Explanation : In this example, methane, , has the lowest molecular weight. Example Question 3 : Intermolecular Forces. Explanation : Melting points of hydrocarbons are determined by two main factors: length of the carbon chain and degree of saturation. Example Question 4 : Intermolecular Forces. Explanation : Polarity is determined by differences in electronegativity between the two atoms involved in a bond.
Example Question 5 : Intermolecular Forces. Possible Answers: CCl 4. Correct answer: H 2 O. Explanation : Of the answers, only H 2 O has a net dipole moment, making water the polar molecule.
Example Question 6 : Intermolecular Forces. Possible Answers: Differences in electronegativity between atoms that changes the overall charge of the molecule. Differences in electronegativity between atoms that result in uneven sharing of electrons. Differences in atomic size between atoms that changes the overall charge of the molecule. Differences in atomic size between atoms that result in uneven sharing of electrons. Correct answer: Differences in electronegativity between atoms that result in uneven sharing of electrons.
Explanation : Dipole-dipole interactions are intermolecular forces that result from attraction of partial charges of atoms. Example Question 7 : Intermolecular Forces. Dipole-dipole interactions can be observed in molecules of: I. Methanol II. Dichloromethane III. Possible Answers: II only. Correct answer: I and II. Explanation : Dipole-dipole interactions occur between a partial positive and negative charge.
Example Question 8 : Intermolecular Forces. Which of the following molecules contain intramolecular hydrogen bonds?
Possible Answers: Dimethyl ether. Correct answer: Ortho-nitrophenol. Explanation : The question is asking for intramolecular hydrogen bonds, meaning which of the following molecules will contain hydrogen bonds between the atoms within a single molecule.
Example Question 9 : Intermolecular Forces. Possible Answers: Bond between the hydrogen from ammonia and the oxygen from methanol. Bond between the nitrogen from ammonia and the hydrogen from the hydroxyl group of methanol. Bond between the nitrogen from ammonia and the oxygen from methanol. Correct answer: Bond between the nitrogen from ammonia and the hydrogen from the hydroxyl group of methanol.
Explanation : A hydrogen bond forms between a hydrogen bond donor hydrogen and a hydrogen bond acceptor nitrogen, oxygen, or fluorine. Example Question 10 : Intermolecular Forces. Which of the three molecules cannot participate in hydrogen bonding? Possible Answers: I only. Explanation : Hydrogen bonding is an intermolecular force that occurs between a hydrogen bond donor hydrogen and a hydrogen bond acceptor nitrogen, oxygen, or fluorine.
Copyright Notice. Jonathan Certified Tutor. Alexandra Certified Tutor. Andrew Certified Tutor. Report an issue with this question If you've found an issue with this question, please let us know. Opposite charges attract; like charges repel. As two molecules approach each other, an instantaneous dipole in one molecule will attract opposite charges in the other molecule and create a weak dipole in its neighbor.
The two weak dipoles now attract each other. Although the instantaneous dipole of the first will continue to change, the induced dipole in the second molecule will follow suit, so the weak attraction between the two molecules will persist. This type of intermolecular interaction is called a London dispersion force. Generally, larger molecules are easier to polarize, so they experience stronger London forces than smaller molecules.
London forces are the only intermolecular force that propane molecules experience. Propane molecules are relatively small, so the London forces between them are weak -- too weak to hold them together in solid or liquid phase at room temperature.
They can quickly run up smooth walls and across ceilings that have no toe-holds, and they do this without having suction cups or a sticky substance on their toes. And while a gecko can lift its feet easily as it walks along a surface, if you attempt to pick it up, it sticks to the surface. How are geckos as well as spiders and some other insects able to do this?
The huge numbers of spatulae on its setae provide a gecko, shown in Figure 8. In , Kellar Autumn, who leads a multi-institutional gecko research team, found that geckos adhered equally well to both polar silicon dioxide and nonpolar gallium arsenide. This proved that geckos stick to surfaces because of dispersion forces—weak intermolecular attractions arising from temporary, synchronized charge distributions between adjacent molecules.
By curling and uncurling their toes, geckos can alternate between sticking and unsticking from a surface, and thus easily move across it. Further investigations may eventually lead to the development of better adhesives and other applications. In order for a substance to enter the gas phase, its particles must completely overcome the intermolecular forces holding them together.
Therefore, a comparison of boiling points is essentially equivalent to comparing the strengths of the attractive intermolecular forces exhibited by the individual molecules. For small molecular compounds, London dispersion forces are the weakest intermolecular forces. Dipole-dipole forces are somewhat stronger, and hydrogen bonding is a particularly strong form of dipole-dipole interaction.
However, when the mass of a nonpolar molecule is sufficiently large, its dispersion forces can be stronger than the dipole-dipole forces in a lighter polar molecule. Thus, nonpolar Cl 2 has a higher boiling point than polar HCl. What intermolecular forces besides dispersion forces, if any, exist in each substance? Are any of these substances solids at room temperature? Covalent network compounds contain atoms that are covalently bonded to other individual atoms in a giant 3-dimensional network.
Covalent molecular compounds contain individual molecules that are attracted to one another through dispersion, dipole-dipole or hydrogen bonding. List the three common phases in the order you are likely to find them—from lowest temperature to highest temperature. List these intermolecular interactions from weakest to strongest: London forces, hydrogen bonding, and ionic interactions.
List these intermolecular interactions from weakest to strongest: covalent network bonding, dipole-dipole interactions, and dispersion forces. Explain how a molecule like carbon dioxide CO 2 can have polar covalent bonds but be nonpolar overall. Sulfur dioxide SO 2 has a formula similar to that of carbon dioxide see Exercise 7 but is a polar molecule overall. What can you conclude about the shape of the SO 2 molecule? London forces, hydrogen bonding, and ionic interactions.
