Answer
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Step 1:: Determine the total number of valence electrons.
For ClF^4 +, chlorine (Cl) is in the 17th group of the periodic table and has 7 valence electrons. Fluorine (F) is in the 17th group and has 7 valence electrons as well. Since there is one positive charge, we remove one electron from the ClF^4 + ion, giving us a total of 27 valence electrons (28 - 1).
Step 2:: Draw a rough sketch of the structure.
Place the central atom, chlorine, in the center and surround it with fluorine atoms. Since there are 4 fluorine atoms, arrange them in a tetrahedral shape around the chlorine atom.
Step 3:: Distribute the valence electrons.
First, put two electrons between each bond to form single covalent bonds. This uses up 8 electrons (4 bonds * 2 electrons/bond). Next, place the remaining electrons on the outer atoms (fluorine) as lone pairs until each fluorine atom has 8 electrons in its outer shell. This uses up an additional 16 electrons (4 fluorine atoms * 4 lone pairs * 2 electrons/lone pair).
Step 4:: Check the formal charges.
Calculate the formal charge for each atom by subtracting the number of valence electrons the atom would have in its neutral state from the total number of electrons around the atom. For chlorine, it has 7 valence electrons and is bonded to 4 fluorine atoms with a total of 8 electrons around it (4 bonds * 2 electrons/bond). The formal charge for chlorine is 7 - 8 = - 1. For fluorine, it has 7 valence electrons and is bonded to 1 chlorine atom with a total of 8 electrons around it (2 electrons from the bond and 6 lone pairs * 2 electrons/lone pair). The formal charge for fluorine is 7 - 8 = - 1. Since the formal charge for each atom is - 1, the Lewis structure is valid.
Step 5:: Draw the final Lewis structure.
Draw the Lewis structure for ClF^4 + with chlorine in the center, surrounded by 4 fluorine atoms in a tetrahedral shape, and include the formal charges on each atom.
Final Answer
The best Lewis structure for ClF^4 + is: \ Chem{ \overset{- 1}{\underset{+}{\ce{Cl}}}-\underset{-\charge}{\ce{F}}-\underset{-\charge}{\ce{F}}-\underset{-\charge}{\ce{F}}-\underset{-\charge}{\ce{F}} } where the formal charge for each atom is shown next to the atom.
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