Ch.10 - Molecular Shapes & Valence Bond TheoryWorksheetSee all chapters
All Chapters
Ch.1 - Intro to General Chemistry
Ch.2 - Atoms & Elements
Ch.3 - Chemical Reactions
BONUS: Lab Techniques and Procedures
BONUS: Mathematical Operations and Functions
Ch.4 - Chemical Quantities & Aqueous Reactions
Ch.5 - Gases
Ch.6 - Thermochemistry
Ch.7 - Quantum Mechanics
Ch.8 - Periodic Properties of the Elements
Ch.9 - Bonding & Molecular Structure
Ch.10 - Molecular Shapes & Valence Bond Theory
Ch.11 - Liquids, Solids & Intermolecular Forces
Ch.12 - Solutions
Ch.13 - Chemical Kinetics
Ch.14 - Chemical Equilibrium
Ch.15 - Acid and Base Equilibrium
Ch.16 - Aqueous Equilibrium
Ch. 17 - Chemical Thermodynamics
Ch.18 - Electrochemistry
Ch.19 - Nuclear Chemistry
Ch.20 - Organic Chemistry
Ch.22 - Chemistry of the Nonmetals
Ch.23 - Transition Metals and Coordination Compounds

Solution: Although I3- is known, F3- is not.Another classmate says F3-– does not exist because F is too small to make bonds to more than one atom. Is this classmate possibly correct?

Problem

Although I3- is known, F3- is not.

Another classmate says F3-– does not exist because F is too small to make bonds to more than one atom. Is this classmate possibly correct?

Solution

Let us start this problem by drawing the Lewis structure of F3-.

Step 1: The steps that are required to be taken:

  1. Firstly, we need to calculate the total number of valence electrons that would be present in the structure. The number of valence electrons per atom of each element will be based on its group number. F is in group 7A so it has 7 valence electrons.
  2. One extra electron should be added for the negative charge present.


Atom               Valence Electrons

           Fluorine       7 e- x 3 = (21 e) + 1 

                                                = 22 valence e- 

   A total of 22 valence electrons are present.  

Step 2 : The Lewis structures have to be drawn.

  • All the atoms in F3are equivalent; therefore, they can be drawn next to each other. 
  • You will then need to subtract two electrons for each bond. 
  • Then place the rest of the electrons in each fluorine atom to make sure the atom fulfills the octet
  • Finally, the two remaining electrons are placed in the central atom.
  • The formal charge for each atom is then calculated using the equation 

Formal Charge = group number – (bonds the element is making + non-bonding electrons)

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