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Chem.200-12.MolForce
General Chemistry Ch. 12 - Intermolecular Forces: Liquids, Solids, Phase changes
| Question | Answer |
|---|---|
| Phase | Any physically distinct, homogeneous part of a system. E.g. ice water consists of two phases: solid water and liquid water. |
| Intramolecular forces | Bonding forces that exist WITHIN each molecule and influence the chemical properties of the substance |
| Intermolecular forces | Nonbonding forces that exist BETWEEN the molecules and influence the physical properties of the substance |
| Phase changes | The changes from one phase to another |
| Roughly, how do phase changes come about? | Intermolecular forces combine with the particles’ kinetic energy to create the properties of each phase as well as phase changes. |
| Whether a substance is a liquid, solid, or gas depends on… | The interplay of the potential energy of the potential energy of the intermolecular attractions (Coulomb’s law) which tends to draw molecules together, and their kinetic energy (prop. to temp.) which tends to disperse them. |
| Gas: shape and volume, compressibility, and viscosity | Conforms to shape and volume of container; high compressibility because molecules are so spread apart; low viscosity (flows and diffuses very easily) |
| Liquid: shape and volume, compressibility, and viscosity | Conforms to shape of container but volume limited by surface; Compressibility is very low; viscosity is moderate. Note: molecules are very close together but their kinetic energy still allows them to tumble over each other. |
| Solid: shape and volume, compressibility, and viscosity | Maintains own shape and volume; almost no compressibility; almost no ability to flow. Note: molecules sit very closely next to each other jiggling in space. |
| List all the types of phase changes | Gas-liquid: condensation; liquid-solid: freezing; solid-liquid: fusion, liquid-gas: vaporization, solid-gas: sublimation, gas-solid: deposition |
| At what point do metals freeze? | At very high temperatures actually. At that point they solidify. |
| Condensing and freezing: endothermic or exothermic? | Exothermic, heat is released as the particles become fixed and solid. |
| Melting and vaporizing: exothermic or endothermic? | Endothermic. |
| Heat of vaporization and fusion | (deltaH_vap & deltaH_fus) The enthalpy change per mole of a substance for a particular phase change. E.g. H2O(l) -> H2O(g) has a deltaH_vap of 40.7 kJ/mol at 100deg C. |
| How to calculate deltaH_subl | This can refer to solid-gas (+) or gas-solid (-) which is technically deposition, but still given the subscript subl but negative. To calculate, total up both fus and vap for the substance. |
| Heating-cooling curve | A curve that shows the changes that occur when heat is added or removed from a particular sample of matter at a constant rate. |
| Change of heat within a phase | Change of heat is accompanied by A CHANGE IN TERMPERATURE, which is associated with a change in average E_k as the most probable speed of the molecule changes. |
| Change of heat during a phase change | Change of heat occurs AT A CONSTANT TEMPERATURE which is associated with a change in E_p as the average distance between the molecules changes. |
| Are phase changes reversible? Can they reach equilibrium? | Phase changes are reversible and can reach equilibrium only in closed containers. |
| When is equilibrium reached in a closed container? What is said about that temperature? | When the rate of condensation is equal to the rate of vaporization. The pressure of the vapor is constant at that temperature. |
| Why the equilibrium in a closed container is called dynamic equilibrium? | Because although it appears static macroscopically, molecularly the particles are leaving and entering the liquid surface at equal rates. |
| What is the pressure exerted by the vapor at equilibrium called? | Vapor pressure. |
| Vapor pressure of a substance depends on its … | Temperature. RAISING THE TEMPERATURE INCREASED THE VAPOR PRESSURE. |
| Vapor pressure also depends on … | The intermolecular forces present. The E_k is the same for different substances at a given temperature. Thus, molecules with weaker intermolecular forces vaporize more easily. |
| Clausius-Clapeyron equation: what does it do? | It gives a way of finding heat of vaporization, the energy needed to vaporized 1 mol of molecules in a liquid state. |
| Clausius-Clapeyron equation | ln(P2/P1) = [(-deltaH_vap)/R]*(1/T2 – 1/T1) |
| Boiling point | The temperature at which the vapor pressure equals the external pressure. |
| Boiling point’s relationship with elevation | Higher the elevation, lower the boiling point. In space even cold water would boil because there is no air pressure. |
| Melting point | The state where melting rate equals the freezing rate. This dynamic equilibrium is called the melting point. |
