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Chapter 8: Problem 41

Which of the following molecules is(are) polar? For each polar molecule, indicate the direction of polarity-that is, which is the negative end, and which is the positive end of the molecule. (a) \(\mathrm{BeCl}_{2}\) (b) \(\mathrm{HBF}_{2}\) (c) \(\mathrm{CH}_{3} \mathrm{Cl}\) (d) \(\mathrm{SO}_{3}\)

Short Answer

Expert verified
\(\mathrm{HBF}_{2}\) and \(\mathrm{CH}_{3}\mathrm{Cl}\) are polar; negative ends are at \(\mathrm{F}\) and \(\mathrm{Cl}\), respectively.

Step by step solution

01

Determine Molecular Geometry

To assess the polarity of a molecule, we start by determining its molecular geometry, which can help us understand the distribution of electron density. For each molecule: - **(a) \(\mathrm{BeCl}_{2}\)**: Linear shape. - **(b) \(\mathrm{HBF}_{2}\)**: Trigonal planar shape. - **(c) \(\mathrm{CH}_{3}\text{Cl}\)**: Tetrahedral shape. - **(d) \(\mathrm{SO}_{3}\)**: Trigonal planar shape.
02

Assess Electronegativity Differences

Examine the electronegativity of atoms within each molecule to predict bond dipoles. - **Be-Cl bonds** in \(\mathrm{BeCl}_{2}\) have a small difference, making individual bonds slightly polar. - **H-B, B-F bonds** in \(\mathrm{HBF}_{2}\) have differences; the B-F bonds are more polar due to the larger electronegativity difference. - **C-H and C-Cl bonds** in \(\mathrm{CH}_{3} \mathrm{Cl}\); the C-Cl bond is significantly polar. - **S-O bonds** in \(\mathrm{SO}_{3}\) exhibit polarity, but are arranged symmetrically.
03

Determine Net Dipole Moment

To ascertain if molecules are polar or nonpolar, consider both bond polarities and geometry to identify the net dipole moment: - **(a) \(\mathrm{BeCl}_{2}\)**: Despite polar bonds, the linear shape results in dipoles canceling each other, making it nonpolar. - **(b) \(\mathrm{HBF}_{2}\)**: Asymmetric shape and bond dipoles result in a net dipole moment, making it polar, with the negative end at fluorine atoms. - **(c) \(\mathrm{CH}_{3}\mathrm{Cl}\)**: The tetrahedral shape with C-Cl bond dipole results in a net dipole moment, making it polar; the negative end is at the chlorine. - **(d) \(\mathrm{SO}_{3}\)**: Despite polar bonds, the symmetry of the molecule results in dipoles canceling, making it nonpolar.

