Lecture 1 – Recapping Important Concepts - … 1 – Recapping Important Concepts ... 2s2 2p3...
Transcript of Lecture 1 – Recapping Important Concepts - … 1 – Recapping Important Concepts ... 2s2 2p3...
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2P32 Winter Term 2015-16 Principles of Inorganic ChemistryDr. M. Pilkington
Lecture 1 – Recapping Important Concepts
Inorganic Chemistry and the Periodic Table
Bonding Models
Shapes of Molecules - Lewis Structures
Valence bond theory: cases of NH3 H2O and BF3
Lewis Acids and Bases
σ and π bonds in CH2=CH2
The Shapes of Molecules – Relationship between Lewis Structure, VSEPR theory and VBT.
Assignment 1 – Drawing Lewis structures and predicting the shapes/geometries of molecules due after class Tuesday 12th January
1. Inorganic Chemistry and the Periodic Table
Carbon is only one element and has limited bonding modes, oxidation states and coordination numbers.But it does CATENATE well and forms MULTIPLE BONDS with itself and other p-block elements especially N and O.
For the rest of the elements:Wide range of electronegativity, oxidation states, coordination numbers, ability to form multiple bonds and catenate etc…
How can we make sense of such wide ranging behaviors?
We have a system called the Periodic Table. The ‘Periodic Law’ 1860-1870 (Mendeleev and Meyer): A periodic repetition of physical and chemical properties occurs when the elements are arranged in order of increasing atomic weight [number]’
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With the development of atomic theory and spectroscopic techniques the modern Periodic Table has evolved:
2P32 Course Outline:
Lectures 1-16Coordination Chemistry of transition Metal ions
Lectures 17 – 34Descriptive Inorganic Chemistry – Main Group Elements.
2. Bonding Models:
In covalent species, electrons are shared between atoms.In an ionic species, one or more electrons are transferred between atoms to form bonds.
Modern views of molecular structure are, based on applying wave mechanics to molecules; such studies provide answers as to how and why atoms combine. Two such methods are:1. Valence Bond (VB) approach- overlap of valence orbitals on atoms to form bonds. 2. Molecular Orbital theory (MO) of bond formation – allocates electrons to molecular orbitals formed by the overlap (interaction) of atomic orbitals.
Familiarity with both VB and MO concepts is necessary as it is often the case that a given situation can be approached using one or the other of these models.
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Lewis structures – you need to be able to draw these.
Lewis presented a simple but useful method of describing the arrangement of valence electrons in molecules.
Lewis structures give the connectivity of an atom in a molecule, the bond order and the number of lone pairs and these maybe used to derive structures.
Revise your first year notes.
3. Shapes of Molecules
Understanding the shapes of molecules is an important step in being able to discuss and predict chemical properties. Although here we discuss the shapes of “simple” molecules, this topic has also important applications in the understanding of the behavior of much larger molecules, e.g the shape of macromolecules in biology is often important with respect to their biochemical function
Test Question Draw the Lewis Structure of the Nitrato ion NO3
-.
How many bonds, how many bonds?
What is the nitrogen-oxygen bond order?
Are there possible resonance structures, can you draw them?
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Bond Order
Single bond - first order Double bond = second order Triple bond third order
Bond order is a measure of the number of bonding electron pairs between atoms. Single bonds have a bond order of 1, double bonds have a bond order of 2 and triple bonds (the maximum number) have a bond order of 3. A fractional bond order is possible in molecules and ions that have resonance structures. In the example of ozone, the bond order would be the average of a double bond and a single bond or 1.5 (3 divided by 2). As the bond order becomes larger, the bond length becomes smaller.
Remember atoms in the 3rd period or below e.g. P, I do not always obey the Octet rule!
The Shape of Ammonia (NH3) – VSEPR is important here.
N HH
H
Lone Pair
Lewis Structure
We have to consider repulsions between the lone pair and valence electronsactual structure:
N
HH
H
H-N-H angle is just slightly smaller than 109.50
The Nitrogen atom is Pyramidal
But why isnt the NHN angle 900?
Ammonia is a polar molecule with N carrying a partial negative charge. Molecular shape is important with respect to determining if a molecule is polar or not.
4. Valence Bond Theory
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Look at Valence Bond Theory (VBT)
The actual shape of NH3 is trigonal pyramidal (approximately tetrahedral minus one atom).
Hybridization of N = sp3N [He] 2s2 2p3
2s 2p
Hybridizationmix the orbitals -" like mixing together a red and white plant"
H HH
N [He] 2s2 2p3
H 1s1
We know that sp3 hybrids have a 109.50 angle
N
HH
H
N
HH
H
Molecular Structure of NH3 - cannot see the lone pair on N but there is a flattened lone pair
Compared to H20 The O in H2O has 2 bond pairs and 2 lone pairs. Two corners of the
tetrahedron are missing because they are occupied by lone pairs, not atoms. The shape is called bent. The H-O-H angle is less than NH3, due to the greater repulsions felt with two lone pairs
Other molecules with 2 bond plus 2 lone pairs include OF2, H2S and SF2. Bond angles vary, but all are significantly less than 109.50.
