![]() ![]() The application of this theory to a molecule depends on stoichiometry, the number of bond pairs, and the number of lone pairs on the central atom. Valence shell electron pair repulsion (VSEPR) theory is the preferred choice when it comes to determining the geometry of a molecule.Īccording to this, the constituent atoms in a molecule arrange themselves in a way that the repulsion arising from the valence shell electrons on all atoms is minimum. Steric number = number of sigma bonds + number of lone pairs Steric numberįor F2, steric number = 1 + 3 = 4, corresponding to sp3. The already fully filled sp3 hybrid orbitals on each atom are the lone pairs.Īnother way of determining the hybridization of the central atom is by using the following formula. This results in the formation of a single bond (also called sigma bond) between the half-filled sp3 hybrid orbitals of two fluorine atoms. This required electron comes from the half-filled sp3 hybrid orbital from the second fluorine atom. Three of these hybrid orbitals are fully filled (with two electrons in each), and the fourth one is half-filled (i.e., it has one electron) and thus can accept one more. In a fluorine atom, the valence orbitals-2s, 2px, 2py, and 2pz-hybridize together to form four identical sp3 orbitals, all of which have the same energy. Let us look at the ground state electronic configuration of fluorine atom(s) in F2 in terms of the orbitals. When the valence orbitals on two atoms in a molecule overlap by sharing a pair of electrons, a chemical bond is said to form between the two atoms. Anyway, for your information, you may refer to the formula for the formal charge as provided below.įormal charge (FC) = Valence electrons – 0.5*bonding electrons – non-bonding electrons In fact, each fluorine atom has a formal charge of 0 on it. This low bond energy of fluorine explains why it is reactive.Īnyway, back to the main topic of the article! The fluorine molecule is neutral, i.e., there is no charge on it. ![]() This is approximately half the energy required to break a carbon–carbon single bond. Both fluorine atoms share one pair of electrons and hence have a single covalent bond between them.ĭo you know how much energy it would take to break this bond? It is 157 KJ/mol. Now, let us construct a skeleton of the F2 molecule on the basis of the information presented in step 2. The three unshared pairs of electrons on each fluorine atom are called the lone pairs. There is no rocket science going into determining that both fluorine atoms can share one pair of electrons and be happy and satisfied with their individual octets achieved! The case of fluorine is really one of the simplest ones. To achieve the octet, each atom needs one more electron. Thus, each fluorine atom has 7 valence electrons. There are 2 electrons in its K shell and 7 electrons in its L shell. ![]() The atomic number of fluorine is 9 therefore, it possesses 9 electrons in its neutral atomic form. Start by calculating the number of valence electrons in each atom of F2 and see how many more electrons each fluorine atom needs to form an octet. Let us take a look at the chemical bonding represented by lewis structure in F2. This is exactly why they are called “noble.” Noble gases already have completely filled valance shells, so they do not need to bond/react with any other atoms/molecules. Hydrogen is an exception, though it seeks a duplet, not octet, because it has only one electron in its K shell, and thus needs only one more to achieve the maximum capacity of K shell. Several atoms tend to seek eight electrons in their valence shell through chemical bonding this is referred to as the octet rule and is reflected in the Lewis structure of a molecule. In a typical Lewis structure, each valence electron is represented as a dot, and a covalent bond between two atoms (formed as a result of sharing of two electrons) is represented as a line. It is to be noted though that this theory about the electronic structure is quite primitive and most limited. The valence electrons in each atom are the ones that participate in the bonding, and hence they are the only ones displayed in the Lewis structures. The Lewis theory of chemical bonding helps us visualize the arrangement of atoms-how they are attached or bonded-in molecules. ![]()
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