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Chemistry Chapter 10
Intermolecular Forces
| Question | Answer |
|---|---|
| Significance of Structure (10.1) | The physical and chemical properties of a compound depend on its structure. |
| Intramolecular Forces (10.1) | Interactions within a molecule (as in chapters eight and nine). The prefix "intra" means on the inside, within. |
| Intermolecular Forces (10.1) | Interactions between molecules, or between molecules and ions. These forces are electrostatic in nature, but weaker than chemical bonds. Influences properties like physical state. |
| Solids, Liquids, & Gases (10.1) | Solids retain the same nearest neighbors over time, their kinetic energy is principally vibrational. Liquids have the ability to slide past one another. Gases typically have enough kinetic energy to overcome all intermolecular forces of attraction, as a result they possess a much fuller range of motion. Dispersion can determine which state diatomic molecules exist in at a certain temperature. |
| Physical Change Temperatures (10.1) | The stronger the attractive forces among the particles in a solid, the greater the amount of energy needed to overcome those forces to cause melting or sublimation. A substance made of strongly interacting particles has higher melting/boiling points. Weak intermolecular attractions = lower melting/boiling points. |
| Boiling Point Periodic Trend (10.2) | Boiling points increase as the atomic numbers increase. The stronger the particles’ attractions for each other, the greater the amount of energy needed to separate them. |
| Fritz London's Theory (10.2) | When atoms approach each other they have an electrostatic interaction that is similar to the formation of a covalent bond. (Their respective electron clouds still repel one another). |
| Induced Dipoles (10.2) | Regions of partial, oppositely coordinating electrical charge that attract each other. Induced (temporary) dipole moments are created when an external electric field distorts the electron cloud of a neutral molecule. When the electrons that constantly orbit the nucleus occupy a similar location by chance. Temporary dipoles can induce a dipole in neighboring molecules, initiating an attraction called a London dispersion force (this can occur in nonpolar molecules). |
| Dispersion Forces, aka London Forces (10.2) | All atoms and molecules experience London Forces to some degree. As # of electrons and protons increases, so too does the strength of interactions. This is because the larger the cloud of electrons around an atom, the more likely it is that they will be polarized or unevenly distributed. Greater polarizability means they are more likely to form temporary dipoles that attract molecules to each other in the liquid phase and inhibit vaporization. |
| Why are larger electron clouds more easily polarized? (10.2) | Electrons in larger atoms are held less tightly by the nucleus because of their greater average distance from the nucleus as well as the screening of the nuclear charge by electrons in lower-energy orbitals. |
| Surface Area (10.2) | Large surface area = more opportunity for dispersion forces to act upon molecules. Less surface area (spherical instead of cylindrical) = less dispersion forces and therefore a lower boiling point. For a set of hydrocarbons with no permanent dipoles and the same chemical formula, more branches = less surface area. |
| Ion-Dipole Interactions (10.3) | An attractive force between an ion and a molecule with a permanent dipole moment. When a salt dissolves in water, these Ion-Dipole interactions cause dissociation of the ionic compound. |
| Sphere of Hydration (10.3) | As an ion is pulled out of its lattice and away from its solid-state neighbors, it becomes surrounded by water molecules. (In a solution where the solvent is H₂O). These ions are said to be "hydrated" when this occurs. Usually, this takes six H₂O, but it can range from 4 to 9. The water molecules closest to the ion are oriented so that their oxygen atoms (negative poles) are directed toward a cation or their hydrogen atoms (positive poles) are directed toward an anion. |
| Sphere of Solvation (10.3) | When an ion is surrounded by non-water molecules after being separated from its lattice, this type of sphere is formed. These ions are said to be "solvated" when this occurs. |
| Dipole-Dipole Interactions (10.3) | Experienced by all polar molecules. Partial positive and negative charges on polar molecules are attracted to one another. This interaction occurs often between H₂O, which are bent linear with a steric # of four. The oxygen's lone pairs are more electronegative than the bonded hydrogens, giving H₂O a permanent (1.855 D) dipole moment. It is mutually attracted to up to two other water molecules' hydrogens. |
