Presentation on theme: "Chemical bond: Sructure and shape. VSEPR Teory: valence shell electron pair repulsion teory Digunakan untuk memprediksi struktur molekul Pasangan elektron."— Presentation transcript:
VSEPR Teory: valence shell electron pair repulsion teory Digunakan untuk memprediksi struktur molekul Pasangan elektron valensi yang memiliki muatan negatif akan saling berjauhan sedemikian sehingga tolakan antar pasangan elektron seminimal mungkin.
Chemical structure and shape Chemical structure Molecular shape
Rule 1. First find the number of bonding pairs of electrons in the molecule. The number of bonding pairs of electrons in the molecule NH3 can be seen in the formula. There must be three bonding pairs of electrons holding the three hydrogens onto the nitrogen. Rule 2. Find the number of valence electrons (electrons in the outer energy level) on an atom of the central atom (The one of which there is only one.) Nitrogen is in group V, so the nitrogen has five electrons in the outer energy level. Rule 3. Find the number of lone pairs on the central atom by subtracting the number of bonding pairs (3) from the valence electrons (5) to find the number of electrons (2) that will make up lone pairs of electrons. Divide this number by 2 to find the number of lone pairs, 2/2 = 1. Rule 4. Distribute all the electron pairs around the central atom and learn the angles they will make from molecules with no lone pairs. Rule 5. Learn that the repulsion between lone pairs of electrons is greater than the repulsion between bonding pairs, and subtract 2o from the bond angles for every lone pair. Rule 6. Learn the names of the shapes. The shapes are named form the position of the atoms and not the position of the orbitals. The Shape of Ammonia, NH 3
Hydrogen chloride, HCl HCl is composed of only 2 atoms, 1 atom of hydrogen and 1 atom of chlorine covalently bonded. HCl is, therefore, diatomic. All diatomic molecules are linear in shape. H-Cl is linear in shape
Hydrogen cyanide, HCN HCN is composed of 3 atoms, 1 atom of hydrogen, 1 atom of carbon and 1 atom of nitrogen covalently bonded. Carbon, C, is the central atom in the molecule. Carbon has 4 valence electrons (electrons that can be used in bonding). 1 of carbon's valence electrons will be used to form a covalent bond with hydrogen. 3 of carbon's valence electrons will be used to form 3 covalent bonds with nitrogen ( a triple bond). The central carbon atom therefore has no lone pairs of electrons. The bonding pairs will repel each other as much as possible, so the molecule will be linear. HCN is linear in shape.
Aluminium chloride, AlCl 3 AlCl3 is composed of 4 atoms, 1 atom of aluminium and 3 atoms of chlorine. Aluminium, Al, is the central atom in the molecule. Aluminium has 3 valence electrons (electrons that can be used in bonding). Each of the aluminium's valence electrons will be used to form a covalent bond with each chlorine atom. The central aluminium atom will therefore have no lone pairs of electrons and 3 bonding pairs of electrons. The bonding pairs of electrons will repel each other as much as possible, so the molecule will be trigonal planar. AlCl3 is trigonal planar in shape.
Methane, CH 4 CH4 is composed of 5 atoms, 1 atom of carbon and 4 atoms of hydrogen covalently bonded. Carbon, C, is the central atom in the molecule. Carbon has 4 valence electrons (electrons that can be used in bonding). Each of carbon's 4 valence electrons will form a bonding pair with 1 of hydrogen's electrons. The central carbon atom will therefore have no lone pairs of electrons and 4 bonding pairs of electrons, The bonding pairs of electrons will repel each other as much as possible, so the molecule will be tetrahedral. CH4 is tetrahedral in shape.
