Chapter 10 Molecular Geometry VSEPR Model: Geometries Valence Shell Electron Pair Repulsion Theory Electron pairs repel and get as far apart as possible Example: Water Four electron pairs Two bonds Two
More informationCHEMICAL BONDING. Chemical Bonds. Ionic Bonding. Lewis Symbols
CHEMICAL BONDING Chemical Bonds Lewis Symbols Octet Rule whenever possible, valence electrons in covalent compounds distribute so that each main-group element is surrounded by 8 electrons (except hydrogen
More informationChapter 9. Molecular Geometry and Bonding Theories
Chapter 9. Molecular Geometry and Bonding Theories PART I Molecular Shapes Lewis structures give atomic connectivity: they tell us which atoms are physically connected to which atoms. The shape of a molecule
More informationChapter 10. The Shapes of Molecules
Chapter 10 The Shapes of Molecules Molecules are visualized using Lewis Structures Molecular formula Step 1 Atom placement Step 2 Add A-group numbers ctet Rule Sum of valence e - Step 3 Remaining valence
More informationCHAPTER 8. Molecular Structure & Covalent Bonding Theories
CAPTER 8 Molecular Structure & Covalent Bonding Theories 1 Chapter Goals 1. A Preview of the Chapter 2. Valence Shell Electron Pair Repulsion (VSEPR) Theory 3. Polar Molecules:The Influence of Molecular
More informationTopic 2. Structure and Bonding Models of Covalent Compounds of p-block Elements
Topic 2 2-1 Structure and Bonding Models of Covalent Compounds of p-block Elements Bonding 2-2 Many different approaches to describe bonding: Ionic Bonding: Elements with large electronegativity differences;
More informationBonding. Honors Chemistry 412 Chapter 6
Bonding Honors Chemistry 412 Chapter 6 Chemical Bond Mutual attraction between the nuclei and valence electrons of different atoms that binds them together. Types of Bonds Ionic Bonds Force of attraction
More information5 Polyatomic molecules
s manual for Burrows et.al. Chemistry 3 Third edition 5 Polyatomic molecules Answers to worked examples WE 5.1 Formal charges in N 2 (on p. 221 in Chemistry 3 ) Use formal charges to decide whether oxygen
More informationChapter 9. Molecular Geometry and Bonding Theories
Chapter 9 Molecular Geometry and Bonding Theories MOLECULAR SHAPES 2 Molecular Shapes Lewis Structures show bonding and lone pairs do not denote shape Use Lewis Structures to determine shapes Molecular
More informationChapters 8 and 9. Octet Rule Breakers Shapes
Chapters 8 and 9 Octet Rule Breakers Shapes Bond Energies Bond Energy (review): The energy needed to break one mole of covalent bonds in the gas phase Breaking bonds consumes energy; forming bonds releases
More informationSubtopic 4.2 MOLECULAR SHAPE AND POLARITY
Subtopic 4.2 MOLECULAR SHAPE AND POLARITY 1 LEARNING OUTCOMES (covalent bonding) 1. Draw the Lewis structure of covalent molecules (octet rule such as NH 3, CCl 4, H 2 O, CO 2, N 2 O 4, and exception to
More informationAdapted from CHM 130 Maricopa County, AZ Molecular Geometry and Lewis Dot Formulas Introduction
Adapted from CHM 130 Maricopa County, AZ Molecular Geometry and Lewis Dot Formulas Introduction A chemical bond is an intramolecular (within the molecule) force holding two or more atoms together. Covalent
More informationBonding. Polar Vs. Nonpolar Covalent Bonds. Ionic or Covalent? Identifying Bond Types. Solutions + -
Chemical Bond Mutual attraction between the nuclei and valence electrons of different atoms that binds them together. Bonding onors Chemistry 412 Chapter 6 Types of Bonds Ionic Bonds Force of attraction
More informationAndrew Rosen *Note: If you can rotate a molecule to have one isomer equal to another, they are both the same
*Note: If you can rotate a molecule to have one isomer equal to another, they are both the same *Note: For hybridization, if an SP 2 is made, there is one unhybridized p orbital (because p usually has
More informationVSEPR. Valence Shell Electron Pair Repulsion Theory
VSEPR Valence Shell Electron Pair Repulsion Theory Vocabulary: domain = any electron pair or bond (single, double or triple) is considered one domain. bonding pair = shared pair = any electron pair that