| Dipole-Dipole vs. Ion-Dipole (10.3) | Dipole-Dipole attractions are NOT as strong as Ion-Dipole, because they only involve partial charges. Ions have exchanged an electron, and thus have one full positive or negative charge at minimum. (Inner hydration spheres are stronger than outer spheres). |
| 2-Methylpropane and Acetone Comparison (10.3) | Both of these molecules have a molar mass of 58 g/mol. They have similar shapes, aside from a different central carbon hybridization. However, the boiling point of acetone is ~70K higher. This is because the C═O bond in acetone creates a large dipole moment (2.88 D) that nonpolar 2-Methylpropane does not have. Therefore, polar acetone has a higher boiling point because of its strong Dipole-Dipole interactions. |
| Dipole-Induced Dipole Interaction (10.3) | Created by a molecule with a permanent dipole when it induces a temporary dipole moment in a nonpolar molecule (by perturbing electron distribution). This is weaker than the Dipole–Dipole force between two polar molecules and is of the same order of magnitude as the dispersion forces between temporary dipoles in nonpolar molecules. |
| Molar Mass Trend (10.3) | Boiling points increase with increasing molar masses. However, H-bonded [NH₃], [H₂O], and [HF] are unusually high compared to their respective series' data. |
| Unusual Strength of H-Bonds (10.3) | Hydrogen atoms share just two electrons, and when they are shared with a HIGHLY electronegative atom of N, O, or F, the electron density on the surface of the hydrogen atom is greatly reduced, leaving the nucleus of the hydrogen atom and its positive charge with very little electron density separating it from an approaching N, O, and F atom on a neighboring polar molecule. This allows the unshielded H atoms to get very close to its H-bonding partner, increasing the strength of the bond. |
| Oligomer (Extra) | A polymer whose molecules consist of relatively few repeating units. It could (theoretically or actually) be derived from copies of a smaller molecule, (its monomer). |
| Dimer (10.3) | An oligomer consisting of two (di-) monomers joined by bonds that can be either strong or weak, covalent or intermolecular. The term homodimer is used when the two molecules are identical (e.g. A–A) and heterodimer when they are not (e.g. A–B). |
| Strongest Dipole-Dipole Interactions (10.3) | Hydrogen bonds. They can be nearly 1/10th the strength of covalent bonds. This is a defining property of water, allowing it absorb a lot of heat. The breaking of hydrogen bonds absorbs heat. Water regulates the thermal energy of its surroundings, which is why the temperature change between seasons is gradual rather than sudden. |
| Continuum of Intermolecular Force Strength (10.3) | Dispersion < Dipole-Induced Dipole < Ion-Induced Dipole < Dipole-Dipole < Hydrogen Bonding < Ion-Dipole |
| Biological Structures & Hydrogen Bonding (10.3) | Within proteins, long chains of atoms tend to fold back and wrap around themselves forming hydrogen bonds with adjacent chains. This creates secondary, tertiary, and quaternary structures. Within DNA, pairs of two nucleotides named guanine (G) and cytosine (C) on adjacent strands form three hydrogen bonds, whereas pairs of adenine (A) and thymine (T) form two hydrogen bonds |
| Rate of Vaporization (10.4) | The higher the temp, the greater the # of molecules with sufficient kinetic energy to break the attractive forces that hold them together in the liquid. The larger the surface area of the liquid, the greater the number of molecules on the surface in a position to enter the gas phase. The stronger the intermolecular forces, the greater the kinetic energy needed for a molecule to escape the surface, and the smaller the number of molecules in the population that have this energy. |
| Rate of Condensation (10.4) | Within an enclosed space: some molecules will condense after colliding with the surfaces that confine them. As concentration of vapor increases, so does the rate of condensation (more collisions occur). |
| Dynamic Equilibrium (10.4) | Eventually, within an enclosed encounter, the rates of condensation and vaporization equalize. Molecules begin to enter and leave the liquid phase at the same rate. This means there will be no net change in the amount of liquid within the container. |
| Vapor Pressure (10.4) | As the number of water molecules in the vapor increases at constant temperature (T), so does the partial pressure (Pₕ₂ₒ). At a certain temperature (T), the partial pressure of water vapor at equilibrium with liquid water (within its closed system) is called its vapor pressure. |
