AP Chemistry Chapter 10 Test FRQ
Ammonia, NH3 , is very soluble in water, whereas phosphine, PH3 , is only moderately soluble in water.
Ammonia has hydrogen-bonding intermolecular forces, whereas phosphine has dipole-dipole and/or dispersion intermolecular forces. Water also has hydrogen-bonding intermolecular attractive forces. Ammonia is more soluble in water than phosphine because ammonia molecules can hydrogen-bond with water molecules, whereas phosphine molecules cannot hydrogen-bond with water molecules.
The melting point of NaF is 993°C, whereas the melting point of CsCl is 645°C
Both NaF and CsCl are ionic compounds with the same charges on the cations and anions. The ionic radius of Na+ is smaller than the ionic radius of Cs+ and the ionic radius of F− is smaller than the ionic radius of Cl−. Therefore, the ionic centers are closer in NaF than in CsCl. Melting occurs when the attraction between the cation and the anion are overcome due to thermal motion. Since the lattice energy is inversely proportional to the distance between the ion centers (Coulomb's Law), the compound with the smaller ions will have the stronger attractions and the higher melting point.
At 25°C and 1 atm, F2 is a gas, whereas I2 is a solid
Both are nonpolar molecules that only expierence dispersion forces, however, the electrons in I2 occupy a larger volume than F2 because I2 has valence electrons in a higher subshell, causing its electron cloud to be larger and more polarizable which increases the strength of the LDFs. A greater number of electrons gives it more of a chance of an instantaneous dipole (polarizability)
Dichloromethane has a greater solubility in water than carbon tetrachloride has. Account for this observation in terms of the intermolecular forces between each of the solutes and water.
Dichloromethane is a polar molecule and therefore interacts with water via dipole-dipole interactions, which are stronger than carbon tetrachloride's weak dipole-induced LDFS that interact with water
Ethanol is completely soluble in water, whereas ethanethiol has limited solubility in water. Account for the difference in solubilities between the two compounds in terms of intermolecular forces.
Ethanol is a polar molecule that experiences LDFs, dipole-dipole, and H-bonding. Thus, the hydrogen bonds increases the attraction between molecules of ethanol and molecules of water, making them more soluble in each other. Ethanethiol is also a polar but only experiences LDFs and dipole-dipole interactions, and is therefore not as polarizable in water which limits its solubility.
Predict which of the two processes shown in the table has the greater change in entropy. Justify your prediction. I2(s) to I2(g) or Br2 (l) to Br2 (g)
I2(s) - I2(g) should have the greater change in entropy. The sublimation of I2 may be thought of as a combination of fusion and vaporization. The conversion from solid to liquid would involve an increase in entropy, as would the conversion from liquid to gas. Br2 is only undergoing the liquid to gas conversion and so will undergo a smaller entropy increase.
In which layer, water or hexane, would the concentration of I3- be higher? Explain.
I3- would be more soluble in water because it experiences LDFs and dipole-dipole forces, thus because of the ion-dipole interactions that would occur between the ions and the polar water molecules it would be more soluble. These interactions are not possible in the nonpolar hexane
) I2(s) and Br2(l) can react to form the compound IBr(l). Predict which would have the greater molar enthalpy of vaporization, IBr(l) or Br2(l). Justify your prediction.
IBr(l). Two reasons may be given. First, IBr is polar, and dipole-dipole forces would tend to increase the enthalpy of vaporization. Second, IBr should have stronger London dispersion forces because of the greater number of electrons in the larger IBr molecule.
Why does hydrogen bonding in a molecule affect its solubility
If a compound has intramolecular h-bonding, it is capable of forming intermolecular hydrogen bonds with H2O, increasing its solubility. This is because hydrogen bonds are extremely polar due to the large electronegativity difference between H and N, O, or F. This causes N, O, or F to hold electrons tighter and hold a partial negative charge, while H has a partial positive charge. This partial positive H atom readily associates with the minus end of another oxygen atom in a water molecule.
SO2 melts at 201 K, whereas SiO2 melts at 1,883 K. Account for the difference in melting points. You must discuss both of the substances in your answer.
In the solid phase, SO2 consists of discrete molecules with dipole-dipole and London (dispersion) forces among the molecules. These forces are relatively weak and are easily overcome at a relatively low temperature, consistent with the low melting point of SO2 . In solid SiO2 , a network of Si and O atoms, linked by strong covalent bonds, exists. These covalent bonds are much stronger than typical intermolecular interactions, so very high temperatures are needed to overcome the covalent bonds in SiO2 . This is consistent with the very high melting point for SiO2 .
MgO melts at a much higher temp than NaF
MgO and NaF are both ionic solids, however, according to Coulomb's Law, the attraction between +2 and -2 ions is greater than +1 ions and -1 ions. Greater charges means greater attraction and thus requires a greater amount of lattice energy to split the intermolecular attraction
The boiling point of liquid propene (226 K) is lower than the boiling point of liquid vinyl chloride (260 K). Account for this difference in terms of the types and strengths of intermolecular forces present in each liquid.
Propene only experiences LDFs while vinyl has LFDs and Dipole-dipole forces, which are stronger and require more energy to overcome. These stronger intermolecular forces cause vinyl to have a larger electron cloud, which increases the polarizability and creates a larger dipole moment (increase in dipole moment increases the attractive forces between the molecules), and thus increases its boiling point
Structures of the pyridine molecule and the benzene molecule are shown below. Pyridine is soluble in water, whereas benzene is not soluble in water. Account for the difference in solubility. You must discuss both of the substances in your answer.
