Which of the following incorrectly shows the bond polarity?...

Which of the following incorrectly shows the bond polarity? Show the correct bond polarity for those that are incorrect.
a. $^{2}+H-F^{2}$
b. $^{3+} \mathrm{Cl}-1^{5-}$
c. $^{3+} S i-S^{3-}$
d. $^{3+} B r-B r^{3-}$
e. $^{3+} \mathrm{O}-\mathrm{P}^{3-}$

Key Concepts

Bond Polarity

Bond polarity refers to the uneven distribution of electron density between two atoms in a covalent bond. When the electronegativity of the atoms differs, the shared electrons are drawn more toward the more electronegative atom, creating a dipole with partial positive (?+) and partial negative (??) charges. This concept is fundamental for understanding molecular interactions, reactivity, and properties such as solubility and boiling point.

Electronegativity Differences

Electronegativity is the tendency of an atom to attract electrons in a chemical bond. The difference in electronegativity between two bonded atoms determines the magnitude of the bond's polarity. A larger difference leads to a more polarized bond, whereas similar electronegativities result in a more nonpolar or evenly shared electron cloud.

Oxidation States Versus Partial Charges

Oxidation states (or oxidation numbers) are a formalism used to keep track of electrons in redox reactions and provide an indication of electron loss or gain in a compound. While they are integers assigned based on standard rules, they differ from partial charges, which are actual electron density shifts in a bond due to differences in electronegativity. Confusing oxidation states with partial charges can lead to misrepresentation of the true electron distribution in a molecule.

Dipole Moment Representation

The dipole moment is a quantitative measure of a bond's polarity, calculated as the product of the magnitude of the partial charge and the distance between the charges. Representing dipole moments accurately is crucial for predicting molecular properties and interactions. In diagrams, the signs (?+ and ??) indicate the direction of electron pull, emphasizing that the more electronegative atom carries the partial negative charge.

Transcript

00:01 All right, in this question, we are given a number of bonds and we are told the polarities and we're asked to determine if these are likely correct or incorrect.

00:15 And so just to orient you, these plus signs indicate a partially positive charge and these minus signs indicate a partially negative charge.

00:23 So this is a electron negativity question.

00:27 Just to recap as we go across the periodic table, electronegativity increases, and as we go down the periodic table, electro negativity decreases.

00:39 So hydrogen fluorine, we know that fluorine is the most electronegative element on the periodic table, so it would make sense that it's partially negative and the hydrogen is partially positive because fluorine is just going to pull the electrons more than the hydrogen.

00:57 Next, we have chlorine versus iodine.

00:59 So chlorine is here, iodine is here.

01:02 Chlorine should have a greater electronegativity because it is higher up on the periodic table and in the same column.

01:10 So in this case, we would expect chlorine to have the negative charge instead of the iodine because chlorine is more electronegative.

01:20 Next, we're comparing silicon versus sulfur.

01:24 So sulfur is over here.

01:26 Silicon is over here.

01:28 Sulfur is to the right of the periodic table of silicon while being on the same row.

01:33 So we would expect sulfur to have the greater electronegativity, and therefore sulfur should have a negative, partially negative charge...

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