Molecular Geometry Part 3. David A. Katz Pima Community College Tucson, AZ
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1 Concepts of Chemical Bonding and Molecular Geometry Part 3 David A. Katz Pima Community College Tucson, AZ
2 Theories of Chemical Bonding
3 Valence Bond Theory and dobit Orbital loverlap We think of covalent bonds forming through the sharing of electrons by adjacent atoms. In this approach, we assume bonding occurs when orbitals on two atoms overlap. s orbital p orbital d orbital Overlaps can occur between s, p, and d orbitals p, p, on two different atoms
4 Orbital Overlap and Bonding Increased overlap brings the electrons and nuclei closer together while simultaneously decreasing electronelectron repulsion. However, if atoms get too close, the internuclear repulsion greatly raises the energy. This diagram is known as a Heitler-London diagram
5 Orbital Overlap Some examples of orbital overlaps are shown in the diagram below. The areas of bond overlap are shaded. This represents increased electron density of the chemical bond s-s s overlap s-p overlap p-p p overlap
6 Valence Bond Theory There are two ways orbitals can overlap to form bonds between atoms. Sigma bonds, symbolized by Pi bonds, symbolized by Hybridization, which we will look at later, is a major player in this approach to bonding.
7 Sigma ( ) Bonds Sigma bonds are characterized by Head-to-head overlap. Cylindrical symmetry of electron density about the internuclear axis.
8 Pi ( ) Bonds Pi bonds are characterized by Side-to-side overlap. Electron density above and below the internuclear axis. Although there are two areas of overlap, this is a single 2-electron bond
9 Single Bonds Single bonds are always bonds, because overlap is greater, resulting in a stronger bond and more energy lowering.
10 Multiple Bonds In a multiple bond one of the bonds is a bond and the rest are bonds.
11 Orbital Overlap and Molecular l Shapes There is a problem with the simple orbital overlap model. The shapes of the molecules are linear due to s-s, s-p, or p-p (end to end overlaps) The shapes are predicted to be at 90 angles if there are two or three different p orbitals involved on one atom In order to account for tetrahedral, trigonal bipyramidal, and other geometries arising from the atomic orbitals we recognize, we use a model known as HYBRID ORBITALS.
12 Hybrid Orbitals Consider beryllium: In its ground electronic state, it would not be able to form bonds because it has no singly-occupied orbitals.
13 Hybrid Orbitals But if it absorbs the small amount of energy needed to promote an electron from the 2s to the 2p orbital, it can form two bonds.
14 Hybrid Orbitals Mixing the s and p orbitals yields two degenerate Mixing the s and p orbitals yields two degenerate orbitals that are hybrids of the two orbitals. These are called sp hybrid orbitals and have two lobes like a p orbital. One of the lobes is larger and more rounded as is the s orbital.
15 Hybrid Orbitals These two degenerate orbitals would align themselves 180 from each other. This is consistent with the observed geometry of beryllium compounds: linear.
16 Hybrid Orbitals With hybrid orbitals the orbital diagram for beryllium would look like this. The sp orbitals are higher in energy than the 1ss orbital but lower than the 2p. p
17 Hybrid Orbitals Using a similar model for boron leads to The mixing of one s and two p orbitals gives rise to three sp 2 hybrid orbitals
18 Hybrid Orbitals three degenerate sp 2 orbitals. The three sp 2 orbitals arrange The three sp orbitals arrange to form trigonal planar bonds
19 Hybrid Orbitals With carbon we get The mixing of one s and three p orbitals gives rise to four sp 3 hybrid orbitals
20 Hybrid Orbitals four degenerate sp 3 orbitals. The four sp 3 orbitals arrange to form tetrahedral bonds
21 Hybrid Orbitals For geometries involving expanded octets on the central atom, we must use d orbitals in our hybrids. The mixing of one s and three p and one d orbitals gives rise to five sp 3 d hbid hybrid orbitals bitl
22 Hybrid Orbitals This leads to five degenerate sp 3 d orbitals. This forms a trigonal bipyramidal molecule In a similar manner, the mixing of one s, three p and two d orbitals gives rise to six degenerate sp 3 d 2 hybrid orbitals This forms an octahedral molecule
23 Hybrid Orbitals Once you know the electron-domain geometry, you know the hybridization state of the atom.
24 Multiple Bonds Utilize both hybridized and unhybridized orbitals In the formaldehyde d molecule (shown at left) an sp 2 orbital on carbon overlaps in a bond with the corresponding orbital on the oxygen. Hydrogen atoms overlap with the remaining sp 2 orbitals on the carbon. The unhybridized p orbitals overlap in bond.
25 Multiple Bonds In triple bonds, as in acetylene, two sp orbitals form a bond between the carbons, and two pairs of p orbitals overlap in fashion to form the two bonds.
26 Delocalized Electrons: Resonance Previously, we observed, when writing Lewis structures for species like the nitrate ion, we drew resonance structures to more accurately reflect the structure of the molecule or ion.
27 Delocalized Electrons: Resonance In reality, each of the four atoms in the nitrate ion has a p orbital. The p orbitals on all three oxygens overlap with the p orbital on the central nitrogen.
28 Delocalized Electrons: Resonance This means the electrons are not localized between the nitrogen and one of the oxygens, but rather are delocalized throughout the ion.