| Difference in Vapor Pressures (10.4) | Liquids have their own characteristic vapor pressures. Some substances, such as [Br₂], have higher pressures than water because the bonds that hold them together are weaker than the hydrogen bonds that must be broken to vaporize water. For many substances, as (T) increases so does vapor pressure. |
| Normal Boiling Point (10.4) | The temperature at which the vapor pressure of a substance reaches 1 atm or 760 torr. Most chemical reactions in nature, the human body, and the laboratory take place at around 1 atm of atmospheric pressure. |
| Pressure Unit [atm] (10.4) | A standard atmosphere. Equal to the average atmospheric pressure (weight exerted by the atmosphere) at sea level. Specifically 1 atm = 101,325 pascals, which is the SI unit of pressure. |
| Pressure Unit [Pa] (extra) | A pascal is one newton per square meter. It measures internal pressure and stress. The kilopascal is also commonly used (1 kPa = 1000 Pa). The pascal measures the stiffness, tensile strength and compressive strength of materials. Also equivalent to the SI unit of energy density, the joule per cubic meter. |
| Pressure Unit [Torr] (10.4) | A unit of pressure used in measuring partial vacuums, equal to 133.32 pascals. It is defined to be exactly 1/760 of one standard atmosphere [atm]. |
| Volatile Liquids (10.4) | Liquid substances with significant vapor pressures and vaporization rates at room temperature. This can be used as an adjective: as in, "most volatile" or "nonvolatile" depending on the behavior of the substance. For example, Ethylene Glycol is a nonvolatile substance with a normal boiling point of 197.3° C, which allows it to be used in automobile cooling systems. |
| Clausius-Clapeyron Equation (10.4) | Plotting the natural logarithm of its vapor pressure (Pᵥₐₚ) versus (1/T) creates a straight line. The equation of this line fits the Clausius-Clapeyron equation. Which is: [ln(Pᵥₐₚ) = -(ΔHᵥₐₚ/R) (1/T) + C]. [ΔHᵥₐₚ] is the enthalpy of vaporization, [R] is the gas constant, [T] is in kelvin, and [C] is a constant dependent on the identity of the liquid. |
| Gas Constant (10.4) | 8.313 J / (mol·K) |
| Using the Clausius-Clapeyron Equation (10.4) | Solve for [C] by separating the equation into each desired [T] value. (Page 522 for example). Then, rearrange into: [ln(Pᵥₐₚ,T₂/Pᵥₐₚ,T₁) = -(ΔHᵥₐₚ/R)(1/T₂ - 1/T₁)]. This provides a way to solve for ΔHᵥₐₚ if vapor pressure at two [T]s are known, or the vapor pressure at a given temperature (Pᵥₐₚ,T₂ via T₂) or vice versa. |
| Phase Diagrams (10.5) | Used to display which phase of a substance is stable via the temperature (x-coordinates) and pressure (y-coordinates) it is exposed to. These graphs have three sections for each of the phases, and a fourth labelled the "supercritical" region. The lines separating each region are called equilibrium lines (on them, the substance is at equilibrium). |
| Why is water's phase diagram unique? (10.5) | The liquid/solid equilibrium line shows that there is a DECREASE in melting point as pressure increases. This is because water expands when it freezes. Most other substances contract when they freeze, but the hydrogen bonds in ice create a more open structure. Applying force (pressure) to this "open structure" solid may force a shift into the liquid state. |
| Triple Point (10.5) | Identifies the temp/pressure at which all three physical states are at equilibrium with one another. |
| Critical Point (10.5) | The pressure at which the liquid and gaseous states are indistinguishable from each other. Located at the end of the boiling/condensation line. Thermal expansion at this high temperature causes the liquid to be less dense, while the high pressure confines the gas into a smaller area increasing its density. Thus gas and liquid become visibly equivalent. |
| Supercritical Fluid (10.5) | A substance that is exposed to temp/pressure above its critical temperature and pressure. This fluid has both gaseous and liquid properties. For example, it can dissolve substances like a liquid and also penetrate materials like a gas. |
| Application of Supercritical Fluids (10.5) | These fluids have several different intermolecular forces acting upon them. Supercritical carbon dioxide and water are sometimes mixed to generate fluids that can selectively dissolve specific substances from a mixture while leaving others untouched. |
| Sublimation/Deposition of Carbon Dioxide (10.5) | At a pressure of 1 atm, carbon dioxide shifts directly from gas to liquid or vice versa. This is why it is called "dry ice" and creates no "wet" liquid in our normal atmosphere. |