Ammonia, NH 3 NH3 is composed of 4 atoms, 1 nitrogen atom and 3 hydrogen atoms covalently bonded. Nitrogen, N, is the central atom in the molecule. Nitrogen has 5 valence electrons. 3 of nitrogen's valence electrons will be used to form bonding pairs of electrons with hydrogen (3 covalent bonds). These bonding pairs repel each other equally and will try to get as far away from each other as possible. 2 of nitrogen's valence electrons will be unused for bonding, these are a lone pair of electrons. lone pair-lone pair repulsion is greater than bonding pair-bonding pair repulsion, so the lone pair pushes the bonding pairs closer together than in a tetrahedral arrangement of the 'electron clouds'. This distorted tetrahedral arrangement is called trigonal pyramidal. NH3 is trigonal pyramidal in shape. In this representation, the solid triangles represent bonds coming out of the plane of the screen, the broken lines represented a bond going behind the plane of the screen.
Water, H 2 O Water is composed of 3 atoms, 1 atom of oxygen and 2 atoms of hydrogen covalently bonded. Oxygen, O, is the central atom. Oxygen has 6 valence electrons. 2 of oxygen's electrons will be used to form bonding pairs of electrons with hydrogen (2 covalent bonds). These bonding pairs repel each other equally and will try to get as far away from each other as possible. 4 of oxygen's valence electrons will not be used for bonding, these will remain as 2 lone pairs of electrons. These lone pairs of electrons repel each other equally and will try to get as far away from each as possible. lone pair-lone pair repulsion is greater than lone pair-bonding pair or bonding pair-bonding pair repulsion, so the lone pairs of electrons push the bonding pairs of electrons closer together than in a tetrahedral arrangement of the 'electron clouds'. This distorted tetrahedral arrangement is call bent. H2O is bent in shape. In this representation, solid lines represent bonds that are in the same plane as the screen.
berilium khlorida BeCl2 dan timah (II) khlorida SnCl2?
Pada CHCH, tentukan rumus lewis dan tentukan orbital yang terlibat, membentuk hibridisasi apa? Apakah yang dimaksud isomer optik? Secara eksperimen bagaimana cara membuktikannya?
BONDING IN METHANE AND ETHANE Methane, CH 4 The simple view of the bonding in methane You will be familiar with drawing methane using dots and crosses diagrams, but it is worth looking at its structure a bit more closely. There is a serious mis-match between this structure and the modern electronic structure of carbon, 1s 2 2s 2 2p x 1 2p y 1. The modern structure shows that there are only 2 unpaired electrons for hydrogens to share with, instead of the 4 which the simple view requires. You can see this more readily using the electrons-in-boxes notation. Only the 2-level electrons are shown. The 1s 2 electrons are too deep inside the atom to be involved in bonding. The only electrons directly available for sharing are the 2p electrons. Why then isn't methane CH 2 ?
There is only a small energy gap between the 2s and 2p orbitals, and an electron is promoted from the 2s to the empty 2p to give 4 unpaired electrons. The extra energy released when these electrons are used for bonding more than compensates for the initial input. The carbon atom is now said to be in an excited state. Now that we've got 4 unpaired electrons ready for bonding, another problem arises. In methane all the carbon-hydrogen bonds are identical, but our electrons are in two different kinds of orbitals. You aren't going to get four identical bonds unless you start from four identical orbitals. Promotion of an electron
Hybridisation The electrons rearrange themselves again in a process called hybridisation. This reorganises the electrons into four identical hybrid orbitals called sp 3 hybrids (because they are made from one s orbital and three p orbitals). You should read "sp 3 " as "s p three" - not as "s p cubed". sp 3 hybrid orbitals look a bit like half a p orbital, and they arrange themselves in space so that they are as far apart as possible. You can picture the nucleus as being at the centre of a tetrahedron (a triangularly based pyramid) with the orbitals pointing to the corners. For clarity, the nucleus is drawn far larger than it really is.