Pyridine is polar (and capable of forming hydrogen bonds with water), while the nonpolar benzene is not capable of forming hydrogen bonds. Pyridine will dissolve in water because of the strong hydrogen bonds (or dipole-dipole intermolecular interactions) that exist between the lone pair of electrons on pyridine's nitrogen atom and the solvent water molecules. No such strong intermolecular interaction can exist between benzene and water, so benzene is insoluble in water.
Account for the fact that Si melts at a much higher temp than Cl2
Si is a network covalent solid with strong covalent bonds between atoms. Cl2 has discrete molecules with weak LDFs between the molecules, thus more energy is required to break the stronger bonds of Si
Compare the strength of the dipole-dipole forces in liquid H2S to the strength of the dipole-dipole forces in liquid H2O. Explain.
Since S is a less electronegative atom than O, the strength of the dipole-dipole forces in liquid H2S are weaker than the dipole-dipole forces in liquid H2O. The net dipole moment of the H2S molecule is less than that of the H2O molecule, which results from the lesser polarity of the H-S bond compared with that of the H-O bond.
Explain why the hexane layer is light purple while the water layer is virtually colorless. Your explanation should reference the relative strengths of interactions between molecules of I2 and the solvents H2O and C6H14 , and the reasons for the differences.
The entrance of the I2 into water requires disruption of the hydrogen bonds in water, which are much stronger than the LDFs in hexane. In hexane, the I2 can dissolve because hexane is nonpolar and I2 is nonpolar, and like dissolves like. Since water is a highly polar molecule and I2 is a nonpolar molecule, they will only interact through a dipole-induced LDFs, but they aren't attractive enough to overcome the sufficient H-bonding bonds between H20 molecules. Meanwhile, the London dispersion forces between I2 and hexane would be stronger than the London dispersion forces between I2 and water.
The normal boiling point of Cl2 (l) (238 K) is higher than the normal boiling point of HCl(l) (188 K). Account for the difference in normal boiling points based on the types of intermolecular forces in the substances. You must discuss both of the substances in your answer.
The intermolecular forces in liquid Cl2 are London (dispersion) forces, whereas the intermolecular forces in liquid HCl consist of London forces and dipole-dipole interactions. Since the boiling point of Cl2 is higher than the boiling point of HCl, the London forces among Cl2 molecules must be greater than the London and dipole-dipole forces among HCl molecules. The greater strength of the London forces between Cl2 molecules occurs because Cl2 has more electrons than HCl, meaning it has a larger electron cloud, increased polarizability, and stronger LDFs (the strength of the London interaction is proportional to the total number of electrons because large atoms with many electrons typically have very diffuse electron clouds and large atomic radii that limit the interaction of their external electrons and the nucleus. The greater the number of electrons, the less control the nuclear charge has on charge distribution, and thus the increased polarizability of the atom. As polarizability increases, the dispersion forces also become stronger. Thus, molecules attract one another more strongly and melting and boiling points of covalent substances increase with larger molecular mass)
Structures of the dimethyl ether molecule and the ethanol molecule are shown below. The normal boiling point of dimethyl ether is 250 K, whereas the normal boiling point of ethanol is 351 K. Account for the difference in boiling points. You must discuss both of the substances in your answer.
The intermolecular forces of attraction among molecules of dimethyl ether consist of London (dispersion) forces and weak dipole-dipole interactions. In addition to London forces and dipole-dipole interactions that are comparable in strength to those in dimethyl ether, ethanol can form hydrogen bonds between the H of one molecule and the O of a nearby ethanol molecule. Hydrogen bonds are particularly strong intermolecular forces, so they require more energy to overcome during the boiling process. As a result, a higher temperature is needed to boil ethanol than is needed to boil dimethyl ether.
Two types of intermolecular forces present in liquid H2S are London (dispersion) forces and dipoledipole forces. (i) Compare the strength of the London (dispersion) forces in liquid H2S to the strength of the London (dispersion) forces in liquid H2O. Explain.
The strength of the London forces in liquid H2S is greater than that of the London forces in liquid H2O. The electron cloud of H2S is larger and has more electrons and is thus more polarizable than the electron cloud of the H2O molecule.
Why does CS2 have a higher boiling point than COS? Describe in terms of intermolecular forces
While CS2 only has LDFs, and COS has LDFs and dipole-dipole interactions, the S atom in CS2 is larger than the O atom in COS, thus CS2 has a greater electron cloud, meaning it is more polarizable, and thus has stronger LDFs that require more energy to overcome than COS's combined LDFs and dipole-dipole forces
Explain why the enthalpy changes of I2(g) from I2(s) is larger than the formation of Br2(g) from Br2(l). In your explanation identify the type of particle interactions involved and a reason for the difference in magnitude of those interactions.
While both non-polar molecules only experience LDFs, I2 has stronger LDFs because it is a larger molecule and has more electrons, thus has a larger electron cloud and a higher polarizability and stronger London dispersion forces. The second reason is that since ΔH of sublimation is approximately ΔH of fusion plus ΔH of vaporization, I2(g) should have a larger ΔH° of formation since it involves sublimation, whereas Br2(g) formation involves only vaporization.
In terms of intermolecular forces, explain why dichloromethane has a higher vapor pressure than carbon tetrachloride.
dichloromethane is a smaller molecule, thus it has a smaller electron cloud and has weaker LDFs than carbon tetrachloride. Carbon tetrachloride is a larger molecule, and therefore has a greater polarizability due to its valence electrons being held less tightly by the nucleus (weaker effective nuclear charge = larger electron cloud) and can more easily form temporary dipoles. This causes it to have stronger LDFs that require more energy to overcome than dichloromethane's LDFs and dipole-dipole forces.
isomer
molecules or polyatomic ions with identical molecular formula - that is, same number of atoms of each element - but distinct arrangements of atoms in space