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Chapter 11
Liquids, Solids, and Intermolecular Froces
| Term | Definition |
|---|---|
| Intermolecular Forces | Attractive forces that exist between all molecules and atoms |
| Properties of Gases | Low Density, Indefinite Shape, Indefinite Volume, Weak Strength of Intermolecular Forces |
| Properties of Liquids | High Density, Indefinite Shape, Definite Volume, Moderate Strength of Intermolecular Forces |
| Properties of Solids | High Density, Definite Shape, Definite Volume, Strong Strength of Intermolecular Forces |
| Two Types of Solids | 1. Crystalline.. 2. Amorphous |
| Crystalline Solids | The atoms or molecules that compose them are arranged in a well ordered three-dimensional array. Regular ordered structure. Long range order |
| Amorphous Solids | The atoms of molecules that compose them have no long range order. No long range order |
| Dispersion Forces(London Forces) | An intermolecular force exhibited by all atoms and molecules that results from fluctuations in the electron distribution. An instantaneous dipole on any one helium atom induces instantaneous dipole on neighboring atoms, which then attract one another |
| Molar Mass Effect on Dispersion Forces | If all other variables are held constant, the dispersion force increases with increasing molar mass because molecules or atoms of higher molar mass generally have more electrons dispersed over a greater volume |
| Dispersion Force and Molecular Shape | The straight shape of n-pentane molecules allows them to interact with on another along the entire length of the molecules(So Stronger DP). The nearly spherical shape of neopentane molecules allows for only a small area of interaction(So Weaker DP). |
| Dipole - Dipole Force | Exists in all molecules that are polar. An intermolecular force exhibited by polar molecules that results from the uneven charge distribution |
| Permanent Dipole | A permanent separation of charge; a molecule with a permanent dipole always has a slightly negative charge at one end and a slightly positive charge at the other |
| Polar Molecule's Boiling Point Compared to Nonpolar Molecule's Boiling Point | Polar molecules have higher melting and boiling points than nonpolar molecules of similar molar mass |
| Miscibility (Polar-P) (Nonpolar-NP) | The ability to mix without separating into two states. The polarity of molecules composing liquids is important in determining miscibility of liquids. P liquids are miscible with other P liquids and NP liquids are miscible with other NP liquids |
| Hydrogen Bonding | A strong dipole-dipole attractive force between a hydrogen bonded to O, N, or F and one of these electronegative atoms on a neighboring molecule. Not a chemical bond |
| Ion-Dipole Force | Occurs when an ionic compound is mixed with a polar compound; it is especially important in aqueous solution of ionic compounds. An intermolecular force between an ion and the oppositely end of a polar molecule |
| 1. Summarizing Intermolecular Forces | Dispersion forces are present in all molecules and atoms and increase with increasing molar mass. These forces are always weak in small molecules but can be significant in molecules with high molar masses |
| 2. Summarizing Intermolecular Forces | Dipole-dipole forces are present in polar molecules |
| 3. Summarizing Intermolecular Forces | Hydrogen bonds, the strongest of the intermolecular forces that can occur in pure substances(second only to ion-dipole forces in general), are present in molecules containing hydrogen bonded directly to fluorine, oxygen, or nitrogen |
| 4. Summarizing Intermolecular Forces | Ion-dipole forces are present in mixtures of ionic compounds and polar compounds. These are very strong and are especially important in aqueous solution of ionic compounds |
| Strength of Intermolecular Forces from Weakest to Strongest | Dispersion(All molecules and atoms) ---> Dipole-dipole(Polar molecules) ---> Hydrogen bonding(Molecules containing H bonded to F, O, or N) ---> Ion-dipole(Mixtures of ionic compounds and polar compounds) |
| Surface Tension | The tendency of liquids to minimize their surface area |
| Surface Tension of a Liquid | The energy required to increase the surface area by a unit amount |
| Viscosity | The resistance of a liquid to flow. Measured in a unit called poise(P), defined as 1g/ cm*s. The viscosity of water at room temperature is approximately one centipoise(cP) |
| Capillary Action | The ability of a liquid to flow against gravity up a narrow tube |
| Cohesive Forces | The attraction between molecules in a liquid |
| Adhesive Forces | The attraction between these molecules and the surface of the tube |
| Vaporization | The phase transition from liquid to gas. Is an endothermic process because it takes energy to vaporize the molecules in a liquid |
| Condensation | The phase transition from gas to liquid. Is an exothermic process because heat is released when a gas condenses to a liquid |
| Volatile | Tending to vaporize easily |
| Nonvolatile | Not easily vaporized |
| Summarizing the Process of Vaporization | 1. The rate of vaporization increases with increasing temperature.. 2. The rate of vaporization increases with increasing surface area.. 3. The rate of vaporization increases with decreasing strength of intermolecular forces |
| Heat of Vaporization | The amount of heat required to vaporize one mole of a liquid to a gas is its heat of vaporization(ΔHvap). The heat of vaporization of water at its normal boiling point of 100 degrees Celcius or +40.7 kJ/mol |
