Chapter 7. Covalent Bonds and Molecular Structure

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 • Electrical 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 Non-polar covalent, polar covalent, ionic Electronegativity • 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 4 4 2.1 0 1 4 3 Non-polar covalent Polar covalent Ionic No (or very small difference in electronegativity Electronegativity difference larger than 0.4-2 Electronegativity difference larger than 2 • • • Covalent Bond 4 1.9 Red = high electron density Blue = low electron density Green = in between 8 Bond Lengths and Strengths • There is an optimum distance that will maximize attraction. This is the bond length which is calculated by adding the radii of two atoms. 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) 1200 triple 1000 800 double 600 single 400 200 0 100 120 140 160 180 200 220 Bond length (pm ) Bond Lengths and Strengths Lewis Dot Symbols Bond Strength < < Bond Length > > 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. • 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. Mainly used for covalent compounds. Not used for d-block compounds (which are ionic anyway). Electron-Dot Structures Electron-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). Procedure to give each atom an octet (in most cases):
Count total valence electrons • Add or subtract electrons for polyatomic ions  Draw an atomic skeleton  Place electron pairs (single bonds) between bonded atoms  Place remaining electrons on the outside atoms, then the central atom  Shift electrons, as necessary, to make multiple bonds and satisfy the octet rule and # of valence electrons 14 Exceptions to the Octet Rule A couple of hints for drawing Lewis Structures • Hydrogen will always be at the end of a molecule • Odd-Electron Molecules – This is because H can only hold 2 electrons in its 1s orbital. – It will always have a duet. – Draw a Lewis structure for NO and NO2 – Why does NO2 combine with itself to form N2 O 4 ? – Odd-electron molecules are very reactive. They are called radicals. You will see these much less often than even-number molecules. • 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 • Incomplete Octets – Certain central molecules don’t need an octet. – Draw structures for BeCl2, BH3, BF3, AlCl3 – Always try single bonds first 16 Exceptions 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: SF4 SF6 IF4+ XeF4 XeF2 PF5 ClF3 BrF5 Cl • Where do we find central atoms that expand their octets? 18 17 The Concept of Resonance The Concept of Resonance – Ozone • 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. 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. – Examples: O3, SO3, NO3-, CO32- O O O 121 pm? 148 pm? Experimental evidence indicates that both bond lengths are exactly the same, 128 pm. 19 The Concept of Resonance – Ozone O O O Resonance • How many different valid Lewis formulas can you draw for the following molecules or ions? O O O The structure of ozone can be best described by using both structures simultaneously. SO2 The structures do not flip “back-and-forth” NO3- SO3 CO32NCS- 22 Resonance Formal Charge What about structures with non-equivalent resonance structures? Take the structure for CO2. 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 Lewis Structure Practice Molecular Shapes: The VSEPR Model • Valence-Shell Electron-Pair Repulsion •CO2 •SO2 •BF3 •CH4 •NH4+ •CCl4 •NH3 •H2O •SF2 •COCl2 •PCl5 •SF4 •ClF3 •XeF2 •OF2 •SF6 •BrF5 •AlCl4•AsH3 • Electron 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 • Electron domains occur as far apart as possible • Notice the single, double, and triple bonds each count as ONE electron domain. 25 VSEPR Theory 26 Five Fundamental VSEPR Geometries • Use A, B, E notation: A = central atom; B = # outer atoms; E = # lone e- pairs – CH4 = AB4, NH3 = AB3E, H2O = AB2E2 • Can predict the angles between electron domains (charge clouds, areas of electron density): • 2 domains - linear (180o) • 3 domains - trigonal planar (120o) • 4 domains - tetrahedral (109.5o) • 5 domains - trigonal bipyramidal (90o & 120o) • 6 domains - octahedral (90o) 27 28 Molecular and Electronic Geometries Using these geometries, we can determine two different types of shapes. Electronic Geometry Shape that is made electron density make from a central atom. H O H Molecular Geometry Shape that is made when we are concerned only with the shape of the bonding electrons. H O H Four areas of electron density = Angle of bonds Tetrahedral Bent • A = central atom; B = # outer atoms; E = # lone e- pairs Molecular Geometries Electronic Geometries 1 lone pair 2 lone pair AB2 Linear AB3 Trigonal Planar AB2E Bent AB4 Tetrahedral AB3E Trigonal Pyramidal AB2E2 Bent AB5 Trig. Bipyramidal AB4E See-Saw AB3E2 T-Shaped AB6 Octahedral AB5E Square Pyramid AB4E2 Square Planar You need to know! What are the molecular geometries of the following molecules? •CO2 •SO2 •BF3 •CH4 •NH4+ •CCl4 •NH3 •H2O •SF2 •COCl2 Molecular Shapes • Should CO2 and SO2 have the same