More information8.3 Bonding Theories > Chapter 8 Covalent Bonding. 8.3 Bonding Theories. 8.1 Molecular Compounds 8.2 The Nature of Covalent Bonding
Chapter 8 Covalent Bonding 8.1 Molecular Compounds 8.2 The Nature of Covalent Bonding 8.3 Bonding Theories 8.4 Polar Bonds and Molecules 1 Copyright Pearson Education, Inc., or its affiliates. All Rights
More informationLecture Presentation. Chapter 10 Chemical Bonding II: Molecular Shapes, Valence Bond Theory, and Molecular Orbital Theory
Lecture Presentation Chapter 10 Chemical Bonding II: Molecular Shapes, Valence Bond Theory, and Molecular Orbital Theory Predicting Molecular Geometry 1. Draw the Lewis structure. 2. Determine the number
More information4 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.
CHEMISTRY & YOU Chapter 8 Covalent Bonding 8.1 Molecular Compounds 8.2 The Nature of Covalent Bonding 8.3 Bonding Theories 8.4 Polar Bonds and Molecules 1 Copyright Pearson Education, Inc., or its affiliates.
More informationHybridization and Molecular Orbital (MO) Theory
ybridization and Molecular Orbital (MO) Theory Chapter 10 istorical Models G.N.Lewis and I. Langmuir (~1920) laid out foundations Ionic species were formed by electron transfer Covalent molecules arise
More informationEXPERIMENT 12: MOLECULAR ARCHITECTURE
Name Section EXPERIMENT 12: MLECULAR ARCITECTURE PRE-LABRATRY QUESTINS The following preparatory questions should be answered before coming to lab. They are intended to introduce you to several ideas important
More informationChapter 9. Covalent Bonding: Orbitals
Chapter 9 Covalent Bonding: Orbitals Localized electron model A bond is made when a half-filled orbital of one atom overlaps with a half-filled orbital of another.! Bond: orbitals overlap straight on p
More informationStructure and Bonding of Organic Molecules
Chem 220 Notes Page 1 Structure and Bonding of Organic Molecules I. Types of Chemical Bonds A. Why do atoms forms bonds? Atoms want to have the same number of electrons as the nearest noble gas atom (noble
More informationCovalent Bonding Introduction, 2. Chapter 7 Covalent Bonding. Figure 7.1 The Hydrogen Molecule. Outline. Covalent Bonding Introduction, 1. Figure 7.
Covalent Bonding Introduction, 2 William L. Masterton Cecile N. Hurley http://academic.cengage.com/chemistry/masterton Chapter 7 Covalent Bonding Electron density Electrons are located between nuclei Electrostatic
More informationChapter 6 Molecular Structure
hapter 6 Molecular Structure 1. Draw the Lewis structure of each of the following ions, showing all nonzero formal charges. Indicate whether each ion is linear or bent. If the ion is bent, what is the
More informationbond energy- energy required to break a chemical bond -We can measure bond energy to determine strength of interaction
bond energy- energy required to break a chemical bond -We can measure bond energy to determine strength of interaction ionic compound- a metal reacts with a nonmetal Ionic bonds form when an atom that
More informationValence Bond Model and Hybridization
Valence Bond Model and ybridization APPENDIX 4 1 Concepts The key ideas required to understand this section are: Concept Book page reference VSEPR theory 65 More advanced ideas about electronic structure
More informationChapter 9 Molecular Geometry and Bonding Theories
Chapter 9 Molecular Geometry and Bonding Theories molecular shapes the VSEPR model molecular shape and molecular polarity covalent bonding and orbital overlap hybrid orbitals multiple bonds 9.1 Molecular
More informationLab Lecture on VSEPR and SPARTAN Chem 141 Lab Dr Abrash 10/3/2011
Q: What is the purpose of this lab? Lab Lecture on VSEPR and SPARTAN Chem 141 Lab Dr Abrash 10/3/2011 To learn two methods to study and predict the shapes of molecules. One is a rule based paper method
More informationRESONANCE STRUCTURE When a molecule has more than one possible structure. Draw all possible structures and place a double end arrow ( ) in between.
