Chapter 7. Comparing Ionic and Covalent Bonds. Ionic Bonds. Types of Bonds. Quick Review of Bond Types. Covalent Bonds

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1 Comparing Ionic and Covalent Bonds Chapter 7 Covalent Bonds and Molecular Structure Intermolecular forces (much weaker than bonds) must be broken Ionic bonds must be broken 1 Ionic Bonds Covalent Bonds Ionic bonds are very strong, so separating ions requires a lot of energy high melting points high boiling points Ionic bonds Solids are usually relatively soft. low melting points low boiling points Intermolecular forces, NOT bonds, break to melt, boil etc. between ions must break to melt, boil, etc. Crystals are hard and brittle Crystal lattice is an arrangement of ions of opposite charge surrounding one another in three dimensions lectrical insulators when solid, electrical conductors when molten or dissolved in water Properties arise because molecules are not connected to other molecules by bonds, but by intermolecular forces. Usually composed of nonmetals 3 Quick Review of Bond Types Types of Bonds Classify the following substances by the 4 type of bond: CaF2 CuCl2 NCl3 H2O NH4Cl K2SO4 5 CaF2 CuCl2 NCl3 H2O NH4Cl K2SO4 ionic ionic covalent covalent ionic and covalent ionic and covalent 6

2 lectronegativity Non-polar covalent, polar covalent, ionic The ability of an atom in a molecule to attract shared electrons to itself. Relative values no units (4 for F is highest). Note the relative trend. H belongs between B and C Non-polar covalent Polar covalent Ionic No (or very small difference in electronegativity lectronegativity difference larger than lectronegativity difference larger than 2 Red = high electron density Blue = low electron density Green = in between 8 Covalent Bond There is an optimum distance that will maximize attraction. This is the bond length which is calculated by adding the radii of two atoms. Bond Lengths and Strengths What is the general trend of bond strength vs. bond length? What is the trend for number of bonds? Notice that the trend is not absolute Bond Strength (KJ/mol) triple double single Bond length (pm) Bond Lengths and Strengths Lewis Dot Symbols Bond Strength Bond Length < < > > Covalent bonding focuses on interactions of valence electrons of two or more atoms Can use Lewis Dot Symbols to represent the numbers of valence electrons for each atom (based on electron configurations) Usually used for predicting structures of covalent molecules. In general, triple bonds are stronger (and shorter) than double bonds which are stronger than single bonds. Also, in general, shorter bonds are stronger than longer bonds. As with anything else in chemistry, there are exceptions. Mainly used for covalent compounds. Not used for d-block compounds (which are ionic anyway).

3 lectron-dot Structures Used to show how electrons are shared between nonmetals in a covalent bond. Use valence electrons to give each atom an octet, with a few exceptions (e.g., H). lectron-dot Structures Procedure to give each atom an octet (in most cases): ❶Count total valence electrons Add or subtract electrons for polyatomic ions ❷Draw an atomic skeleton ❸Place electron pairs (single bonds) between bonded atoms ❹Place remaining electrons on the outside atoms, then the central atom ❺Shift electrons, as necessary, to make multiple bonds and satisfy the octet rule and # of valence electrons 14 A couple of hints for drawing Lewis Structures Hydrogen will always be at the end of a molecule This is because H can only hold 2 electrons in its 1s orbital. It will always have a duet. Carbon will be a central atom Fluorine will always be a peripheral atom Most times, carbon will not have lone pairs as a stable molecule. There are a very few exceptions. CO and CN - are two exceptions. Central atoms are usually less electronegative than peripheral atoms Only draw multiple bonds if the structure cannot be correctly drawn with single bonds Always try single bonds first xceptions to the Octet Rule Odd-lectron Molecules Draw a Lewis structure for NO and NO 2 Why does NO 2 combine with itself to form N 2 O 4? Odd-electron molecules are very reactive. They are called radicals. You will see these much less often than even-number molecules. Incomplete Octets Certain central molecules don t need an octet. Draw structures for BeCl 2, BH 3, BF 3, AlCl 3 16 xceptions to the Octet Rule What do you do if all of your atoms have an octet but there are still electrons left over? Draw Lewis Structures for the following: SF 4 SF 6 IF + 4 XeF 4 XeF 2 PF 5 ClF 3 BrF 5 Where do we find central atoms that expand their octets? Cl 18 17