| Why does sublimation occur? | Because the combination of intermolecular attractions and atmospheric pressure is not great enough to keep the particles near one another when they leave the solid state |
| Phase diagram | A diagram which combines the liquid-gas, solid-liquid, and solid-gas curves |
| What do the lines between the regions in the phase diagrams represent? | At any point on the line shows the pressure and temperature where the bordering phases exist at equilibrium. |
| How is the slope of line between solid and liquid phases different in H2O as compared to most other substances? | In most substances the line slopes to the right indicating that as pressure is increased at a constant temperature the substance will become a solid. But water slopes to the left indicating the opposite. This is because ice = less dense than water. |
| Critical point | At critical temperature and critical pressure the density of the liquid and vapor in a closed container are equal to each other and the phase boundary disappears. Beyond the critical temperature a supercritical fluid exists. |
| Triple point | The pressure and temperature at which three phases are in equilibrium. |
| Relative strengths of intramolecular and intermolecular forces | Intramolecular forces are relatively strong because they involve larger charges and are closer together than that of the relatively smaller intermolecular forces. |
| List and identify all distances in a molecule | Covalent radius; bond length, van der Waals radius; van der Waals distance. Pg. 450 for illustration. Note: van der Walls radius is always longer than covalent radius and determines the shortest distance over which intermolecular forces operate. |
| Types of intermolecular forces (in order of strength, strongest-weakest) | Ion-dipole; H-bonding; dipole-dipole; ion-induced dipole; dipole-induced dipole; and dispersion (London) forces |
| Ion-dipole force | When an ion and a nearby polar molecule (dipole) attract each other an ion-dipole force results. |
| Dipole-dipole force | When polar molecules lie near one another, as in liquids and solids, their partial charges act as tiny electric fields that orient them and give rise to dipole-dipole forces |
| H-bond | Partially positive H of one molecule is attracted to the partially negative lone pair on the N, O, or F of another molecule. |
| Polarizability | The ease with which the electron cloud of a particle can be distorted. Smaller atoms (or ions) are less polarizable than larger ones. |
| Polarizability trends | Increases down a group, decreases across a period. Cations are less polarizable than anions because cations are smaller. |
| Dipole-induced and ion-induced dipoles | Dipoles caused by the distortion of molecules’ electron clouds by other particles. |
| The most universal intermolecular force | Dispersion forces |
| Dispersion force | The intermolecular force primarily responsible for the condensed states of nonpolar substances. Caused by momentary oscillations of electron charge in atoms and are present between all particles. |
| Another way to describe dispersion forces | Instantaneous dipole-induced dipole forces. |
| What is the dominant intermolecular force? | Dispersion forces are dominant because they’re most prevalent. They make up most of the bonding forces of most substances. |
| The relative strength of dispersion forces depends on | The polarizability of the particles, which closely correlates with molar mass. I.e., as molar mass increases so do dispersion forces. |
| What else affects the relative strength of dispersion forces? | Molecular shape due to surface area differences. E.g. spherical molecules have much smaller dispersion forces than cylindrical molecules of the same molar mass. |
| What is the least understood phase of matter? | Liquid |
| Why does a liquid surface tend to have the smallest possible area? | Most particles in a liquid are attracted at all sides, except for the surface particles which are only attracted at one side, this creates a “net downward force”. |
| How does surface tension relate to the strength of the particles attraction to each other? | The greater the strength of the attractions, the greater the surface tension. |
| How do surfactants work to decrease surface tension of water? | Surfactants (surface-active agents), such as soaps, decrease the surface tension of water by congregating at the surface and disrupting the H-bonds. |
| Capillarity | The rising of a liquid through a narrow space against the pull of gravity is called capillary action, or capillarity. It results from a competition between the intermolecular forces within the liquid and those between the liquid and the tube walls. |
| Differences in the meniscuses of H2O and Hg | H2O is concave, Hg is convex. More H2O is binding to the wall of the glass, Hg however is more attracted to each other than the glass. |
| Viscosity | Resistance to flow. |
| Relation of viscosity to temperature. | Viscosity decreases with heating. This is because the molecules have higher E_k and can overcome resistance more efficiently. E.g. oil in a cooking pan being heated moves more freely. |
