Electronega+vity Review

Transcription

1 Electronega+vity Review Remember from the first course that electronega+vity is an es+mate of how atoms pull electrons towards themselves in a molecule. The higher the electron affinity, the more the element can pull electrons towards itself.

2 Dipoles In methyl lithium, C 3 Li, an atom of lithium is bonded directly to an atom of carbon. Lithium has an electronega+vity of 1.0 whereas carbon has an electronega+vity of 2.5. This means that carbon is beier at pulling electron density to itself than lithium. The splijng of electric charge along a bond or molecule is called a dipole. The signs δ + and δ are used to show a par+al charge. The dipole makes methyl lithium a polar molecule. methyl lithium methanol In methanol, C 3 O, carbon (2.5) becomes δ+ and oxygen (3.5) is δ. In this case, oxygen is also pulling electron density from hydrogen. ydrogen atoms are osen associated with a strong dipole since they are so small and lose their electron density easily. The resul+ng dipole lies between the C 3 group and the atom making methanol a polar molecule. In ethane, there is no dipole as both atoms on either side of the C C bond are carbon and have the same electronega+vity. Addi+onally, carbon and hydrogen have very similar electronega+vi+es (2.5 and 2.1, respec+vely) and so C dipole is very small. The lack of any strong dipole makes ethane a non- polar molecule. Dipoles are important because they are responsible for many chemical and physical proper+es of compounds.

3 Molecular Shape Understanding the shapes of molecules is very important for understanding their chemical and physical proper+es. The shape of a molecule helps us to understand how dipoles will be arranged in a molecule and how the molecules will interact with each other in liquids, solids and in chemical reac+ons. There are various experimental methods we can use to find the shapes of molecules. Diatomic molecules can only be linear, but triatomic and higher order molecules can make different shapes (geometries). We can combine these experimentally determined shapes with our knowledge of atomic orbitals to make theories of how molecules form.

4 ow Atoms Make Molecules Valence Bond Theory We already know that atoms form molecules by sharing pairs of electrons but how does this actually happen? If the orbitals of two different atoms approach closely enough, they can overlap. This will result in greater, nega+vely charged, electron density between the two posi+vely charged nuclei and the two nuclei will be pulled together more strongly. This airac+on between the nuclei and the central electron density lowers the energy of both atoms and a molecule is formed. This diagram shows the case for the s- orbitals of two hydrogen atoms. Red arrows show repulsion between like charges and blue arrows show airac+on between opposite charges. The nuclei will approach each other un+l the long range airac+ve forces balance the short range repulsive forces. The distance between the nuclei when all forces balance is called the bond length.

5 The Lennard- Jones Poten+al Forming a bond releases energy, breaking a bond absorbs energy. If we plot the poten+al energy of two atoms, with respect to the separa+on of their nuclei, we get an important curve known as the Lennard- Jones poten+al. The shape of the curve illustrates how the atoms are airacted to each other at large separa+ons but repel each other at small separa+ons, with a minimum at the bond length.

6 Electron repulsion VSEPR Theory Overlapping orbitals explains the forma+on of chemical bonds, but how does it affect the shape of the molecule? Although single electrons can form pairs if they have opposite spins, pairs of electrons will always repel other pairs. Beryllium hydride, Be 2 has two pairs of electrons in two bonds. The minimum repulsion energy is found when the electron pairs are on opposite sides of the beryllium atom. This gives Be 2 a linear shape. Repulsion between electron pairs pushes bonds apart. A linear molecule has minimum repulsion between bonding pairs. This way of explaining molecular shapes is known as Valence Shell Electron Pair Repulsion theory, or VSEPR theory. Now we understand why Be 2 is linear but why is 2 O, an apparently similar molecule, bent? To understand that, first we have to understand the shape of methane, C 4.

7 The Shape of Methane We know that C 4 has four pairs of electrons, one for each C bond. ow can we minimise the repulsion energy? The answer is to arrange the pairs in a tetrahedron. This arrangement equalises all the repulsion forces, in other words, it gives us the minimum repulsion energy between the four pairs of electrons in the four bonds. Any other arrangement of electron pairs will result in 1) unequal repulsion forces so the bonding pairs will rearrange themselves into a tetrahedron and 2) bond angles of less than these configura+ons have stronger repulsion, which makes them higher energy than a tetrahedron and therefore less stable !80 unequal repulsion => unstable Any devia+on from exactly 180 /90 will result in unequal forces => unstable.

8 The Shape of Water Now we can understand the shape of water. Remember that, in water, oxygen has four pairs of electrons in its outer shell, two bonding pairs and two lone pairs. O Since we have four pairs of electrons, the most stable arrangement of electron pairs is a tetrahedral arrangement. owever, the ^ angle is less than because the lone pairs repel each other more strongly than the bonding pairs. F B F F N F F F We can also understand the why boron triflouride (BF 3 ) is trigonal planar while ammonia (N 3 ) is pyramidal.

9 Electron groups When we use VSEPR to predict the shape of a molecule, we do not need to care about the difference between lone pairs, single, double and triple bonds. We only need to consider electron groups. One lone pair is one electron group. One single bond, one double bond or one triple bond is one electron group. Now lets consider the shape of ethane, ethene, ethyne and ethanol: Ethane Both carbons of ethane are surrounded by 4 single bonds, i.e. 4 electron groups. We therefore expect these carbons to be tetrahedral. C Ethene C Both carbons of ethene are surrounded by 2 single bonds and one double bond for a total of 3 electron groups. We therefore expect these carbons to be trigonal planar.

