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Definition
particles are spread out
have no set shape or volume |
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Definition
particles are close together
have a constant volume, but no set shape |
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Definition
particles are close together and arranged in a crystal lattice
have a constant volume and shape |
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| a measure of the amount of energy needed to expand the surface area of a liquid |
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Definition
| surrounded by molecules to which it is attracted--no net attraction |
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Definition
| feels a net attraction toward the interior |
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Definition
| the number of particles changing from liquid to vapor will eventually equal the number of particles changing from vapor to liquid |
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| depends on the IMF and the KE of the liquid, not on the volume of the container |
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| equilibrium vapor pressure (VP) |
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Definition
| the pressure exerted by vapor of the liquid |
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Term
| What is the difference between intramolecular forces and intermolecular forces? |
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Definition
| Intramolecular forces are those within the molecule, the bonds between atoms, whereas intermolecular forces are those between molecules. |
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Term
| What must be present in a molecule if it is a polar molecule? |
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Definition
| If a molecule is to be polar, it must have a polar-covalent bond, a bond between two atoms of differing electronegativity. |
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Term
| Why is the bond between carbon and oxygen polar in carbon dioxide? If the carbon-oxygen bond is polar, why is the CO2 molecule not polar? |
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Definition
| The electrons shared between carbon and oxygen are drawn closer to oxygen because it has the greater electronegativity (O = 3.5, C = 2.5). Both carbon-oxygen bonds are polar, but CO2 is a linear molecule orienting the two bonds to be opposite each other. This cancels the polarity of each bond, making the molecule nonpolar overall. |
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Term
Which of the following molecules has polar bonds and is a polar molecule?
(a) O3 – bent
(b) HCN – linear
(c) SO3 – trigonal planar
(d) CH3Br – tetrahedral |
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Definition
(a) no polar bonds, nonpolar molecule
(b) polar bonds and polar molecule
(c) polar bonds but nonpolar molecule
(d) polar bonds and polar molecule |
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Term
| What is meant by a change of state? Why is the decomposition of H2O(l) to form two gases, H2(g) and O2(g) not a change of state? |
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Definition
| A change of state is a physical change from one state—solid, liquid, or gas—to another. The decomposition of liquid water to form H2(g) and O2(g) is a chemical change, a decomposition. |
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Term
What is the term used to describe each of the following phase changes?
(a) Br2(l) " Br2(s)
(b) CO2(s) " CO2(g)
(c) H2O(g) " H2O(s)
(d) Hg(s) " Hg(l) |
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Definition
(a) freezing
(b) sublimation
(c) condensation or deposition
(d) melting or fusion
(e) vaporization or evaporation
(f) condensation |
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List the type(s) of intermolecular forces that would be present in the following substances:
(a) Cl2(l)
(b) Cl2CH2(l)
(c) HCF3(l)
(d) Cl-NH2(l)
(e) CO2(s) |
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Definition
(a) London
(b) London, dipole-dipole
(c) London, dipole-dipole
(d) London, dipole-dipole, and hydrogen bonding
(e) London
(f) London |
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Term
| Why is the boiling point of a polar liquid generally higher than the boiling point of a nonpolar liquid of similar molecular mass? |
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Definition
| If the molecular masses of the molecules are similar, then the London forces in each should be similar. But if one molecule can also engage in dipole-dipole attractive forces, it will have the greater total intermolecular attractive forces that will cause it to have the higher boiling point. |
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| Why does water have such a high boiling point? |
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Definition
| The high boiling point of water is because of the extensive hydrogen bonding in liquid water. Each water molecule can engage in a maximum of four hydrogen bonds: two with its own hydrogen atoms, and two between its oxygen atom and hydrogen atoms from two other water molecules. |
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Term
| If you spill nail polish remover on your skin, you sense a cooling sensation as it quickly evaporates. Why? |
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Definition
| As the liquid evaporates, intermolecular forces between molecules at the surface are broken, and this requires energy. That energy can be drawn from your body, which you sense as cooling. Evaporation is a cooling process. |
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Term
| What must occur if intermolecular forces are to have any appreciable effect in a substance? |
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Definition
| The molecules must be very close together or touching. |
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Term
| Under what condition is the vapor pressure of liquids measured? Does the vapor pressure of liquids increase or decrease as temperature increases? |
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Definition
| The vapor pressure of liquids is measured at the point of equilibrium between the liquid and gaseous phases. The vapor pressure of liquids increases with increasing temperature. |
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Term
| The boiling point of NH3(l) is –33°C, and that of H2S(l) is –60°C. Which would have the greater vapor pressure at –70°C? |
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Definition
| The one with the lower boiling point will have the higher vapor pressure at –70°C: H2S. |
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Term
| What does it mean when fusion and vaporization are described as constant temperature processes? |
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Definition
| In fusion and vaporization, all the heat energy goes into breaking down intermolecular forces, not changing the temperature of the substance. |
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Term
| Compare the crystal hardness and melting temperatures of molecular solids and ionic solids. |
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Definition
| Molecular solids have softer crystals with lower melting temperatures than ionic solids. |
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Term
| Why are the melting points of covalent-network solids so very high? |
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Definition
| If a covalent-network solid is to melt to a liquid, many covalent bonds must be simultaneously broken, and that requires a great deal of energy, thus the high temperature. |
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Term
What is the general difference between the following?