29 The organic molecule benzene has six bonds and a p orbital on each carbon atom. Resonance
30 Resonance In reality the electrons in benzene are not localized, but delocalized. The even distribution of the electrons in benzene makes the molecule unusually stable.
31 Molecular Orbital (MO) Theory Though valence bond theory effectively conveys most observed properties of ions and molecules, there are some concepts better represented by molecular orbitals. Molecular orbitals are considered to be formed from a linear combination of atomic orbitals (LCAO) If two atomic orbitals combine, then two molecular orbitals will be formed
32 Molecular Orbital (MO) Theory In MO theory, we invoke the wave nature of electrons. If waves interact constructively, the resulting orbital is lower in energy: a bonding molecular orbital.
33 Molecular Orbital (MO) Theory If waves interact destructively, the resulting orbital is higher in energy: an antibonding molecular orbital.
34 MO Theory These are the atomic orbitals on two different H atoms These are the two resulting molecular orbitals In H 2 the two electrons go into the bonding molecular orbital. The bond order is one half the difference between the number of bonding and antibonding electrons.
35 MO Theory For hydrogen, with two electrons in the bonding MO and none in the antibonding MO, the bond order is 1 - = 2 (2 0) 1
36 MO Theory In the case of He 2, the bond order would be 1 (2-2) = 0 2 Therefore, He 2 does not exist.
37 MO Theory For atoms with both s and p orbitals, there are two types of interactions: The s and dthe p orbitals that face each other overlap in fashion. The other two sets of p orbitals overlap in fashion.
38 MO Theory The resulting MO diagram looks like this. There are both and bonding molecular orbitals and * and * antibonding molecular orbitals.
39 MO Theory The smaller p-block elements in the second period have a sizeable interaction between the s and p orbitals. This flips the order of the s and p molecular l orbitals in these elements.
40 Second-Row MO Diagrams
41 Metallic Bonds
42 Physical Properties of Metals Conduct heat and electricity. Malleable (can be pressed or hammered into sheets). Ductile (can be drawn into wire). Atoms can slip past each other. So metals aren t as brittle as other solids.
43 Electron-Sea Model Metals can be thought of as cations suspended in sea of valence electrons. Attractions hold electrons near cations, but not so tightly as to impede their flow.
44 Electron-Sea Model This explains properties of metals Conductivity of heat and electricity Deformation
45 Molecular Orbital Model Electron-sea model does not explain observed trends in melting point, boiling point, heat of fusion, etc. Suggests these properties should increase with increasing number of valence electrons.
46 Molecular Orbital Model These trends can be explained by energy bands created by large number of molecular orbitals formed as metal atoms bond with each other.
47 Molecular Orbital Model As with nonmetals, bond order apexes in center of row, then decreases. Thus, attractions (and melting point, etc.) apex in center of transition metals. (Group 6B)
48 Alloys Mixtures of elements that have properties characteristic of metals. Many ordinary uses of metals involve alloys.
49 Solution Alloys Components of alloys are dispersed uniformly Substitutional alloys: Interstitial titi alloys: Particles close in size. Solute particles smaller than solvent. Solute particles take place of Solute particles find their way into solvent metal atoms. holes between solvent metal atoms.
50 Intermetallic Compounds Homogeneous alloys with definite properties and compositions. Co 5 Sm Used for permanent magnets in headsets and speakers.
51 Transition Metals Many important metals are included in this group. Comprised of elements in d block of periodic table.
52 Physical Properties of Transition Metals Some of their properties (such as ionization i energy, atomic radius, etc.) are suggestive of isolated atoms. Others (such as density, melting point, etc.) suggest bulk solid metal.
53 Atomic Radii Trends are similar across all three rows of transition metals. While Z eff increases across row, so does number of nonbonding electrons. These repel each other and increase radius.
54 Electron Configurations and Oxidation Statest Transition metals often have more than one common oxidation state. Most have +2 state t due to loss of s electrons. Oxidation numbers greater than 2 are due to loss of d electrons as well as s.
55 Electron Configurations and Oxidation Statest Many form compounds that have colors.
56 Electron Configurations and Oxidation Statest Many have significant magnetic properties. In diamagnetic elements, all electron spins are paired. Therefore, there is no net magnetic moment.
57 Electron Configurations and Oxidation Statest In paramagnetic atoms and ions, there are unpaired spins. The magnetic fields are randomly arranged, though, unless placed in an external magnetic field.
58 Electron Configurations and Oxidation Statest In ferromagnetic substances the orientations of magnetic fields from unpaired electrons are affected by spins from electrons around them.
59 Electron Configurations and Oxidation Statest When an external field is applied and then removed, the substance maintains the magnetic moment and becomes a permanent magnet.
60 Chromium Oxidized by HCl or H 2 SO 4 to form blue Cr 2+ ion. Cr 2+ oxidized by O 2 in air to form green Cr 3+. Cr also found in +6 state as in CrO 2 4 and dthe strong oxidizer Cr 2 O 2 7.
61 Iron Exists in solution in +2 or +3 state. Elemental iron reacts with nonoxidizing acids to form Fe 2+, which oxidizes in air to Fe 3+.
62 Iron Brown water running from a faucet is caused by insoluble Fe 2 O3. Fe 3+ soluble in acidic solution, but forms a hydrated oxide as red-brown gel in basic solution. o
63 Copper In solution exists in +1 or +2 state. +1 salts generally white, insoluble. +2 salts commonly blue, water-soluble.
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