| Slope of Solid/Liquid Equilibrium Line (10.5) | If this line seems to go backwards, toward the y-axis, that suggests that a substance expands when it is in a solid state (like water). If this line goes forward (similarly to the liquid/gas line) then the substance expands when it melts. This concept can be used to predict if solid substances will float on their liquid counterparts (water will, carbon dioxide won't). |
| Surface Tension (10.6) | The resistance of a liquid to an increase in its surface area. This represents the energy required to move molecules apart on the surface of a sample of liquid water so that an object denser than water can break through the surface and settle to the bottom. When surface tension >> downwards force, the object exerting the force atop the water will float. In H₂O, this is made possible by hydrogen bonding (water has a high surface tension). Water forms a concave meniscus due to the silicon dioxide in glass. |
| Surface Tension of H₂O (10.6) | 7.29 x 10²² J/m² at 25°C (energy required to break through) Remember that 23 kJ/mol is needed to break hydrogen bonds. |
| Cohesive Force (10.6) | Interactions between the same type of molecule (ex. cohesive forces are hydrogen bonds between water molecules). |
| Adhesive Force (10.6) | Interactions between unlike molecules (ex. dipole-dipole interactions between water molecules and polar Si―O―Si groups on the surface of glass). |
| Capillary Action (10.6) | Within a thin enough glass tube the cohesive and adhesive forces of water will pull it directly upward. It stops moving upward when the forces are balanced by gravity. |
| Viscosity (10.6) | Resistance to flow. Water is more viscous than gasoline (which is large and nonpolar) because of the strength of its hydrogen bonds. Dispersion and dipole forces effect viscosity. |
| Solid H₂O's Structure (10.6) | When water cools into a solid, the molecules form an extensive and open hexagonal network in the solid phase that leaves, on average, more space between the molecules than in liquid water. The maximum density of water occurs at 4°C which creates autumnal turnover and safe hibernation/insulation for aquatic creatures. Water has a significant thermocline depending on the season (a difference in temperature between layers), which creates a rich ecological environment. |
| Nonpolar Compounds (10.7) | Have limited solubility in water and other polar solvents. This is because the dissolution process relies on dipole–induced dipole interactions, which are much weaker than dipole–dipole interactions. |
| Solubility (10.7) | Is generally determined by competition between the strengths of solvent–solvent and solute–solute interactions that inhibit the dissolving process and the solute–solvent interactions that promote it. [Remember: like dissolves like] |
| Miscible (10.7) | The property of two substances to mix in all proportions, forming a homogeneous mixture. The term is most often applied to liquids but also applies to solids and gases. For example, water and ethanol are miscible because they mix in all proportions. |
| Volume Unit [μg] (10.7) | Microgram; a unit of mass equal to one millionth (1×10⁻⁶) of a gram. |
| Surfactants (10.7) | Composed of long hydrocarbon chains bonded to ionic groups, create micelles to attract and dissolve grease/fat/oil. A usual ingredient found in household soap and detergent. |
| Hydrophilic & Hydrophobic (10.7) | Water "loving" and water "fearing" behaviors (a characteristic of the biological phospholipid bilayer, one of the most important cellular structures). |
| Solubility of Gases in Water (10.8) | Most gases become less soluble in water as temperature increases. It also decreases with decreasing pressure (opening a bottle containing a carbonated beverage creates the rapid expulsion of some previously-soluble gases). At a higher temperature, [O₂] molecules have more kinetic energy, which allows them to overcome the dipole-induced dipole interactions that keep them in solution. |
| Henry's Law (10.8) | The relationship between gas solubility in a liquid and the partial pressure of the gas in the environment surrounding is described by the equation [Cᵍᵃˢ = kₕPᵍᵃˢ]. Where [Cᵍᵃˢ] represents the concentration (solubility) of a gas in a solvent, [kₕ] is the Henry’s law constant for the gas in solvent, and [Pᵍᵃˢ] is the partial pressure of the gas in the environment surrounding the solvent. When [Cᵍᵃˢ] is expressed in M and [Pᵍᵃˢ] in atm, the units of Henry’s law constant are mol/(L·atm). |
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