geometry? •PCl5 •SF4 •ClF3 •XeF2 •OF2 •SF6 •BrF5 •AlCl4•AsH3 31 Valence Bond Theory 32 Creating Covalent Bonds • Lewis structures and VSEPR give information about the shapes of molecules and the distributions of electrons. They don’t explain why a bond forms. How exactly do orbitals interact to create covalent bonds? • Valence-bond theory considers both bond formation and molecular shape Half-filled orbitals overlap so that 2 electrons can share space and form a covalent bond. • Looks at how electrons are shared in a covalent bond We can combine two s orbitals (H2) • VB theory considers the atomic orbitals occupied by the valence electrons 33 Creating Covalent Bonds We can combine two p orbitals (F2) Hybridization zebra + donkey lion + tiger We can combine an s and a p orbital (HCl) Zonkey Liger Hybrid orbitals Hybridization in Methane 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, CH4, for example….carbon has 4 valence electrons to pair up with 1 valence electron of 4 hydrogens. • Need 4 equivalent orbitals to form the 4 single (σ) bonds (based on VSEPR and experiments) • Ground state configuration: 2p 2s 2p 2s 2p E The four bonds formed with hydrogen in methane would appear to form from two different types of orbitals. 2p Promotion of E an electron 2s 2s 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 Hybridization in Boron Trifluoride This new model can account for the experimental evidence for methane that suggests that all of the bonds in CH4 are equivalent. • • • • Hybridization Boron trifluoride: BF3 Lewis Dot Structure? Boron is central atom Electron configuration of boron? – [He] 2s2 2p1 • VSEPR shape? E sp3 – Trigonal planar Each orbital allows a single bond to form Hybridization in Boron Trifluoride Hybridization in Beryllium Chloride • Need 3 equivalent orbitals to form the 4 single (s) bonds (based on VSEPR and experiments) • Ground state configuration: • • • • Beryllium chloride: BeCl2 Lewis Dot Structure? Beryllium is central atom Electron configuration of beryllium? – [He] 2s2 • VSEPR shape? 2p Promotion & hybridization E 2s p E sp2 Notice one unhybridized p orbital is left over – Linear Hybridization in Beryllium Chloride Hybridization in Phosphorus Pentachloride • Need 2 equivalent orbitals to form the 4 single (s) bonds (based on VSEPR and experiments) • Ground state configuration: • • • • Phosphorus pentachloride: PCl5 Lewis Dot Structure? Phosphorus is central atom Electron configuration of phosphorus? – [Ne] 3s2 3p3 • VSEPR shape? 2p Promotion & hybridization E p E – Trigonal bipyramidal sp 2s Hybridization in Phosphorus Pentachloride • How many single bonds in PCl5? • Where does the 5th orbital come from? – Expanded octet – sp3d Hybridization in Sulfur Hexafluoride • • • • Sulfur Hexafluoride: SF6 Lewis Dot Structure? Sulfur is central atom Electron configuration of sulfur? – [Ne] 3s2 3p4 • VSEPR shape? 3d E 3d Promotion & hybridization 2p – Octahedral E sp3d 2s Hybridization in Sulfur Hexafluoride • Need 6 equivalent orbitals to form the 6 single bonds. • Where does the 6th orbital come from? – Expanded octet – sp3d2 3d E 2p 2s 3d Promotion & hybridization E sp3d2 Hybridization and Geometry Once the electronic geometry of a molecule is known, the hybridization can be predicted Linear sp Trigonal Planar sp2 Tetrahedral sp3 Trigonal Bipyramidal sp3d Octahedral sp3d2 Hybridization in Multiple Bonds: Ethylene • • • • Hybridization in Ethylene Based on the Lewis structure, there should be a double bond. How does a double bond form? Ethylene: CH2CH2 Lewis Dot Structure? Carbon is central atom Electron configuration of carbon? Each Carbon 2p – [He] 2s2 2p2 E • VSEPR shape? Each Carbon 2p Promotion & hybridization E sp2 2s – Trigonal planar Hybridization in Ethylene Hybridization in Ethylene Unhybridized p orbital from C sigma bond (σ σ): direct overlap of bonding orbitals. pi bond (π π): “sideways” overlap of bonding orbitals (usually between unhybridized p orbitals) sp2 2p sp2 2p s orbital from H 3 Hybridized sp2 orbital from C bond formed by p-p overlap Hybridization in Acetylene • • • • Acetylene: HCCH Lewis Dot Structure? Carbon is central atom Electron configuration of carbon? Hybridization in Acetylene sp sp 2p – [He] 2s2 2p2 • VSEPR shape? bond formed by p-p overlap (π π) – Linear bond formed by p-p overlap (π π) 2p Hybridization in Acetylene Summary • sp hybridization accounts for single (σ) bonds, but what about the unhybridized 2p orbitals? • Those 2 unhybridized 2p orbitals help make the two π bonds. 2 Hybridized sp orbitals from C s orbital from H 2 Unhybridized p orbitals from C Summary • The number of hybridized orbitals equals the number of electron domains around a central atom (starting with s) • • • • – sp, sp2, sp3, sp3d, sp3d2 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) • 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 (BeCl2 and HCN) – sp2  3 orbitals combined (BF3 and CH2O) – sp3  4 orbitals combined (CH4) – sp3d  5 orbitals combined (PCl5) – sp3d2  6 orbitals combined (SF6)