| Dynamic Equilibrium | The point at which the rate of the reverse reaction or process equals the rate of the forward reaction of process. The condensation and vaporization continue at equal rates and the concentration of the water vapor above the liquid is constant |
| Vapor Pressure | The partial pressure of a vapor in dynamic equilibrium with its liquid |
| General Statement Describing The Tendency of a System in Dynamic Equilibrium to Return to Dynamic Equilibrium | When a system in dynamic equilibrium is disturbed, the system responds so as to minimize the disturbance and return to a state of equilibrium |
| Le Chatelier's Principle | If the pressure above a liquid-vapor system in equilibrium is decreased, some of the liquid evaporates, restoring the equilibrium pressure. If pressure is increased, some of the vapor condenses. bringing the pressure back down to the equilibrium pressure |
| Boiling Point of a Liquid | The temperature at which its vapor pressure equals the external pressure |
| Normal Boiling Point of a Liquid | The temperature at which its vapor pressure equals 1 atm |
| Critical Temperature | The temperature above which a liquid cannot exist, regardless of pressure. |
| Critical Pressure | The pressure required to bring about a transition to a liquid at the critical temperature |
| Supercritical Fluid | When the gas and the liquid states commingle |
| Sublimation | The transition from solid to gas |
| Deposition | The transition from gas to solid |
| Melting Point | The temperature at which the molecules of a solid have enough thermal energy to overcome intermolecular forces and become a liquid |
| Melting or Fusion | The transition from solid to liquid |
| Freezing | The transition from liquid to solid |
| Heat of Fusion(ΔHfus) | The amount of heat required to melt 1 mol of a solid |
| Heat of Fusion and Heat of Vaporization | In general, the heat of fusion is significantly less than the heat of vaporization |
| Heating from Solid To Gas | Can be divided into 5 segments... 1. Solid warming.. 2. Solid melting into liquid(temperature constant).. 3. Liquid warming.. 4. Liquid Vaporizing into Gas(temperature constant).. 5. Gas warming |
| Calculating Heat Needed to Go From Solid to Gas | Add all values for "q" 1. Solid Heating(q = m*Cs,solid*ΔT in J).. 2. Solid melting into liquid(q = n*ΔHfus in kJ).. 3. Liquid Heating(q = m*Cs,liquid*ΔT in J).. 4. Liquid vaporizing into Gas(q = n*ΔHvap in KJ).. 5. Gas warming(q = m*Cs,solid*ΔT in J) |
| Phase Diagram | A map of the state or phase of a substance as a function of pressure(on the y-axis) and temperature(on the x axis) |
| Triple Point in a Phase Diagram | Represents the unique set of conditions at which three states are equally stable and in equilibrium |
| Critical Point in a Phase Diagram | Represents the temperature and pressure above which a supercritical fluid exists |
| X-ray Diffraction | A powerful laboratory technique that allows for the determination of the arrangement of atoms in a crystal and the measuring of the distance between them |
| Crystalline Lattice | The regular arrangement of atoms in a crystalline solid |
| Unit Cell | The smallest divisible unit of a crystal that, when repeated in three dimensions, reproduces the entire crystal lattice |
| Coordination Number | The number of atoms with which each atom in a crystal lattice is in direct contact. This is a characteristic feature of any unit cell. The number of atoms with which a particular atom can strongly interact |
| Packing Efficiency | The percentage of the volume of the unit cell occupied by the spheres. The higher the coordination number, the greater the packing efficiency |
| Body-Centered Cubic Unit Cell | Consists of a cube with one atom at each corner and one atom in the very center of the cube |
| Face-Centered Cubic Unit Cell | A cube with one atom at each corner and one atom(of the same kind) in the center of each cube face |
| Molecular Solids | Those solids whose composite units are molecules |
| Ionic Solids | Those solids whose composite units are ions |
| Atomic Solids | Solids whose composite units are atoms; they can be classified into three categories... 1. Nonbonding atomic solids.. 2. Metallic atomic solids.. 3. Network-covalent solids |
| Nonbonding Atomic Solids | Are held together by relatively weak dispersion forces. Have low melting points that increase uniformly with molar mass |
| Metallic Atomic Solids | Held together by metallic bonds, which in the simplest model are represented by the interaction of metal cations with the sea of electrons that surround them. Ex: iron and gold |
| Network Covalent Atomic Solids | Held together by covalent bonds. Ex: diamond, graphite, and silicon dioxide |
| Band Theory | A model for bonding in atomic solids that comes from molecular orbital theory in which atomic orbitals combine and become delocalized over the entire crystal |
| Band Gap | An energy gap that exists between the valence band and conduction band of semiconductors and insulators |
| n-Type Semiconductor | A semiconductor that employs negatively charged electrons in the conduction band as the charge carriers |
| p-Type Semiconductor | A semiconductor that employs positively charged "holes" in the valence band as the charge carriers |
| p-n Junctions | Tiny areas in electronic circuits that have p-type semiconductors on one side and n-type on the other |
| Diodes | Circuit elements that allow the flow of electrical current in only one direction |
| Amplifiers | Elements that amplify a small electrical current into a larger one |
Created by:
TimChemistry2
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