4 The Concept of Resonance While Lewis structures do help to predict the structures of many molecules, there are some structures that cannot be satisfactorily represented with a single Lewis structure. The Concept of Resonance Ozone Based on the Lewis structure we would expect O O bond (148 pm) to be longer than O = O bond (121 pm) Lewis formulas don t always accurately represent bonds. Sometimes it takes multiple formulas to adequately represent the electron distribution. xamples: O 3, SO 3, NO 3-, CO 3 2- O O O 121 pm? 148 pm? xperimental evidence indicates that both bond lengths are exactly the same, 128 pm. 19 The Concept of Resonance Ozone O O O O O O The structure of ozone can be best described by using both structures simultaneously. The structures do not flip back-and-forth Resonance How many different valid Lewis formulas can you draw for the following molecules or ions? SO 2 SO 3 CO 3 2- NO 3 - NCS - 22 Resonance Formal Charge What about structures with non-equivalent resonance structures? Take the structure for CO 2. Are there any resonance structures? If so, how do we choose the correct structure? Formal Charge = number of valence electrons in an atom (number of nonbonding electrons and ½ of bonding electrons assigned to atom) 23

5 Lewis Structure Practice CO 2 SO 2 BF 3 CH 4 NH 4 + CCl 4 NH 3 H 2 O SF 2 COCl 2 PCl 5 SF 4 ClF 3 XeF 2 OF 2 SF 6 BrF 5 AlCl 4 - AsH 3 Molecular Shapes: The VSPR Model Valence-Shell lectron-pair Repulsion lectron pairs (or groups of pairs) try to avoid one another because of repulsions between like-charged particles Regions where electrons are likely to be found are called electron domains: Lone electron pairs Single, double, and triple bonds lectron domains occur as far apart as possible Notice the single, double, and triple bonds each count as ON electron domain VSPR Theory Five Fundamental VSPR Geometries Use A, B, notation: A = central atom; B = # outer atoms; = # lone e - pairs CH 4 = AB 4, NH 3 = AB 3, H 2 O = AB 2 2 Can predict the angles between electron domains (charge clouds, areas of electron density): 2 domains - linear (180 o ) 3 domains - trigonal planar (120 o ) 4 domains - tetrahedral (109.5 o ) 5 domains - trigonal bipyramidal (90 o & 120 o ) 6 domains - octahedral (90 o ) Using these geometries, we can determine two different types of shapes. lectronic Geometry Molecular Geometry Molecular and lectronic Geometries A = central atom; B = # outer atoms; = # lone e - pairs Molecular Geometries lectronic Geometries 1 lone pair 2 lone pair Shape that is made electron density make from a central atom. Shape that is made when we are concerned only with the shape of the bonding electrons. AB 2 Linear AB 3 Trigonal Planar AB 2 Bent O H H Four areas of electron H O H AB 4 Tetrahedral AB 5 Trig. Bipyramidal AB 3 Trigonal Pyramidal AB 4 See-Saw AB 2 2 Bent AB 3 2 T-Shaped density = Tetrahedral Angle of bonds Bent AB 6 Octahedral AB 5 Square Pyramid You need to know! AB 4 2 Square Planar

6 What are the molecular geometries of the following molecules? CO 2 SO 2 BF 3 CH 4 NH 4 + CCl 4 NH 3 H 2 O SF 2 COCl 2 PCl 5 SF 4 ClF 3 XeF 2 OF 2 SF 6 BrF 5 AlCl 4 - AsH 3 Molecular Shapes Should CO 2 and SO 2 have the same geometry? Valence Bond Theory Lewis structures and VSPR give information about the shapes of molecules and the distributions of electrons. They don t explain why a bond forms. Valence-bond theory considers both bond formation and molecular shape Looks at how electrons are shared in a covalent bond Creating Covalent Bonds How exactly do orbitals interact to create covalent bonds? Half-filled orbitals overlap so that 2 electrons can share space and form a covalent bond. We can combine two s orbitals (H 2 ) VB theory considers the atomic orbitals occupied by the valence electrons 33 Creating Covalent Bonds We can combine two p orbitals (F 2 ) Hybridization zebra + donkey lion + tiger We can combine an s and a p orbital (HCl) Zonkey Liger

