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Which of the following incorrectly shows the bond polarity? Show the correct bond polarity for those that are incorrect. a. \(^{\delta+} \mathrm{H}-\mathrm{F}^{\delta-} \quad\) d. \(\delta^{+} \mathrm{Br}-\mathrm{Br}^{\delta-}\) b. \(^{\delta+} \mathrm{Cl}-\mathrm{I}^{\delta-} \qquad\) e. \(\quad\) e. \(\quad ^{\delta+}\mathrm{O}-\mathrm{P}^{\delta-}\) c. \(\quad \delta+\mathrm{Si}-\mathrm{S}^{\delta-}\)

Short Answer

Expert verified
The incorrect bond polarities are: - d. \(\delta^{+} \mathrm{Br}-\mathrm{Br}^{\delta-}\), which should be \(\mathrm{Br}-\mathrm{Br}\) - e. \(^{\delta+}\mathrm{O}-\mathrm{P}^{\delta-}\), which should be \(\delta^- \mathrm{O}-\mathrm{P} \delta^+\)

Step by step solution

01

Locate electronegativity values for the involved elements

Use a periodic table with electronegativity values or memorize some common electronegativities. The electronegativity values we need are: - H: 2.2 - F: 3.98 - Cl: 3.16 - I: 2.66 - Si: 1.9 - S: 2.58 - Br: 2.96 - O: 3.44 - P: 2.19 Note that these values are from the Pauling scale and can vary slightly depending on the reference used.
02

Determine the partial charge distribution for each bond

For each molecule, compare the electronegativity values of the atoms involved in the bond: a. H-F: 2.2 - 3.98 --> F is more electronegative, so the bond polarity is correct: \(\delta^+ \mathrm{H}-\mathrm{F} \delta^-\) b. Cl-I: 3.16 - 2.66 --> Cl is more electronegative, so the bond polarity is correct: \(\delta^+ \mathrm{Cl}-\mathrm{I} \delta^-\) c. Si-S: 1.9 - 2.58 --> S is more electronegative, so the bond polarity is correct: \(\delta^+ \mathrm{Si}-\mathrm{S} \delta^-\) d. Br-Br: 2.96 - 2.96 --> Same atoms, so no polarity: \(\mathrm{Br}-\mathrm{Br}\) e. O-P: 3.44 - 2.19 --> O is more electronegative, so the bond polarity is incorrect. The correct bond polarity is: \(\delta^- \mathrm{O}-\mathrm{P} \delta^+\)
03

Identify the incorrect bond polarities and state the correct ones

The incorrect bond polarities are: - d. \(\delta^{+} \mathrm{Br}-\mathrm{Br}^{\delta-}\), which should be \(\mathrm{Br}-\mathrm{Br}\) - e. \(^{\delta+}\mathrm{O}-\mathrm{P}^{\delta-}\), which should be \(\delta^- \mathrm{O}-\mathrm{P} \delta^+\)