| Relation of viscosity to molecular shape | Small, spherical molecules flow more freely because they’re not “entangled” with each other. Long, spaghetti-like molecules have much higher viscosity. |
| Misc: How do towels dry you? | Capillary action of the cotton fibers acting on the water on your skin. |
| Solvent properties of water | The result of its polarity and H-bonding ability. It dissolves ions through ion-dipole forces; it dissolves many polar nonionic forces by forming H-bonds with them; and can even dissolve nonpolar molecules via dipole-dipole/dispersion. |
| How do the solvent properties of water contribute to life? | Aquatic animals could not survive without dissolved O2; aquatic plants couldn’t survive without dissolved CO, etc. |
| Thermal properties of water (1) | Water has high specific heat capacity, higher than almost every other liquid. Because of this, a large amount of energy provided by the sun won’t change the temperature too much. |
| Thermal properties of water (2) | Numerous strong H bonds give water a high heat of vaporization. It can absorb a lot of heat before it evaporates. At night water on earth and in its atmosphere will release its stored heat and warm the environment. |
| Surface properties of water | H bonds give water its high surface tension and high capillarity. Except for metals and molten salts, water has the highest surface tension. Plants use capillarity to absorb liquid from the soil, surface tension is used in numerous ways. |
| Density of solid and liquid water | The large spaces between the H2O molecules when solid give it a smaller density when compared to a liquid. When pressure is applied, hydrogen bonds break and the ice turn to liquid. |
| Crystalline solids | Solids that have a well-defined shape because their particles occur in an orderly arrangement. |
| Amorphous solids | Have poorly defined shapes because their particles lack long-range ordering throughout the sample. |
| Lattice | A regular pattern of particles with identical surroundings. The arrangement of the points within the particles defines the lattice. |
| Unit cell | The smallest portion of a crystal that, if repeated in all three directions, gives the crystal. |
| Coordination number | The coordination number of a particle in a crystal is the number of nearest neighbors surrounding it. |
| How many types of crystal systems are there? Which one will we be focusing on? | There are 7. We will be focusing on the cubic system which gives rise to the cubic lattice. |
| Which substances occur as cubic lattices? | The solid states of the majority of metallic elements, some covalent compounds, and many ionic compounds. |
| Three types of cubic-unit cells within the cubic system | (1) Simple cubic unit cell, (2) body-centered cubic unit cell, (3) face-centered cubic-unit cell |
| Simple cubic unit cell | The centers of eight identical particles define the corners of a cube. Attractions pull them together so they touch along the cube’s edges, but do not touch diagonally or through the center. |
| Coordination number of each particle in the simple cubic unit cell & atom number | 6: 4 in its own layer, 1 in the layer above, and 1 in the layer below. Atom number: 1 |
| Body-centered cubic unit cell | Identical particles lie at each corner and in the center of the cube. Those at the corners do not touch each other, but they all touch the center. |
| Coordination number of each particle in the body-centered cubic unit cell & atom number | 8: four above and four below. Atom number: 2 |
| Face-centered cubic unit cell | Identical particles lie at each corner and in the center of each face but not in the center of the cube |
| Coordination number of each particle in the face-centered cubic unit cell & atom number | 12: 4 in own layer, 4 on top, and 4 on bottom. Atom number: 4 |
| How to calculate atoms per unit cell | Corner particles = 1/8 of an atom, face particles = 1/2 of an atom, center particles = 1 atom |
| The higher the coordination number a crystal is … | The greater the number of particles in a given volume. |
| Packing efficiency: the simple cubic unit cell | Each layer of spheres is placed directly above another one. Only 52% of space is occupied under this arrangement which is low, meaning most atoms aren’t packed this way. |
| Packing efficiency: The body-centered cubic unit cell | Each layer fits above the open spaces of the layer below it. Packing efficiency is 68%. Several metallic elements, including all the alkali metals, have this arrangement. |
| Packing efficiency: Hexagonal and face-centered cubic unit cells | Similar to body-centered cubic unit cell except before the layers are stacked, they are shifted to create even smaller spaces between them. These can be arranged in 2 ways: hexagonal closest packing or cubic closest packing. Both = 74% |
| Most metallic elements crystallize in which arrangement? | In either hexagonal closest packing or cubic closest packing |