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Key Concepts

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

Molecular Geometry
Molecular geometry is the three-dimensional arrangement of atoms within a molecule. It's a crucial concept for understanding molecular polarity, as the spatial configuration affects how electron density is distributed. For example:
  • Linear Geometry: Molecules like \(\mathrm{BeCl}_{2}\) have atoms aligned in a straight line. This symmetry often results in nonpolar molecules, as any dipole moments cancel each other out.
  • Trigonal Planar Geometry: Seen in molecules like \(\mathrm{SO}_{3}\), this geometry allows for some symmetries. It means even if individual bonds are polar, the overall molecule may be nonpolar if the vectors cancel each other.
  • Tetrahedral Geometry: Found in \(\mathrm{CH}_{3}\text{Cl}\), where atoms are arranged symmetrically around a central atom. Here, geometry does not cancel out all dipole moments, contributing to polarity.
Understanding geometry helps predict whether a molecule will be polar, as even minor asymmetries can lead to a net dipole moment. This knowledge is foundational when diagnosing the chemical properties of compounds.
Electronegativity
Electronegativity refers to an atom's ability to attract and hold onto electrons in a chemical bond. This property is essential when discussing molecular polarity. Here's why electronegativity matters:
  • Electronegativity Differences: Atoms with higher electronegativity will pull electron density towards themselves, creating a polar bond. \(\ \mathrm{Cl}\) in \(\mathrm{CH}_{3}\mathrm{Cl} \) makes the chlorine atom partially negative, highlighting the bond's polar nature.
  • Bond Polarity: In \(\mathrm{HBF}_{2}\), the \(\mathrm{B-F}\) bonds are polar because of substantial electronegativity differences. The same element, \(\mathrm{F}\), in the molecule contributes to an overall polar molecule due to its strong electronegativity.
Essentially, greater differences in electronegativity between bonded atoms increase the chances of a molecule being polar. Without understanding this concept, predicting molecular behavior becomes challenging.
Dipole Moment
A dipole moment is a vector quantity that results from asymmetric charge distribution in a molecule, indicating polarity. It's quantified by the product of charge separation and the distance between charges. Consider these points:
  • Net Dipole Moment: In molecules like \(\mathrm{CH}_{3}\text{Cl}\), where asymmetrical distribution of polar bonds (like \(\mathrm{C-Cl}\)) leads to a net dipole moment, resulting in net polarity.
  • Zero Dipole Moment: Molecules such as \(\mathrm{SO}_{3}\) and \(\mathrm{BeCl}_{2}\) display zero net dipole moment due to symmetrical geometry that cancels out any individual bond dipoles.
Understanding dipole moments helps in predicting how molecules will interact in fields such as solvents or in reactions requiring polar or nonpolar conditions. It is a key concept at the intersection of physical chemistry and molecular physics.
Covalent Bonds
Covalent bonds are formed when atoms share electrons. These can be either nonpolar or polar, depending on the sharing equality of the electrons. Key aspects of covalent bonds include:
  • Polar Covalent Bonds: These occur when two different atoms share electrons unevenly. The \(\mathrm{C-Cl}\) bond in \(\mathrm{CH}_{3}\mathrm{Cl}\) is a classic example, where electrons are drawn more towards chlorine making it polar.
  • Nonpolar Covalent Bonds: Occur in symmetric molecules such as \(\mathrm{BeCl}_{2}\) where atoms share electrons equally, and any polar nature in bonds is negated by the molecule's symmetry.
Covalent bonds serve as the basis for molecular structure, especially when considering how electrons are distributed amongst atoms. Understanding these interactions is critical to predicting molecular shape, geometry, and ultimately polarity. This forms a foundation for much of chemistry, explaining why molecules behave the way they do in different environments.

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Most popular questions from this chapter

Draw the resonance structures for the formate ion, \(\mathrm{HCO}_{2}^{-},\) and find the formal charge on each atom. If an \(\mathrm{H}^{+}\) ion is attached to \(\mathrm{HCO}_{2}^{-}\) (to form formic acid), does it attach to \(\mathrm{C}\) or \(\mathrm{O} ?\)

The following molecules or ions all have three oxygen atoms attached to a central atom. Draw a Lewis structure for each one, and then describe the electron-pair geometry and molecular geometry around the central atom. Comment on similarities and differences in the series. (a) \(\mathrm{CO}_{3}^{2-}\) (b) \(\mathrm{NO}_{3}^{-}\) (c) \(\mathrm{SO}_{3}^{2-}\) (d) \(\mathrm{ClO}_{3}^{-}\)

Consider a series of molecules in which carbon is attached by single covalent bonds to atoms of second-period elements: \(\mathrm{C}-\mathrm{O}, \mathrm{C}-\mathrm{F}, \mathrm{C}-\mathrm{N}, \mathrm{C}-\mathrm{C}\) and \(\mathrm{C}-\mathrm{B}\). Place these bonds in order of increasing bond length.

The cyanate ion, OCN \(^{-}\), has the least electronegative atom, \(\mathrm{C},\) in the center. The unstable fulminate ion, \(\mathrm{CNO}^{-},\) has the same molecular formula, but the \(N\) atom is in the center. (a) Draw the three possible resonance structures of \(\mathrm{CNO}^{-}\) (b) On the basis of formal charges, decide on the resonance structure with the most reasonable distribution of charge. (c) Mercury fulminate is so unstable it is used in blasting caps for dynamite. Can you offer an explanation for this instability? (Hint: Are the formal charges in any resonance structure reasonable in view of the relative electronegativities of the atoms?)

You are doing an experiment in the laboratory and want to prepare a solution in a polar solvent. Which solvent would you choose, methanol \(\left(\mathrm{CH}_{3} \mathrm{OH}\right)\) or toluene \(\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{CH}_{3}\right) ?\) Explain your choice. EQUATION CANT COPY
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