CHEMISTRY NOTES 6.1 COVALENT BONDS Objectives Explain the role and location of electrons in a covalent bond. Describe the change in energy and stability that takes place as a covalent bond forms. Distinguish
More informationMolecular shape is determined by the number of bonds that form around individual atoms.
Chapter 9 CH 180 Major Concepts: Molecular shape is determined by the number of bonds that form around individual atoms. Sublevels (s, p, d, & f) of separate atoms may overlap and result in hybrid orbitals
More informationChapter 9: Chemical Bonding I: Lewis Theory. Lewis Theory: An Overview
Chapter 9: Chemical Bonding I: Lewis Theory Dr. Chris Kozak Memorial University of ewfoundland, Canada Lewis Theory: An verview Valence e - play a fundamental role in chemical bonding. e - transfer leads
More informationCHAPTER TEN MOLECULAR GEOMETRY MOLECULAR GEOMETRY V S E P R CHEMICAL BONDING II: MOLECULAR GEOMETRY AND HYBRIDIZATION OF ATOMIC ORBITALS
CHAPTER TEN CHEMICAL BONDING II: AND HYBRIDIZATION O ATOMIC ORBITALS V S E P R VSEPR Theory In VSEPR theory, multiple bonds behave like a single electron pair Valence shell electron pair repulsion (VSEPR)
More informationCHEM 101: CHAPTER 11: CHEMICAL BONDS: THE FORMATION OF COMPOUNDS FROM ATOMS
1 CHEM 101: CHAPTER 11: CHEMICAL BONDS: THE FORMATION OF COMPOUNDS FROM ATOMS PERIODIC TRENDS: See pages 214-216, 221 Table 11.3, and 227 + 228 of text. Lewis Structures of Atoms: The Lewis Dot Diagram
More informationLewis Theory of Shapes and Polarities of Molecules
Lewis Theory of Shapes and Polarities of Molecules Sulfanilamide Lewis Structures and the Real 3D-Shape of Molecules Molecular Shape or Geometry The way in which atoms of a molecule are arranged in space
More information(A) 1 bonding pair (B) 1 bonding pair and 1 lone pair (C) 2 bonding pairs (D) 2 bonding pairs and 2 lone pairs
AP Chemistry - Problem Drill 13: Lewis Structures and VSPER No. 1 of 10 1. Lewis structure is used to model covalent bonds of a molecule or ion. Covalent bonds are a type of chemical bonding formed by
More informationMolecular Geometries. Molecular Geometries. Remember that covalent bonds are formed when electrons in atomic orbitals are shared between two nuclei.
Molecular Geometries Lewis dot structures are very useful in determining the types of bonds in a molecule, but they may not provide the best insight into the spatial geometry of a molecule, i.e., how the
More informationChemical Bonding II: Molecular Geometry and Hybridization of Atomic Orbitals
Chemical Bonding II: Molecular Geometry and Hybridization of Atomic Orbitals Chapter 10 Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. 1 Valence shell electron
More informationMolecular shape is only discussed when there are three or more atoms connected (diatomic shape is obvious).
Chapter 10 Molecular Geometry (Ch9 Jespersen, Ch10 Chang) The arrangement of the atoms of a molecule in space is the molecular geometry. This is what gives the molecules their shape. Molecular shape is
More informationA DOT STRUCTURE FOR A LARGER MOLECULE ETHANOL! Count valence electrons
212 A DOT STRUCTURE FOR A LARGER MOLECULE Count valence electrons Pick central atom and draw skeletal structure - central atom is usually the one that needs to gain the most electrons! - skeletal structure
More information