7 Hybrid orbitals Although we know that orbitals (s, p, d, f) overlap to form covalent bonds, there are situations where this simplistic model seems to fall apart. Take methane, CH 4, for example.carbon has 4 valence electrons to pair up with 1 valence electron of 4 hydrogens. Hybridization in Methane Need 4 equivalent orbitals to form the 4 single (σ) bonds (based on VSPR and experiments) Ground state configuration: The four bonds formed with hydrogen in methane would appear to form from two differen types of orbitals. Promotion of an electron But we know from experimental evidence that all 4 bonds are exactly the same! Same length, same strength, and same bonding angles! Hybridization in Methane This new model can account for the experimental evidence for methane that suggests that all of the bonds in CH 4 are equivalent. Hybridization sp 3 Hybridization in Boron Trifluoride Boron trifluoride: BF 3 Lewis Dot Structure? Boron is central atom lectron configuration of boron? [He] 2 1 VSPR shape? Trigonal planar ach orbital allows a single bond to form Hybridization in Boron Trifluoride Need 3 equivalent orbitals to form the 4 single (s) bonds (based on VSPR and experiments) Ground state configuration: Promotion & hybridization sp 2 p Hybridization in Beryllium Chloride Beryllium chloride: BeCl 2 Lewis Dot Structure? Beryllium is central atom lectron configuration of beryllium? [He] 2 VSPR shape? Linear Notice one unhybridized p orbital is left over

8 Hybridization in Beryllium Chloride Need 2 equivalent orbitals to form the 4 single (s) bonds (based on VSPR and experiments) Ground state configuration: Promotion & hybridization sp p Hybridization in Phosphorus Pentachloride Phosphorus pentachloride: PCl 5 Lewis Dot Structure? Phosphorus is central atom lectron configuration of phosphorus? [Ne] 3s 2 3p 3 VSPR shape? Trigonal bipyramidal Hybridization in Phosphorus Pentachloride How many single bonds in PCl 5? Where does the 5 th orbital come from? xpanded octet sp 3 d 3d Promotion & hybridization 3d sp 3 d Hybridization in Sulfur Hexafluoride Sulfur Hexafluoride: SF 6 Lewis Dot Structure? Sulfur is central atom lectron configuration of sulfur? [Ne] 3s 2 3p 4 VSPR shape? Octahedral Hybridization in Sulfur Hexafluoride Need 6 equivalent orbitals to form the 6 single bonds. Where does the 6 th orbital come from? xpanded octet sp 3 d 2 3d Promotion & hybridization 3d sp 3 d 2 Hybridization and Geometry Once the electronic geometry of a molecule is known, the hybridization can be predicted Linear Trigonal Planar Tetrahedral Trigonal Bipyramidal Octahedral sp sp 2 sp 3 sp 3 d sp 3 d 2

9 Hybridization in Multiple Bonds: thylene Hybridization in thylene thylene: CH 2 CH 2 Lewis Dot Structure? Based on the Lewis structure, there should be a double bond. How does a double bond form? Carbon is central atom lectron configuration of carbon? [He] 2 2 VSPR shape? Trigonal planar ach Carbon Promotion & hybridization ach Carbon sp 2 Hybridization in thylene sigma bond (σ): direct overlap of bonding orbitals. pi bond (π): sideways overlap of bonding orbitals (usually between unhybridized p orbitals) Hybridization in thylene Unhybridized p orbital from C sp 2 sp 2 s orbital from H 3 Hybridized sp 2 orbital from C bond formed by p-p overlap Hybridization in Acetylene Hybridization in Acetylene Acetylene: HCCH Lewis Dot Structure? Carbon is central atom lectron configuration of carbon? [He] 2 2 VSPR shape? Linear sp sp bond formed by p-p overlap (π) bond formed by p-p overlap (π)

10 Hybridization in Acetylene sp hybridization accounts for single (σ) bonds, but what about the unhybridized orbitals? Those 2 unhybridized orbitals help make the two π bonds. s orbital from H 2 Hybridized sp orbitals from C 2 Unhybridized p orbitals from C Summary According to Valence Bond Theory, covalent bonds form when: there are two electrons in each orbital; one electron from each atom these two orbitals overlap The number of hybridized orbitals equals the number of atomic orbitals that are combined. sp 2 orbitals combined (BeCl 2 and HCN) sp 2 3 orbitals combined (BF 3 and CH 2 O) sp 3 4 orbitals combined (CH 4 ) sp 3 d 5 orbitals combined (PCl 5 ) sp 3 d 2 6 orbitals combined (SF 6 ) Summary The number of hybridized orbitals equals the number of electron domains around a central atom (starting with s) sp, sp 2, sp 3, sp 3 d, sp 3 d 2 A single bond has 1 σ bond (the same goes for a lone pair of electrons) A double bond has 1 σ bond and 1 π bond A triple bond has 1 σ bond and 2 π bonds Unhybridized p orbitals participate in π bonding (to make double and triple bonds)

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