10 Electron Groups cont. Ethyne Both carbon atoms in ethyne are surrounded by one single bond and one triple bond for a total of two electron groups. We therefore expect these carbons to be linear. C Ethanol O O The oxygen atom in ethanol is surrounded by two single bonds and two lone pairs for a total of four electron groups. This means the electron groups will be arranged tetrahedrally (but the tetrahedron will be slightly distorted because lone pairs repel other electron groups more strongly than bonding pairs). owever, the molecular shape at the oxygen atom is bent, not tetrahedral, because we only consider atoms when naming a molecular shape.

11 The VSEPR / Atomic Orbitals Problem VSEPR is simple and easy to use but the shapes it predicts do not seem to match the shapes of atomic orbitals. s orbital p orbitals We know that covalent bonds are made by overlapping orbitals but a quick look at the orbitals of carbon ([e]2s 2 2p 2 ) do not show any easy way to make a tetrahedral molecule. energy 2p x C 2p y 2p z 2s 1s Y => X => Z C So how can we make the shapes that VSEPR predicts using atomic orbitals? X Not the shape of methane.

12 Superimposing Orbitals Remember that orbitals behave like 3- D waves and waves can be superimposed. When calcula+ng the shapes of atomic orbitals using the Scrodinger equa+on, we can mathema+cally superimpose one orbital onto another and make a hybrid orbital. 1D waves: + = = 1D hybrid wave 3D waves: + = 3D hybrid wave

13 Orbital ybridisa+on sp 3 Orbitals The atomic orbital geometry problem was solved by combining the equa+ons for one s orbital and three p orbitals to make four sp 3 hybrid orbitals (s x p x p x p = sp 3 ). These hybrid orbitals have the right orienta+on to make a tetrahedral molecule. By overlapping four sp 3 orbitals of a carbon atom with the s orbitals from four hydrogen atoms, we can get the tetrahedral shape of methane. p y s p z p x Four sp 3 orbitals Note, we never lose or gain orbitals when making combina+ons four atomic orbitals (e.g. s + (3 x p)) always make four hybrid orbitals (e.g. 4 x sp 3 orbitals). (a single sp 3 orbital can be thought of as made from ¼ s, and ¾ p orbitals).

14 Orbital ybridisa+on sp 2 Orbitals VSEPR tells us that BF 3 should be triangular but how can we make orbitals that fit that shape? If we mix an s orbital with only two p orbitals, we can make three sp 2 orbitals with one p orbital les over (remember: three atomic orbitals make three hybrid orbitals). If we place a fluorine atom at the end of each sp 2 orbital, and leave the p orbital empty, we get the shape of boron trifluoride (we already know from the octet rule that boron is missing two electrons in BF 3 so there is no problem with having an empty p orbital in this case). 3 sp 2 orbitals + 1 p orbital top view BF 3 3 boron sp 2 orbitals overlap with p orbitals from 3 fluorine atoms

15 sp 2 Orbitals and Double Bonds. In ethene, 2 C=C 2, two carbon atoms are joined by a double bond. For VSEPR, a double bonds is only one electron group so we can predict that ethene will have two trigonal planar carbon atoms. In a double bond, the sp 2 orbitals of two carbon atoms overlap to make a σ bond (a bond that lies directly along the internuclei axis) and the two p orbitals overlap to make a π bond (a bond that lies above and below the internuclei axis). PuJng a hydrogen atom at the end of the four outside sp 2 orbitals makes ethene. σ- bond C C π- bond Note that half of the electron density of the π bond is above the internuclear axis and half is below.

16 sp Orbitals and Triple Bonds Ethyne, C C, has two carbon atoms joined by a triple bond. To make a triple bond, we combine one s orbital and one p orbital to make two sp orbitals, with two p orbitals les over. The two sp orbitals overlap between the carbon atoms to make a σ bond and the remaining p orbitals overlap to make two π bonds. PuJng a hydrogen atom at the end of the two outside sp orbitals gives us ethyne. C C

17 Valence Bond Theory Now we can imagine that atoms make molecules by simply overlapping their orbitals, whether the orbitals are normal atomic orbitals or hybrid orbitals. Bonding electrons spend all their +me in orbitals between atoms (bonding orbitals) and lone pair electrons spend all their +me in an orbital on only one atom (a non- bonding orbital). This way of thinking is called valence bond theory. It is quite easy to imagine how valence bond theory works so it is a very popular way of thinking about molecular bonding. It is most commonly used in organic chemistry.

18 (Problems with Valence Bond Theory) This theory of hybrid orbitals explained many things about about chemistry but there were s+ll some problems. For instance, no one could find any experimental evidence to prove hybrid orbitals really exist. For this reason, valence bond theory has been replaced by the more accurate but more complex molecular orbital theory. Even so, valence bond theory is a very simple concept and so chemists osen find it useful when discussing molecular geometry and electronic proper+es. There are some things that valence bond theory cannot explain, however. For instance, an analy+cal technique called spectroscopy tells us that dioxygen (O 2 ) has one single bond and two unpaired electrons but valence bond theory predicts that O 2 would have a double bond and be sp 2 hybridised. There are many other chemical and spectroscopic observa+ons that cannot be explained by valence bond theory. valence bond theory predicts: O O but reality is: O O O 2 is a diradical (it has two unpaired electrons) and does not have a double bond.