What happens when the bonds are formed? Remember that hydrogen's electron is in a 1s orbital - a spherically symmetric region of space surrounding the nucleus where there is some fixed chance (say 95%) of finding the electron. When a covalent bond is formed, the atomic orbitals (the orbitals in the individual atoms) merge to produce a new molecular orbital which contains the electron pair which creates the bond. Four molecular orbitals are formed, looking rather like the original sp 3 hybrids, but with a hydrogen nucleus embedded in each lobe. Each orbital holds the 2 electrons that we've previously drawn as a dot and a cross. The principles involved - promotion of electrons if necessary, then hybridisation, followed by the formation of molecular orbitals - can be applied to any covalently-bound molecule
The shape of methane When sp 3 orbitals are formed, they arrange themselves so that they are as far apart as possible. That is a tetrahedral arrangement, with an angle of 109.5°. Nothing changes in terms of the shape when the hydrogen atoms combine with the carbon, and so the methane molecule is also tetrahedral with 109.5° bond angles.
The formation of molecular orbitals in ethane Ethane isn't particularly important in its own right, but is included because it is a simple example of how a carbon-carbon single bond is formed. Each carbon atom in the ethane promotes an electron and then forms sp 3 hybrids exactly as we've described in methane. So just before bonding, the atoms look like this: The hydrogens bond with the two carbons to produce molecular orbitals just as they did with methane. The two carbon atoms bond by merging their remaining sp 3 hybrid orbitals end-to-end to make a new molecular orbital. The bond formed by this end-to-end overlap is called a sigma bond. The bonds between the carbons and hydrogens are also sigma bonds. In any sigma bond, the most likely place to find the pair of electrons is on a line between the two nuclei.
The shape of ethane around each carbon atom The shape is again determined by the way the sp 3 orbitals are arranged around each carbon atom. That is a tetrahedral arrangement, with an angle of 109.5°. When the ethane molecule is put together, the arrangement around each carbon atom is again tetrahedral with approximately 109.5° bond angles. Why only "approximately"? This time, each carbon atoms doesn't have four identical things attached. There will be a small amount of distortion because of the attachment of 3 hydrogens and 1 carbon, rather than 4 hydrogens.
Hybridisation of ethena In the case of ethene, there is a difference from, say, methane or ethane, because each carbon is only joining to three other atoms rather than four. When the carbon atoms hybridise their outer orbitals before forming bonds, this time they only hybridise three of the orbitals rather than all four. They use the 2s electron and two of the 2p electrons, but leave the other 2p electron unchanged.
The new orbitals formed are called sp 2 hybrids, because they are made by an s orbital and two p orbitals reorganising themselves. sp 2 orbitals look rather like sp 3 orbitals that you have already come across in the bonding in methane, except that they are shorter and fatter. The three sp 2 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.
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. These are sigma bonds - just like those formed by end-to-end overlap of atomic orbitals in, say, ethane. http://www.chemguide.co.uk/basicorg/bonding/ethene.html
The p orbitals on each carbon aren't pointing towards each other, and so we'll leave those for a moment. In the diagram, the black dots represent the nuclei of the atoms. 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. For clarity, the sigma bonds are shown using lines - each line representing one pair of shared electrons. The various sorts of line show the directions the bonds point in. An ordinary line represents a bond in the plane of the screen (or the paper if you've printed it), a broken line is a bond going back away from you, and a wedge shows a bond coming out towards you.
The shape of ethene The shape of ethene is controlled by the arrangement of the sp 2 orbitals. Notice two things about them: They all lie in the same plane, with the other p orbital at right angles to it. When the bonds are made, all of the sigma bonds in the molecule must also lie in the same plane. Any twist in the molecule would mean that the p orbitals wouldn't be parallel and touching any more, and you would be breaking the pi bond. There is no free rotation about a carbon-carbon double bond. Ethene is a planar molecule. The sp 2 orbitals are at 120° to each other. When the molecule is constructed, the bond angles will also be 120°. (That's approximate! There will be a slight distortion because you are joining 2 hydrogens and a carbon atom to each carbon, rather than 3 identical groups.)