(a) a gas and a vapor
(b) vaporization and boiling
(c) a crystalline solid and an amorphous solid |
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Definition
(a) A vapor is the gaseous phase of a substance that is normally a liquid at room temperature (H2O(g), CCl4(g)), while a substance regarded as a gas is normally a gas at room temperature (N2, O2). Both vapors and gases behave as gases.
(b) Vaporization is the transfer of molecules on the surface of a liquid to the gas phase; boiling occurs when vaporization occurs throughout the liquid as it converts to a gas.
(c) In a crystalline solid, the particles are packed in a highly organized way, but in an amorphous solid, they are packed together randomly. |
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Term
| List the following compounds in order of increasing London dispersion force: CCl4, SiF4, CH4, and CBr4. |
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Definition
| They are all tetrahedral, nonpolar molecules and the strength of the London dispersion force increases with molecular mass. The order of increasing London force is: CH4 (16.0 amu) < SiF4 (104.1 amu) < CCl4 (153.8 amu) < CBr4 (331.6 amu). |
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| Arrange the following in order of increasing London forces: He, OCS, H2O, HCl, and SO3. |
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Definition
| The strength of the London dispersion force increases with molecular mass. The order of increasing London force is: He < H2O < HCl < OCS < SO3. |
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| What kind of species exhibit London dispersion forces? |
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Definition
| As long as one electron is in the species, it will exhibit the London dispersion force, and that includes everything except the hydrogen ion, H+. |
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Term
In which of the following is hydrogen bonding an important intermolecular force?
(a) FCH3
(b) H2S
(c) C6H12O6 – glucose
(d) H2O2 |
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Definition
| Hydrogen bonding is important in (c) glucose, a sugar, (d) H2O2, and (e) C2H5OH, ethyl alcohol. |
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Term
The boiling point of liquid chlorine, Cl2(l), is –34°C. Which of the following would be a reasonable boiling point for liquid bromine, Br2(l)?
59°C
–170°C
–50°C |
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Definition
| 59°C is the best choice. The London forces increase with molecular mass, so the boiling point for Br2(l) should be greater than that for Cl2(l). The other temperatures are lower and not expected. |
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Term
The molar heat of vaporization of ethane, C2H6, a nonpolar molecule, is 14.7 kJ/mole. Which of the following values would you expect to be the molar heat of vaporization of propane, C3H8, also a nonpolar molecule? Justify your selection.
8.2 kJ/mole
18.8 kJ/mole
12.7 kJ/mole |
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Definition
| Both ethane and propane are nonpolar, but propane, being a larger molecule, will exhibit larger London dispersion forces. Expect propane to have a higher value, 18.8 kJ/mole. |
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Term
| What is general relationship between the vapor pressure of a liquid and the strength of the intermolecular forces in the liquid? |
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Definition
| The stronger the intermolecular forces in a liquid, the lower its vapor pressure at a given temperature. The stronger the intermolecular forces, the more energy it requires to escape the liquid phase and enter the gas phase. |
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Term
| Arrange these molecules in order of increasing strength of the dipole-dipole force: HI, HBr, HF, HCl. Justify your sequence with an explanation. |
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Definition
| These are all two-atom molecules, so the strength of dipole-dipole force will depend only on the polarity of the bond. The polarity increases with the increasing electronegativity difference between the two bonded atoms. Knowing that the electronegativity of H is 2.1, you can determine the differences in electronegativity are: HI (2.5 – 2.1 = 0.4), HBr (2.8 – 2.1 = 0.7), HCl (3.0 – 2.1 = 0.9), and HF (4.0 – 2.1 = 1.9). The order of increasing strength of dipole-dipole force is: HI < HBr < HCl < HF. |
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| How much heat energy is required to vaporization 0.850 kg of octane, C8H18? The molar heat of vaporization of octane is 41.5 kJ/mole. The molar mass of C8H18 is 114.2 g. |
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Definition
309 kJ of heat is absorbed as 0.850 kg (850 g) of octane is vaporized.
moles C8H18 =(850 g C8H18) X 1mole C8H18/114.2 g C8H18 = 7.44 moles C8H18
amount of heat absorbed = (7.44 moles C8H18) X (41.5 kJ C8H18/1 moled C8H18) = 308 kJ |
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| The molar heat of fusion for water is 6.01 kJ/mole. How much heat energy is needed to melt (fuse) 1802 g (100. moles) of water at 0°C? |
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Definition
601 kJ of heat energy is required.