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Electronegativity
Electronegativity is a key concept in understanding chemical bonds. It describes an atom's ability to attract and hold on to electrons within a bond. Each element has an electronegativity value. These values help predict how electrons are shared in bonds. Atoms with high electronegativity, like fluorine, tend to pull electrons closer. This leads to the formation of a partial negative charge around them. Conversely, atoms with lower electronegativity, like hydrogen, end up with a partial positive charge. This difference is crucial for understanding bond polarity. Electronegativity values are instrumental in determining bond characteristics, such as whether a bond is more ionic or covalent. Bonds between atoms with significantly different electronegativities often show stronger dipoles, while differences of less than 0.5 generally indicate nonpolar bonds.
Pauling Scale
The Pauling scale is a commonly used method to quantify electronegativity. Developed by Linus Pauling, it assigns values to elements based on their electron-attracting power. These values typically range from about 0.7 for the least electronegative element—like cesium—to 3.98 for fluorine, the most electronegative. Pauling's work allows chemists to compare elements effectively when predicting bond properties. Understanding the Pauling scale helps reveal why certain molecules behave differently in chemical reactions. The scale provides insight into the distribution of electrons in a compound, guiding predictions about molecule shape, reactivity, and interactions. While other scales also exist, Pauling's remains the most widely used due to its comprehensive nature.
Partial Charge Distribution
Partial charge distribution occurs in molecules when atoms have different electronegativities. Electrons tend to gravitate towards the more electronegative atom, creating a dipole—a separation of partial charges. In a simple diatomic molecule, the atom with higher electronegativity becomes partially negative (δ⁻), while the less electronegative atom turns partially positive (δ⁺). This concept is essential in understanding the behavior of polar molecules, as it affects how they interact with other compounds, including solvents and reagents in reactions. It’s important to note that the distinction between the δ⁺ and δ⁻ is subtle yet critical in chemistry. This charge imbalance can influence boiling points, solubility, and biological activity of molecules. Identifying the correct partial charge distribution enables students to predict and explain a molecule's properties accurately.

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Most popular questions from this chapter

The most common type of exception to the octet rule are compounds or ions with central atoms having more than eight electrons around them. PF_. \(\mathrm{PF}_{5}, \mathrm{SF}_{4}, \mathrm{ClF}_{3}\) and \(\mathrm{Br}_{3}^{-}\) are examples of this type of exception. Draw the Lewis structure for these compounds or ions. Which elements, when they have to, can have more than eight electrons around them? How is this rationalized?

An ionic compound made from the metal \(M\) and the diatomic gas \(X_{2}\) has the formula \(M_{a} X_{b},\) in which \(a=1\) or 2 and \(b=1\) or \(2 .\) Use the data provided to determine the most likely values for \(a\) and \(b,\) along with the most likely charges for each of the ions in the ionic compound. Data (in units of \(\mathrm{kJ} / \mathrm{mol} )\) Successive ionization energies of \(\mathrm{M} : 480 ., 4750\) . Successive electron affinity values for \(\mathrm{X} :-175,920\) . Enthalpy of sublimation for \(\mathrm{M}(s) \rightarrow \mathrm{M}(g) : 110\) . Lattice energy for MX \(\left(\mathrm{M}^{+} \text { and } \mathrm{X}^{-}\right) :-1200\) . Lattice energy for \(\mathrm{MX}_{2}\left(\mathrm{M}^{2+} \text { and } \mathrm{X}^{-}\right) :-3500\) . Lattice energy for \(\mathrm{M}_{2} \mathrm{X}\left(\mathrm{M}^{+} \text { and } \mathrm{X}^{2-}\right) :-3600\) . Lattice energy for \(\mathrm{MX}\left(\mathrm{M}^{2+} \text { and } \mathrm{X}^{2-}\right) :-4800\) .

List all the possible bonds that can occur between the elements P, Cs, O, and H. Predict the type of bond (ionic, covalent, or polar covalent) one would expect to form for each bond.

When comparing the size of different ions, the general radii trend discussed in Chapter 7 is usually not very useful. What do you concentrate on when comparing sizes of ions to each other or when comparing the size of an ion to its neutral atom?

Write Lewis structures for the following. Show all resonance structures where applicable. a. \(\mathrm{NO}_{2}^{-}, \mathrm{NO}_{3}^{-}, \mathrm{N}_{2} \mathrm{O}_{4}\left(\mathrm{N}_{2} \mathrm{O}_{4} \text { exists as } \mathrm{O}_{2} \mathrm{N}-\mathrm{NO}_{2} .\right)\) b. \(\mathrm{OCN}^{-}, \mathrm{SCN}^{-}, \mathrm{N}_{3}^{-}\) (Carbon is the central atom in \(\mathrm{OCN}^{-}\) and \(\mathrm{SCN}^{-} . )\)

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