The yellow groups are non-bonding electron pairs. When there are non bonding pairs, other shape can arise (see the above chart). The hybridization is sp 3d or sp 3d 2, with a trigonal bipyramidal or octahedral shape respectively. Recall that some elements, especially sulfur and phosphorus, can bond with five or six groups. When they bond, carbon has no non-bonding pairs, but sulfur has one. Carbon has four valence electrons, sulfur has six. The difference is in their valence shells. Carbon dioxide is linear, and sulfur dioxide is bent. Both have two oxygen atoms attached with double covalent bonds. Take a look at sulfur dioxide (SO 2) and carbon dioxide (CO 2). Both are attached to two groups, so it depends on how many non-bonding pairs the central atom has. Some elements will have a bent shape, others have a linear shape. No non-bonding pairs causes sp hybridization, leading to a linear shape. The two remaining p orbitals sit above and to the side of the beryllium atom (they are empty). This time, one 2s and one 2p orbital combine to form two sp hybrid orbitals. An example is beryllium chloride, which has two chlorine atoms situated on opposite sides of a beryllium atom. These repel each other at cos -1(-1) = 180°, forming a linear molecule. The remaining p orbital is empty and sits above and below the plane of the molecule.īeryllium, on the other hand, forms only two pairs of valence electrons. A different set of hybrid orbitals is formed in this molecule: the 2s and two 2p orbitals combine to form three sp 2 hybrid orbitals. Each fluorine atom is separated by cos -1(-1/2) = 120°. In boron trifluoride (BF 3), there are only three electron pairs which repel each other to form a flat plane. This occurs in electron deficient atoms such as boron and beryllium, which don't conform to the octet rule (they can have 6 and 4 valence electrons respectively). In some molecules, there are less than four pairs of valence electrons. Linear and Planar Shapes Electron-Poor Atoms The two lone pairs both push the bonds closer together, giving a smaller angle than in ammonia. The angle between the hydrogen atoms is 104.5°, which is what we expect from our model. The electron pairs repel each other into a tetrahedral shape. In water (H 2O) there are two hydrogen atoms and two lone pairs, again making four groups in total. Testing this assumption with water provides further evidence. The lone pair occupies more space than a bonding pair, decreasing the angles. It is trigonal pyramidal because it is attached to four "things": the three hydrogens and a non-bonding pair of electrons (to fulfill nitrogen's octet). Ammonia (NH 3) is not trigonal planar, however. For example, methane (CH 4) is tetrahedral-shaped because the carbon is attached to four hydrogens. The number of groups is how many atoms or electron pairs are bonded to the central atom. Those "things" can be other atoms or non-bonding pairs of electrons. The hybridization is determined by how many "things" are attached to the central atom. Table of Geometries Orbital Hybridization Modern methods of quantitatively calculating the most stable (lowest energy) shapes of molecules can take several hours of supercomputer time, and is the domain of computational chemistry. This can allow us to predict the geometry of similar molecules, making it a fairly useful model. The VSEPR model is by no means a perfect model of molecular shape! It is simply a system which explains the known shapes of molecular geometry as discovered by experiment. Wikipedia has related information at VSEPR Within this model, the AXE method is used in determining molecular geometry by counting the numbers of electrons and bonds related to the center atom(s) of the molecule. A scientific model, called the VSEPR (valence shell electron pair repulsion) model can be used to qualitatively predict the shapes of molecules. There are several different methods of determining molecular geometry. This repulsion causes covalent molecules to have distinctive shapes, known as the molecule's molecular geometry. Being mutually negatively charged, the electron pairs repel the other electron pairs and attempt to move as far apart as possible in order to stabilize the molecule. Covalent molecules are bonded to other atoms by electron pairs.
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