The two covalent bonds are oriented in such a way that their dipoles cancel out. Learning Objectives Define phase. Identify the types of interactions between molecules. Covalent Network Materials Substances with the highest melting and boiling points have covalent network bonding. Diamond, a form of pure carbon, has covalent network bonding. Ionic Compounds The strongest force between any two particles is the ionic bond , in which two ions of opposing charge are attracted to each other.
Solid NaCl is held together by ionic interactions. Covalent Molecular Compounds There are two different covalent structures: molecular and network. Dipole-dipole Intermolecular Forces As discussed in Section 4. The electrons in the HF molecule are not equally shared by the two atoms in the bond. Because the fluorine atom has nine protons in its nucleus, it attracts the negatively charged electrons in the bond more than the hydrogen atom does with its one proton in its nucleus.
Thus, electrons are more strongly attracted to the fluorine atom, leading to an imbalance in the electron distribution between the atoms. Such a bond is called a polar covalent bond. Although the individual bonds in both CO 2 and CCl 4 are polar, their effects cancel out because of the spatial orientation of the bonds in each molecule. As a result, such molecules experience little or no dipole-dipole interaction.
Solution CO and N 2 are both diatomic molecules with masses of about 28 amu, so they experience similar London dispersion forces. Hydrogen Bonding Intermolecular Forces Molecules with hydrogen atoms bonded to electronegative atoms such as O, N, and F tend to exhibit unusually strong intermolecular interactions due to a particularly strong type of dipole-dipole attraction called hydrogen bonding.
Answer The melting point and boiling point for methylamine are predicted to be significantly greater than those of ethane. London Dispersion Forces Finally, there are forces between all molecules that are caused by electrons being in different places in a molecule at any one time, which sets up a temporary separation of charge that disappears almost as soon as it appears.
Solution Applying the skills acquired in the chapter on chemical bonding and molecular geometry, all of these compounds are predicted to be nonpolar, so they may experience only dispersion forces: the smaller the molecule, the less polarizable and the weaker the dispersion forces; the larger the molecule, the larger the dispersion forces.
Applications: Geckos and Intermolecular Forces Geckos have an amazing ability to adhere to most surfaces. Boiling Points and Bonding Types In order for a substance to enter the gas phase, its particles must completely overcome the intermolecular forces holding them together. Because a hydrogen atom is so small, these dipoles can also approach one another more closely than most other dipoles. A hydrogen bond is usually indicated by a dotted line between the hydrogen atom attached to O, N, or F the hydrogen bond donor and the atom that has the lone pair of electrons the hydrogen bond acceptor.
Because each water molecule contains two hydrogen atoms and two lone pairs, a tetrahedral arrangement maximizes the number of hydrogen bonds that can be formed. In the structure of ice, each oxygen atom is surrounded by a distorted tetrahedron of hydrogen atoms that form bridges to the oxygen atoms of adjacent water molecules.
The bridging hydrogen atoms are not equidistant from the two oxygen atoms they connect, however. Instead, each hydrogen atom is pm from one oxygen and pm from the other. The resulting open, cagelike structure of ice means that the solid is actually slightly less dense than the liquid, which explains why ice floats on water rather than sinks.
Each water molecule accepts two hydrogen bonds from two other water molecules and donates two hydrogen atoms to form hydrogen bonds with two more water molecules, producing an open, cagelike structure.
The structure of liquid water is very similar, but in the liquid, the hydrogen bonds are continually broken and formed because of rapid molecular motion. Because ice is less dense than liquid water, rivers, lakes, and oceans freeze from the top down. In fact, the ice forms a protective surface layer that insulates the rest of the water, allowing fish and other organisms to survive in the lower levels of a frozen lake or sea.
If ice were denser than the liquid, the ice formed at the surface in cold weather would sink as fast as it formed. Bodies of water would freeze from the bottom up, which would be lethal for most aquatic creatures. Draw the hydrogen-bonded structures. A Of the species listed, xenon Xe , ethane C 2 H 6 , and trimethylamine [ CH 3 3 N] do not contain a hydrogen atom attached to O, N, or F; hence they cannot act as hydrogen bond donors.
B The one compound that can act as a hydrogen bond donor, methanol CH 3 OH , contains both a hydrogen atom attached to O making it a hydrogen bond donor and two lone pairs of electrons on O making it a hydrogen bond acceptor ; methanol can thus form hydrogen bonds by acting as either a hydrogen bond donor or a hydrogen bond acceptor. The hydrogen-bonded structure of methanol is as follows:. Compounds such as HF can form only two hydrogen bonds at a time as can, on average, pure liquid NH 3.
Consequently, even though their molecular masses are similar to that of water, their boiling points are significantly lower than the boiling point of water, which forms four hydrogen bonds at a time. Identify the intermolecular forces in each compound and then arrange the compounds according to the strength of those forces. Electrostatic interactions are strongest for an ionic compound, so we expect NaCl to have the highest boiling point. To predict the relative boiling points of the other compounds, we must consider their polarity for dipole—dipole interactions , their ability to form hydrogen bonds, and their molar mass for London dispersion forces.
Helium is nonpolar and by far the lightest, so it should have the lowest boiling point. Consequently, N 2 O should have a higher boiling point. Because the boiling points of nonpolar substances increase rapidly with molecular mass, C 60 should boil at a higher temperature than the other nonionic substances. The predicted order is thus as follows, with actual boiling points in parentheses:. Although C—H bonds are polar, they are only minimally polar.
The most significant intermolecular force for this substance would be dispersion forces. Although this molecule does not experience hydrogen bonding, the Lewis electron dot diagram and VSEPR indicate that it is bent, so it has a permanent dipole. The most significant force in this substance is dipole-dipole interaction.
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