amount of heat absorbed = (100. mole H2)) X (6.01 kJ/1mole H2O) = 601 kJ |
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Term
| Thinking in terms of the kinetic energy of molecules, why does the vapor pressure of a liquid increase with increasing temperature? |
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Definition
| If a molecule is to escape from a liquid and enter the gas phase, it must overcome the intermolecular forces holding it in the liquid phase. As temperature increases, the average kinetic energy of the molecules in the liquid increases—giving more molecules the energy to overcome the forces holding them in the liquid phase, and allowing them to evaporate, increasing the concentration of molecules in the vapor phase. Thus, vapor pressure increases. |
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Term
| Why is the vapor pressure of ethyl alcohol, C2H5OH(l), higher than that of water, H2O(l), at the same temperature? |
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Definition
| Both liquids, water and ethyl alcohol, can engage in all three intermolecular forces, but water can hydrogen bond more extensively than can the alcohol (water has two –O–H bonds; the alcohol, has one), giving water stronger intermolecular forces to overcome in vaporization. At a given temperature with lower intermolecular forces, it will be easier for the alcohol to escape the liquid phase than water, so it will always have a greater vapor pressure. |
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Term
| Both dimethyl ether and ethyl alcohol have the same formula, C2H6O. The molar heats of vaporization are 21.5 kJ/mole and 38.7 kJ/mole, respectively. Only one of the two can engage in hydrogen bonding. Which one is it? Justify your choice. |
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Definition
| The greater the forces of attraction in the liquid, the greater the heat of vaporization. The greater heat of vaporization for the alcohol indicates that it has the greater forces of attraction; this would be so if ethyl alcohol engaged in hydrogen bonding, which it does. |
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Term
| Explain why the normal boiling point of HF (19.4°C) is higher than the normal boiling point of HBr (–66.7°C), while the boiling point of Br2 (58.8°C) is higher then that of F2 (–188°C). |
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Definition
| The boiling point of HF is greater than that of HBr because HF can engage in strong hydrogen bonding, something HBr cannot do. It's a different story with the diatomic elements. The only intermolecular force in these nonpolar molecules is the London force, which is greater in Br2 (159.8 amu) than in F2 (38.0 amu), causing Br2 to have the higher boiling point. |
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| How many grams of solid benzene, C6H6, can be melted if 1500. kJ of heat energy is absorbed at its melting temperature of 5.5°C? The molar heat of fusion for benzene is 9.87 kJ/mole. The molar mass of benzene is 78.1 g. |
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Definition
1.19 × 104 g of benzene will be melted. First, calculate the number of moles of benzene that can be melted with 1500. kJ of heat energy, and then convert moles of benzene to mass.
moles C6H6 = (1500.kJ) X (1mole C6H6/9.87 kJ) = 152 moles C6H6
mass C6H6 = (152 moles C6H6) X (78.1 g C6H6/1 moles C6H6)= 1.19 X 10^4 g c6H6 |
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| How much heat energy is evolved as 250. g of liquid ammonia freezes to form solid ammonia at its normal freezing point? The molar heat of fusion of ammonia is 5.65 kJ/mole. The molar mass of ammonia is 17.0 g. State the answer in terms of a change in heat energy, DH. |
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Definition
DH = –83.1 kJ
moles NH3 = (250 g NH3) X (1mole NH3/17.0 g NH3) = 14.7 moles NH3
amount of heat evolved = (14.7 moles NH3) X (5.65 kJ/1 mole NH3) = 83.1 kJ |
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| What mass of carbon disulfide, CS2, can be vaporized by absorbing 1.35 × 103 kJ of heat energy at its boiling point? The molar heat of vaporization of CS2 is 27.4 kJ/mole. The molar mass of CS2 is 76.15 g. |
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Definition
3.75 × 103 g of CS2 can be vaporized. First calculate the number of moles of CS2 that can be vaporized with 1.35 × 103 kJ of heat energy, and then convert moles of CS2 to mass.
mole CS2 = (1.35 x 10^3 kJ) X (1.00 mole CS2)/27.4 kJ) = 49.3 moles CS2
mass CS2 - (49.3 moles CS2) X ( 76.15 g CS2/1 mole CS2) = 3754 g = 3.75 X 10^3 g CS2 |
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| How many kilojoules of heat energy are evolved as 5.00 kg of water is frozen to ice at 0°C? The molar heat of fusion of water is 6.01 kJ/mole, and the molar mass of water is 18.02 g. State the answer in terms of a change in heat energy, DH. |
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Definition
277 moles of water evolve 1.67 × 103 kJ of heat energy. DH = – 1.67 × 103 kJ
moles H2) = (5.00 x 10^3 g H2O) X (1 mole H20/18.02 g H20) = 277 moles H20
amount of heat evolved = (277 moles H20) X ( 6.01 kJ / 1 mole H20) = 1.67 X 10^3 kJ |
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Term
What is the strongest intermolecular force present for each of the following compounds?
