What is lowest energy structure, in terms of formal charge?
What is lowest energy structure, in terms of formal charge?
Smaller the formal charge on the atoms, lower is the energy of the str
It is true that structure with the least formal charge should be lower in energy and thereby be the better Lewis structure.
Consider the molecule H2SO4. There are 3 possible Lewis structures for this molecule. We can determine which is better by determining which has the least formal charge.
Formula to calculate formal charges : V - (N + B/2)
Where V is the number of valence electrons of the atom in the free atom; N is the number of non-bonding electrons on this atom in the molecule; and B is the total number of electrons shared in covalent bonds with other atoms in the molecule. The basic Lewis structure of H2SO4 with formal charge on each atom is shown below.
The 3 possible structures for H2SO4 is shown below.
The bottom resonance structure is the best because all the formal charges are zero
The formal charge is the difference in the number of valence electrons in the atom and the number of valence electrons in the Lewis structure.
The formal charge of any atom in a molecule can be calculated by the following equation:
FC = V-N-(B/2)
Where V is the number of valence electrons of the atom in the free atom; N is the number of non-bonding electrons on this atom in the molecule; and B is the total number of electrons shared in covalent bonds with other atoms in the molecule.
For example: In O3 molecule
Formal charge helps to select most stable structure, i.e. , the one with least energy out of the different possible Lewis structures. The most stable is the one which has the smallest formal charge on the atoms. For example if we take the case of SO2 molecule.
Four resonating structures are possible for SO2. Out of the four, structure IV is most favourable since none of the atom is having any formal charge. Structure I and III are then less stable than IV but more stable than II since they have full octet around all the atoms but two of the atoms in each structure have formal charge. II structure is least stable because in this structure S does not have complete octet and also all the three atoms are having formal charge.
Formal Charge The formal charge is an idea of accounting for the distribution of electrons in an atom. This can help in two ways.
1. It can help us decide which of several Lewis dot structures is closest to representing the properties of the real compound.
2. It can help us envision where there might be regions of positive or negative charge in a molecule.
First, how does it help to decide between different structures? Our general rule is that the best structure minimizes the formal charges. This is because minimizing the formal charges leads to the electrons being most evenly distributed about the different atomic centers in a molecule. Having electrons concentrated in one area will lead to regions of negative charge. The atoms that are now "missing" electrons will be positive in charge. Separating positive and negative charges costs energy and thus we conclude that the lowest energy (best structure) would minimize having separated charges.
How do we find these charges? We look at how many valence electrons the atoms "has" in the molecule compared to how many it has on its own. It is important to know that this is a very general idea that grossly over simplifies the quantum mechanics. The electrons in a molecule have no memory of where they came from or to which atom they "belong". They are simply spread throughout the molecule. None the less, these simple ideas can help us to arrive at the best structures as well as understand something about the charge distributions.
In a molecule, we assign each atom a formal charge. This charge is the number of electrons it had as valence electrons minus the number it "has" in the molecule. The number it has in the molecule is a combination of the lone pair electrons and the shared bonding electrons. For each atom, we will count all of the lone pair electrons but only half of the bonded electrons (as they are shared). This is easiest to account for by just counting the number of bonds. So the formal charge is
Formal charge = Valence Electrons - [Lone Pair Electrons + (# of Bonds)]
So for example if we look at CO2 each oxygen has two lone pairs (4 electrons) and 2 bonds (double bond). Oxygen has 6 valence electrons. So the formal charge on each oxygen atom will be 6-(4+2)=0.
It is true that structure with the least formal charge should be lower in energy and thereby be the better Lewis structure.Consider the molecule H2SO4. There are 3 possible Lewis structures for this molecule. We can determine which is better by determining which has the least formal charge.
Formula to calculate formal charges : V - (N + B/2)
Where V is the number of valence electrons of the atom in the free atom; N is the number of non-bonding electrons on this atom in the molecule; and B is the total number of electrons shared in covalent bonds with other atoms in the molecule.
The basic Lewis structure of H2SO4 with formal charge on each atom is shown below.
The 3 possible structures for H2SO4 is shown below.
The bottom resonance structure is the best because all the formal charges are zero
The formal charge is the difference in the number of valence electrons in the atom and the number of valence electrons in the Lewis structure.
The formal charge of any atom in a molecule can be calculated by the following equation:
FC = V-N-(B/2)
Where V is the number of valence electrons of the atom in the free atom; N is the number of non-bonding electrons on this atom in the molecule; and B is the total number of electrons shared in covalent bonds with other atoms in the molecule.
For example: In O3 molecule
Formal charge helps to select most stable structure, i.e. , the one with least energy out of the different possible Lewis structures. The most stable is the one which has the smallest formal charge on the atoms. For example if we take the case of SO2 molecule.
Four resonating structures are possible for SO2. Out of the four, structure IV is most favourable since none of the atom is having any formal charge. Structure I and III are then less stable than IV but more stable than II since they have full octet around all the atoms but two of the atoms in each structure have formal charge. II structure is least stable because in this structure S does not have complete octet and also all the three atoms are having formal charge.
The formal charge is an idea of accounting for the distribution of electrons in an atom. This can help in two ways.
1. It can help us decide which of several Lewis dot structures is closest to representing the properties of the real compound.
2. It can help us envision where there might be regions of positive or negative charge in a molecule.
First, how does it help to decide between different structures? Our general rule is that the best structure minimizes the formal charges. This is because minimizing the formal charges leads to the electrons being most evenly distributed about the different atomic centers in a molecule. Having electrons concentrated in one area will lead to regions of negative charge. The atoms that are now "missing" electrons will be positive in charge. Separating positive and negative charges costs energy and thus we conclude that the lowest energy (best structure) would minimize having separated charges.
How do we find these charges? We look at how many valence electrons the atoms "has" in the molecule compared to how many it has on its own. It is important to know that this is a very general idea that grossly over simplifies the quantum mechanics. The electrons in a molecule have no memory of where they came from or to which atom they "belong". They are simply spread throughout the molecule. None the less, these simple ideas can help us to arrive at the best structures as well as understand something about the charge distributions.
In a molecule, we assign each atom a formal charge. This charge is the number of electrons it had as valence electrons minus the number it "has" in the molecule. The number it has in the molecule is a combination of the lone pair electrons and the shared bonding electrons. For each atom, we will count all of the lone pair electrons but only half of the bonded electrons (as they are shared). This is easiest to account for by just counting the number of bonds. So the formal charge is
Formal charge = Valence Electrons - [Lone Pair Electrons + (# of Bonds)]
So for example if we look at CO2 each oxygen has two lone pairs (4 electrons) and 2 bonds (double bond). Oxygen has 6 valence electrons. So the formal charge on each oxygen atom will be 6-(4+2)=0.