| Five most important types of solids | Atomic, molecular, ionic, metallic, network covalent |
| Atomic solids | Individual atoms that are held together by dispersion forces form an atomic solid. The noble gases are the only examples and their physical properties (low melting/boiling point) reflect how weak their bonds are. |
| How do atomic solids crystallize? | In cubic closest packed structure |
| Molecular solids | Individual molecules occupy the lattice points. Various combinations of dipole-dipole, dispersion, and H-bonding forces are at work for molecular solids. Higher melting points than atomic solids. Lower than all else. |
| Example of molecular solid | Methane (CH4) crystallizes in a face-centered cubic structure with the center of each carbon as the lattice point. Note: Molecular solids form into several different crystal structures. |
| Ionic solids | The unit cell contains particles with whole, rather than partial, charges. As a result, interparticle forces (ionic bonds) are stronger than van der Waals in atomic or molecular solids. The smaller ions in the solids are fit in holes created by large ions |
| How does the unit cell of an ionic solid relate to the chemical formula of the ionic compound? | The unit cell has the same cation/anion ratio as the empirical formula. (e.g. the ratio of Na to Cl is 1:1 in NaCl). |
| What type of crystal structures do ionic solids adopt? | Several different structures, but many use cubic closest packing. |
| Ionic solid: Sodium chloride structure: prevalence? Ratio? | Found in many compounds including most alkali metal halides/hydrides, alkaline earth sulfides/oxides, several transition metal oxides and sulfides, and most silver halides. 1/1 ratio |
| Ionic solid: Sodium chloride structure: how does it arise? | It arises when two arrays penetrate each other such that the smaller (e.g. Na+) ions end up in the holes between the larger (e.g. Cl-) ions. Face-centered cubic (cubic closest packing) structure |
| Ionic solid: Zinc blende (ZnS) structure | I can be pictured as two face-centered cubic arrays, one of Zn^2+ ions and the other S^2- ions, interpenetrating such that each ion is tetrahedrally surrounded by four ions of opposite charge (coordination number = 4). 1/1 ratio. |
| Ionic solid: Fluorite (CaF2) structure | Common among salts with a 1/2 cation/anion ratio, especially those having relatively large cations and relatively small anions. Unit cell = face-centered cubic array of Ca^2+ ions with F- occupying all eight available holes. |
| Ionic solid: Antifluorite structure | See in compounds having a cation/anion ratio of 2/1 and a relatively large anion (e.g. K2S). In this structure, the cations occupy all eight holes formed by the cubic closest packing of the anions. |
| The properties of ionic solids are a direct consequence of … | Fixed ions positions and very strong interionic forces which create high lattice energy. Thus properties are: high melting points and low electrical conductivities. |
| Metallic solids | The properties of metals—high electrical and thermal conductivity, luster, and malleability—result from the presence of delocalized electrons. Most metals crystallize in the two closest packed structures. |
| Why are group 2A metals harder and why do they have higher melting points than group 1A metals? | Because they have the same structure but twice as many delocalized valence electrons. |
| Network covalent solids | Atoms linked together in a network via covalent bonds. A diamond can be thought of a single molecule. Two popular network covalent compounds: graphite and diamond |
| Graphite | Stacked flat sheets of hexagonal carbon rings with a strong sigma bond framework and delocalized pi bonds. Extremely high melting point (because you need to break trillions of C-C bonds). Delocalized pi bonds allow it to be conductive. |
| Diamond | Sp3 hybridized carbons. Very high melting point. Bonding in all directions. Fixed thus nonconductive. |
| By far the most important network covalent compounds? | Silicates. They consist of extended arrays of covalently bonded silicon and oxygen atoms. Common example? Quartz (SiO2). Silicates form the structure of clays, rocks, minerals, etc. |
| Amorphous solids | Noncrystalline. Examples: charcoal, rubber, glass. |
| How is glass formed? | Crystalline quartz (SiO2) is melted and then cooled very rapidly before it can return back to its orderly crystalline structure. In this regard, glass is referred to as a supercooled liquid. |
| Band theory | A model of metallic bonding developed in the field of quantum mechanics which is more quantitative and more useful than the electron sea model. |
| How does band theory work? | As you accumulate more atoms of a metal, e.g. Li, you can create a quasi-MO diagram where you have an increasing number of higher energy MOs that are empty (conduction bands) and lower energy MOs that are filled (valence bands). |