water
carbon tetrachloride
ammonia
carbon dioxide
phosphorus trichloride
nitrogen
ethane (C2H6)
acetone (CH2O)
methanol (CH3OH)
borane (BH3) |
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Definition
hydrogen bonding
London dispersion forces
hydrogen bonding
London dispersion forces
dipole-dipole forces
London dispersion forces
London dispersion forces
dipole-dipole forces
hydrogen bonding
dipole-dipole forces |
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Term
For each of the following compounds, determine the main intermolecular force. You may find it useful to draw Lewis structures for some of these molecules:
nitrogen
carbon tetrachloride
H2S
sulfur monoxide
N2H2
boron trihydride
CH4O
SiH2O |
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Definition
nitrogen – Van der Waals forces
carbon tetrachloride – Van der Waals forces
H2S – dipole-dipole forces
sulfur monoxide – dipole-dipole forces
N2H2 – hydrogen bonding
boron trihydride – Van der Waals forces
CH4O – hydrogen bonding
SiH2O – dipole-dipole forces |
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Term
| Motor oil largely consists of molecules that consist of long chains of carbon atoms with hydrogen atoms attached to them. Using your knowledge of intermolecular forces, why wouldn’t it be better to use a compound like glycerol. The formula of glycerol is CHOH(CH2OH)2. |
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Definition
| A casual glance at the structure of glycerol shows that it contains at least one hydrogen bond (and in fact, it contains several) – as a result, we would expect it to have high viscosity. Motor oil, on the other hand, contains only Van der Waals forces, which gives it relatively low viscosity. Because the purpose of motor oil is to quickly coat all surfaces of an engine to keep it from getting too hot, a low viscosity material is better than a high viscosity material. |
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Term
| Rank the following by from lowest to highest anticipated boiling point: C2H4, CH4, Ne, H3COCH3. |
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Definition
The highest is clearly H3COCH3, as it’s the only polar molecule. The three lowest all experience Van der Waals forces, meaning that the biggest one will have the highest boiling point.
Overall, the ranking is Ne (-246.10 C) < CH4 (-161.50 C) < C2H2 (-103.70 C) < H3COCH3 (-23.70 C). |
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Term
| Explain why ethyl alcohol (C2H5OH) has a higher boiling point (78.40 C) than methyl alcohol (CH3OH; 64.70 C). |
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Definition
| Both molecules contain one O-H bond, which means that they do the same amount of hydrogen bonding. However, ethyl alcohol is a larger molecule, which means that Van der Waals forces are stronger in it, giving it a slightly higher boiling point. |
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Term
For each of the following compounds indicate which intermolecular force is most important:
a) FCN
b) HCN
c) C2H6
d) CF2H2 |
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Definition
a) FCN dipole-dipole force
b) HCN dipole-dipole force
c) C2H6 Van der Waals forces
d) CF2H2 dipole-dipole force |
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Term
| What causes dipole-dipole interactions? |
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Definition
| Dipole-dipole interactions are caused by the attraction of a partially-positive atom on one compound to the partially-negative atom on another. |
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Term
| Why is hydrogen bonding only possible with hydrogen? |
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Definition
| Hydrogen is the only atom with an unshielded nucleus when it forms covalent bonds. |
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Term
| Which of the following is least strong? |
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Definition
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Term
| What is the effect of a surfactant on surface tension? |
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Definition
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Term
| Ice is less dense than water because: |
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Definition
| The molecules of ice are locked together via hydrogen bonding, and the resulting structure has large empty spaces within it that liquid water doesn’t have. |
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Term
| Chlorine is a gas, bromine is a liquid, and iodine is a solid because: |
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Definition
| They have different melting points because each has a different atomic mass. Generally, heavier molecules have higher melting points. |
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Term
Rank the molecules from lowest to highest polarity:
PF3, LiOH, SF2, NF3 |
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Definition
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Term
Rank the molecules from lowest to highest polarity:
Ni(OH)3, N2H2, CH3OH, C2H5OH |
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Definition
| N2H2 < C2H5OH < CH3OH < Ni(OH)3 |
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Term
Rank the molecules from lowest to highest polarity:
B2F4, H2C2O4, CuCl2, CF2O |
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Definition
| B2F4 < H2C2O4 < CF2O < CuCl2 |
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Term
Rank the molecules from lowest to highest polarity:
PH3, PF3, NH3, NF3 |
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Definition
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Term
Rank the molecules from lowest to highest polarity:
H2O, H2S, HF, H2 |
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Definition
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