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Treat this as an exception to the octet rule.(An atom obeys the octet rule when it gains, looses or shares electrons to give an outer shell containing eight electrons with the configuration ns2np6). Many molecules such as neutral compounds of Boron simply do not contain enough valence electrons for each atom to be associated with eight electrons.
The Shape of BF3
B
F
F
F
Six electrons around the Boron
2s 2pB 2s2 2p F 2s2 2p5
sp2 this leaves an empty 2p orbital
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This leaves an unused "p orbital" perpendicular to the plane of BF3
F B
F
F
But if we want B to have an octet how can we achieve this?
A hybrid of 4 resonance structures is the best Lewis representation for the real strucure of BF3.
F
B
F F
F
B
F F FB
F
F FB
F
F
However... In this structure with a double bond the fluorine atom is sharing extra electrons with the boron. The fluorine would have a '+' partial charge, and the boron a '-' partial charge, this is inconsistent with the electronegativities of fluorine and boron.
Conclusion - the Octet Rule breaks down here.
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Evidence for a resonance structure comes from the B-F distances measured in the solid state. They are shorter by ~15 pm’s compared to the B-F distances in BF4
-. Generally as we move from a single bond towards a double bond our bond lengths shorten by approximately 15 ppm’s.
F
BF F
F
C-C Distances CH3CH2 155 ppm
CH=CH140 ppm
BF3 Resonance
Rehybridize the F’s to sp2
FBF
F
empty'p' on B
filled'p' on F
FB
F
F
empty'p' on B
filled'p' on F
The MO diagram is complex but the result for BF3 is one π-bond spread over 3 B-F links.
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To Summarize: BF3
The B atom has three bond pairs in its outer shell. Minimizing the repulsion causes this molecule to have a trigonal planar shape, with the F atoms forming an equilateral triangle about the B atom. The F-B-F bond angles are all 120°, and all the atoms are in the same plane.
BF3 reacts strongly with compounds which have an unshared pair of electrons which can be used to form a bond with the boron:
BF3 – Lewis Acid – electron pair acceptor. NH3 – Lewis Base – electron pair donor.
5. Lewis Acids and Bases
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6. σ versus Π-bonding
Ethene, C2H4, sp2
HH
H H
p orbital not used in hybridization
The three sp2 hybrid orbitals arrange themselves as far apart as possible - which is at 120° to each other in a plane. The remaining p orbital is at right angles to them.
C-H overlap to give sigma bonds.
Nodal Planefn = 0 (wave function)i.e. no electron density
Two lobes one with apositive sign the otherwith a negative sign gothough a node.
The two carbon atoms and four hydrogen atoms would look like this before they joined together:
The various atomic orbitals which are pointing towards each other now merge to give molecular orbitals, each containing a bonding pair of electrons.
σ orbital – no nodal planes
Π orbital one nodal plane containing the nuclei.
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Notice that the p orbitals are so close that they are overlapping sideways.
This sideways overlap also creates a molecular orbital, but of a different kind. In this one the electrons aren't held on the line between the two nuclei, but above and below the plane of the molecule. A bond formed in this way is called a pi bond.
Π-orbital above and below nodal plane
The σ -bond is protected but the Π -bond is sticking up and is not protected by the rest of the molecule, hence these electrons are exposed to reacting species and it is why alkenes and alkynes are reactive.
7. Relationship between Lewis Structure, VBT,VSEPR Valence Shell Electron Pair Repulsion Theory (VSEPR) enables us to
predict the shape of the central atoms electron pairs and in turn the hybridization of the central atom.
Lewis Structure
Electron Pair Geometry (VSEPR) - non bonding electrons and bonded atoms
hybridization molecular geometry - only looks at shape ofatoms; not lone pairs
bond overlap
(VBT)
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Methane, Ammonia, Water
Electron pair = nonbonding electrons + bonded atoms
Molecular – only looks at shape of atoms; not lone pairs
Electron pair geometry: Tetrahedral Tetrahedral Tetrahedral
Molecular geometry : Tetrahedral Triangular pyramidal Bent/Angular
109.5 107.5 104.5
Number of Bonded Atoms and Lone Pairs on Central Atom
Shape (e-pair Geometry) Hybridization
2 Linear sp
3 Trigonal Planar/ triangular sp2
4 Tetrahedral sp3
5 Trigonal Bipyramidal sp3d
6 Octahedral sp3d2
Examples:1. H2O
- electron pair geometry = tetrahedral (2 lp and 2 bp)
- molecular shape = bent- O hybrization = sp3
H O H
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2. XeF4 (36 electrons)
Xe
F
F F
F
six pairs of electrons around Xelone pair geometry - octahedralXe = sp3d2 hybridizedF = sp3 hybridized
the lone pairs are far appart therefore the compound as a SQUARE PLANARmolecular geometry.
Xe
F F
F F
A typical midterm/exam question would be:
1. Draw the Lewis Structure of XeF4
2. Give (i) the molecular shape, (ii) the electron pair geometry at the central atom and (iii) the hybridization of the central atom.
Practice Exercise
Draw the Lewis structure of BrOF3.
Give its electron pair geometry, the hybridization of the central
atom and its molecular geometry.