| In metals, the valence and conduction bands are contiguous: what does that mean? | The electrons can jump from the filled valence band to the unfilled conduction band if they receive an infinitesimally small quantity of energy. I.e. the electrons are completely delocalized; they can travel freely throughout the material. |
| Metallic luster | Because the electrons are free to move around, they can absorb and release photons of many frequencies as they move between the valence and conduction bands. |
| Conductors (metals) and how does temperature affect conductivity? | The valence and conduction bands of a conductor have no gap between them, so electrons can flow easily even when a small electrical potential difference is applied. As temperature increases, greater random motion hinders movement. |
| Semiconductors (metalloids) and how temperature affects conductivity? | Relatively small energy gap exists between the valence and conduction bands. Thermally excited electrons can cross the gap, allowing a small current to flow. Thus, conductivity of a semiconductor increases when heated. |
| Insulators (nonmetals) | Gap between the bands is too large for electrons to jump even when the substance is heated, so no current is observed. |
| Superconductivity | At ordinary temperatures, electron flow is restricted by collisions with atoms vibrating in their lattice sites. |
| Crystal defects | A perfectly ordered crystal is attainable only if the crystal is grown very slowly under controlled conditions. If formed more rapidly, crystal defects typically form. E.g. misaligned planes of particles, missing particles, lodged debris, etc. |
| Doping | Pure silicon is a poor conductor at room temp because the gaps between the valence and conduction bands are too large. Its conductivity can be greatly enhanced by doping, adding small amount of other elements. |
| n-type semiconductor | The extra electrons of other elements added will enter an empty orbital in the conduction band, thus bridging the energy gap and increasing conductivity. N-type refers to the fact that *n*egative charges (electrons) are present. |
| p-type semiconductor | If doped with an atom that has less electrons than the semiconductor, a positive site will be created. This induces other electrons to fill the gap and increases conductivity. P-type, referring to *p*ositive charges being present. |
| p-n junction. Describe when current flows and when it doesn’t | In contact with each other, p and n-type semiconductors form a p-n junction. When a negative terminal is connected to the n-type portion and the positive is connected to the p-type portion, current flows n-to-p. Opposite: no current. |
| Rectifier | Since semiconductors are unidirectional they act as rectifiers: a device that converts AC current to DC current. |
| Transistors | p-type semiconductors are sandwiched between two n-type creates an n-p-n transistor. This creates adjacent p-n junctions. Traveling through these, the signal is amplified. |
| Liquid crystals | E.g. from the display in LCDs, these materials flow like liquids but, like crystalline solids, pack at the molecular level with a high degree of order. |
| Anisotropic | Gases and liquids are isotropic: their physical properties are the same in every direction. Liquid crystals are anisotropic, meaning properties depend on direction. |
| Two characteristics of the molecules that form liquid crystal phases: | 1) A long, cylindrical shape, and 2) and structure that allows intermolecular attractions through dispersion and dipole-dipole or H-bonding forces, but that inhibits perfect crystalline packing. |
| Orientation of liquid crystals | Low viscosity parallel to long axis, high viscosity perpendicular. Can be oriented vertically with use of electric fields due to molecular dipoles associated with liquid crystal molecules. |
| Liquid crystals can arise in two general phases: | Thermotropic phase and lyotropic phase |
| Thermotropic phase | Develops as a result of change in temperature. |
| Lyotropic phase | Occurs in a solution as the result of changes in concentration, but the conditions for forming such a phase vary for different substances. |
| Molecules that form liquid crystal phases can exhibit various types of order. Three common types are … | Nematic, cholesteric and smectic |
| Nematic phase | The molecules lie in the same direction but their ends are not aligned, much like a school of fish swimming in synchrony. The nematic phase is the least ordered type of liquid crystal phases. |
| Cholesteric phase | Somewhat more ordered than the nematic phase; the molecules lie in layers that each exhibit nematic-type ordering. Rather than lying in parallel fashion, however, each layer is rotated by a fixed angle with respect to the next layer. |
| Because of the helical (corkscrew) arrangement of the cholesteric phase, what is it often called? | Twisted nematic phase |
| Smectic phase | The molecules lie parallel to each other with their ends aligned. The layers are stacked directly over each other. |
