6.E: Intermolecular Forces and Liquids and Solids (Exercises)
Intermolecular Forces
In terms of their bulk properties, how do liquids and solids differ? How are they similar?
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Liquids and solids are similar in that they are matter composed of atoms, ions, or molecules. They are incompressible and have similar densities that are both much larger than those of gases. They are different in that liquids have no fixed shape, and solids are rigid.
In terms of the kinetic molecular theory, in what ways are liquids similar to solids? In what ways are liquids different from solids?
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In terms of the kinetic molecular theory, in what ways are liquids similar to gases? In what ways are liquids different from gases?
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They are similar in that the atoms or molecules are free to move from one position to another. They differ in that the particles of a liquid are confined to the shape of the vessel in which they are placed. In contrast, a gas will expand without limit to fill the space into which it is placed.
Explain why liquids assume the shape of any container into which they are poured, whereas solids are rigid and retain their shape.
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What is the evidence that all neutral atoms and molecules exert attractive forces on each other?
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All atoms and molecules will condense into a liquid or solid in which the attractive forces exceed the kinetic energy of the molecules, at sufficiently low temperature.
Open the PhET States of Matter Simulation to answer the following questions:
- Select the Solid, Liquid, Gas tab. Explore by selecting different substances, heating and cooling the systems, and changing the state. What similarities do you notice between the four substances for each phase (solid, liquid, gas)? What differences do you notice?
- For each substance, select each of the states and record the given temperatures. How do the given temperatures for each state correlate with the strengths of their intermolecular attractions? Explain.
- Select the Interaction Potential tab, and use the default neon atoms. Move the Ne atom on the right and observe how the potential energy changes. Select the Total Force button, and move the Ne atom as before. When is the total force on each atom attractive and large enough to matter? Then select the Component Forces button, and move the Ne atom. When do the attractive (van der Waals) and repulsive (electron overlap) forces balance? How does this relate to the potential energy versus the distance between atoms graph? Explain.
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Define the following and give an example of each:
- dispersion force
- dipole-dipole attraction
- hydrogen bond
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- Dispersion forces occur as an atom develops a temporary dipole moment when its electrons are distributed asymmetrically about the nucleus. This structure is more prevalent in large atoms such as argon or radon. A second atom can then be distorted by the appearance of the dipole in the first atom. The electrons of the second atom are attracted toward the positive end of the first atom, which sets up a dipole in the second atom. The net result is rapidly fluctuating, temporary dipoles that attract one another (example: Ar).
- A dipole-dipole attraction is a force that results from an electrostatic attraction of the positive end of one polar molecule for the negative end of another polar molecule (example: ICI molecules attract one another by dipole-dipole interaction).
- Hydrogen bonds form whenever a hydrogen atom is bonded to one of the more electronegative atoms, such as a fluorine, oxygen, nitrogen, or chlorine atom. The electrostatic attraction between the partially positive hydrogen atom in one molecule and the partially negative atom in another molecule gives rise to a strong dipole-dipole interaction called a hydrogen bond (example: HF⋯HFHF⋯HF).
The types of intermolecular forces in a substance are identical whether it is a solid, a liquid, or a gas. Why then does a substance change phase from a gas to a liquid or to a solid?
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Why do the boiling points of the noble gases increase in the order He < Ne < Ar < Kr < Xe?
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The London forces typically increase as the number of electrons increase.
Neon and HF have approximately the same molecular masses.
- Explain why the boiling points of Neon and HF differ.
- Compare the change in the boiling points of Ne, Ar, Kr, and Xe with the change of the boiling points of HF, HCl, HBr, and HI, and explain the difference between the changes with increasing atomic or molecular mass.
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Arrange each of the following sets of compounds in order of increasing boiling point temperature:
- HCl, H 2 O, SiH 4
- F 2 , Cl 2 , Br 2
- CH 4 , C 2 H 6 , C 3 H 8
- O 2 , NO, N 2
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- SiH 4 < HCl < H 2 O;
- F 2 < Cl 2 < Br 2 ;
- CH 4 < C 2 H 6 < C 3 H 8 ;
- N 2 < O 2 < NO
The molecular mass of butanol, C 4 H 9 OH, is 74.14; that of ethylene glycol, CH 2 (OH)CH 2 OH, is 62.08, yet their boiling points are 117 °C and 197 °C, respectively. Explain the reason for the difference.
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On the basis of intermolecular attractions, explain the differences in the boiling points of n –butane (−1 °C) and chloroethane (12 °C), which have similar molar masses.
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Only rather small dipole-dipole interactions from C-H bonds are available to hold n -butane in the liquid state. Chloroethane, however, has rather large dipole interactions because of the Cl-C bond; the interaction is therefore stronger, leading to a higher boiling point.Add texts here.
On the basis of dipole moments and/or hydrogen bonding, explain in a qualitative way the differences in the boiling points of acetone (56.2 °C) and 1-propanol (97.4 °C), which have similar molar masses.
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The melting point of H 2 O( s ) is 0 °C. Would you expect the melting point of H 2 S( s ) to be −85 °C, 0 °C, or 185 °C? Explain your answer.
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−85 °C. Water has stronger hydrogen bonds so it melts at a higher temperature.
Silane (SiH 4 ), phosphine (PH 3 ), and hydrogen sulfide (H 2 S) melt at −185 °C, −133 °C, and −85 °C, respectively. What does this suggest about the polar character and intermolecular attractions of the three compounds?
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Explain why a hydrogen bond between two water molecules is weaker than a hydrogen bond between two hydrogen fluoride molecules.
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The hydrogen bond between two hydrogen fluoride molecules is stronger than that between two water molecules because the electronegativity of F is greater than that of O. Consequently, the partial negative charge on F is greater than that on O. The hydrogen bond between the partially positive H and the larger partially negative F will be stronger than that formed between H and O.
Under certain conditions, molecules of acetic acid, CH 3 COOH, form “dimers,” pairs of acetic acid molecules held together by strong intermolecular attractions:
Draw a dimer of acetic acid, showing how two CH 3 COOH molecules are held together, and stating the type of IMF that is responsible.
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Proteins are chains of amino acids that can form in a variety of arrangements, one of which is a helix. What kind of IMF is responsible for holding the protein strand in this shape? On the protein image, show the locations of the IMFs that hold the protein together:
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H-bonding is the principle IMF holding the DNA strands together. The H-bonding is between the N−HN−H and C=OC=O.
The density of liquid NH 3 is 0.64 g/mL; the density of gaseous NH 3 at STP is 0.0007 g/mL. Explain the difference between the densities of these two phases.
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Identify the intermolecular forces present in the following solids:
- CH 3 CH 2 OH
- CH 3 CH 2 CH 3
- CH 3 CH 2 Cl
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- hydrogen bonding and dispersion forces;
- dispersion forces;
- dipole-dipole attraction and dispersion forces