Calculate the formal charge on each of the following Lewis structures a) b) Methoxide ion Methyl carbocation c) d) Tetrahydridoborate ion =ö-H An oxonium ion e) :o: Carbon monoxide :=0: H -H Sulfuric acid :ö: g) h) H :0: H -H Dimethyl sulfoxide Nitromethane
Electronic Effects
The effect of electrons that are located in the chemical bonds within the atoms of the molecule is termed an electronic effect. The electronic effect is also explained as the effect through which the reactivity of the compound in one portion is controlled by the electron repulsion or attraction producing in another portion of the molecule.
Drawing Resonance Forms
In organic chemistry, resonance may be a mental exercise that illustrates the delocalization of electrons inside molecules within the valence bond theory of octet bonding. It entails creating several Lewis structures that, when combined, reflect the molecule's entire electronic structure. One Lewis diagram cannot explain the bonding (lone pair, double bond, octet) elaborately. A hybrid describes a combination of possible resonance structures that represents the entire delocalization of electrons within the molecule.
Using Molecular Structure To Predict Equilibrium
Equilibrium does not always imply an equal presence of reactants and products. This signifies that the reaction reaches a point when reactant and product quantities remain constant as the rate of forward and backward reaction is the same. Molecular structures of various compounds can help in predicting equilibrium.
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![**Calculate the Formal Charge on Each of the Following Lewis Structures**
a) **Methoxide ion**
- Structure: H—C—O with an additional pair of electrons on oxygen, denoted with a negative charge.
- The Lewis structure shows hydrogen bonded to carbon, which is single-bonded to oxygen. Oxygen has three lone pairs.
b) **Methyl carbocation**
- Structure: [H—C—H]
- A positively charged carbon atom with three hydrogens attached by single bonds and no lone pairs.
c) **Tetrahydroborate ion**
- Structure: H—B—H with two additional hydrogens bonded to boron.
- Boron is central with four single bonds to hydrogen and a formal negative charge.
d) **An oxonium ion**
- Structure: [H—O—H]+ single-bonded to an additional hydrogen.
- Oxygen is central, with three single bonds and one lone pair, carrying a positive charge.
e) **Sulfuric acid**
- Structure: H—O—S—O with double-bonded oxygen atoms also bonded to sulfur. The terminal oxygens each have two lone pairs.
- Sulfur is central with four oxygens bonded, two single-bonded and two double-bonded. Hydrogens are single-bonded to two oxygens.
f) **Carbon monoxide**
- Structure: :C≡O:
- Carbon and oxygen are connected by a triple bond. Carbon has a lone pair and oxygen has two lone pairs.
g) **Dimethyl sulfoxide**
- Structure: (CH₃)₂—S=O
- Sulfur is bonded to two methyl groups and double-bonded to oxygen, which has two lone pairs.
h) **Nitromethane**
- Structure: CH₃—NO₂
- A nitrogen atom is bonded to a carbon with a single bond and two oxygens, one with a double bond and one with a single bond (carrying a negative charge).
For each molecule, calculating the formal charge involves assigning electrons from bonds and lone pairs and comparing them to the valence electrons for each atom.](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2F6d43ccb7-7359-44f4-9027-62e63eace4bf%2F3c623001-da55-4d4d-b29e-a4a1200f2046%2Fr2kx5li_processed.jpeg&w=3840&q=75)
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