| Ceramics | Nonmetallic, nonpolymeric solids that are HARDENED BY HEATING to high temperatures. Clay ceramics consist of silicate microcrystals suspended in a glassy cementing medium. |
| Why do ceramics become harder when heated? | They lose water and their atoms rearrange to an extended network of Si-centered and Al-centered tetrahedra of O atoms. |
| Why are clay ceramics (e.g. bricks, porcelain, and glazes) useful? | Because of their hardness and resistance to heat and chemicals. Also modern techniques have allowed for ceramics to be doped to the point that under high voltage they’re conductive. |
| Polymer | An extremely large molecule (macromolecule) consisting of smaller molecules called monomers. |
| Monomer | Repeat unit of a polymer |
| Examples of synthetic polymers | Plastics, rubbers, and specialized glasses. Virtually everything in modern-day life consists of synthetic polymers. |
| Polymer mass (M_polymer): Molar mass (AKA: ___) of a polymer depends on two parameters: 1) ___ 2) ___. Also, equation? | AKA: molecular weight. 1) the molar mass of the repeat unit (M_repeat) and 2) the degree of polymerization (n), or the number of repeat units in the chain. Equation: M_polymer = M_repeat * n. Note M = g/mol |
| Number-average molar mass of polymers | The degree of polymerization varies greatly from chain to chain. Thus an average, one approach in the number-average molar mass: M_n = (total mass of all chains)/(number of moles in chains) |
| The long axis of a polymer chain is called … | The backbone |
| The length of an *extended* backbone is… | The number of repeat units (degree of polymerization) * the length of each repeat unit l_o. I.e. n * l_o |
| Do polymers naturally exist as extended chains? | No, it is far more compact. Each of the repeat units move around randomly until they arrive in aggregate at the *random coil* shape that most polymers adopt. |
| How is the radius of a coiled polymer chain expressed? | It’s expressed by its radius of gyration, R_g |
| Radius of Gyration (R_g) | The average distance from the center of mass of the polymer to the outer edge of the coil. Mathematical expression: sqrt[(nl_o^2)/6] |
| Polymer crystallinity | At best a polymer is semicrystalline because only parts of the molecule align with parts of neighboring molecules (or with other parts of the same chain), while most of the chain remains as a random coil. |
| Flow behavior of polymers | A dissolved polymer increases the viscosity of the solution by interacting with the solvent. As the random coil of a polymer moves through a solution, solvent molecules are attracted to its exterior and interior through intermolecular forces. |
| How does molar mass of the polymer correlate with viscosity of solution? | Direct correlation. |
| Polymer glass | Transition from liquid to glass of polymers occurs over the narrow transition temperature (T_g). Examples include polystyrene in drinking cups and polycarbonate in compact disks. |
| Plastic mechanical behavior of polymers | Refers to a material that, when deformed, it returns to its original shape. |
| Branches | Smaller chains appended to a polymer backbone. As the number of branches increases, the chains cannot pack together as well, so the degree of crystallinity decreases and the polymer becomes less rigid. |
| The more branches… | The more flexible the polymer. E.g. plastic bags in grocery stores have many branches. |
| Dendrimers | The ultimate branched polymers. Prepared from monomers with three or more attachment points, so each monomer forms branches. In essence, these polymers have no backbone, just branches. |
| Crosslinks | Branches that link one chain to another. The extent of the crosslinking can result in remarkable differences in properties. |
| Thermoplastic polymer | In many cases, a small degree of crosslinking yields a thermoplastic polymer, one that still flows at high temperatures. |
| Thermoset polymer | A larger degree of crosslinking will eventually yield a thermoset polymer: one that can no longer flow because it has become a single network. Below their T_g level, some thermosets are extremely rigid and used in bicycle helmets and other things. |
| Elastomers | Above their glass transition temperatures, many thermosets become elastomers, polymers that can be stretched and immediately spring back to their initial shapes when released. |
| Examples of elastomers | Polydimethyl siloxane (breast implants); polybutadiene (rubber bands); polyisoprene (surgical gloves); and polychloroprene (footwear) |
| Homopolymer | Consists of one type of monomer (A-A-A-A-A- …) |
| Copolymer | Consists of two or more types of monomers |
| The simplest copolymer | AB block copolymer because a chain (block) of monomer A and a chain of monomer B are linked at one point: (A-A-A-A-B-B-B-B-B-B) |
| If the intermolecular forces between A and B portions of a AB block copolymer are weaker than those between different regions within each portion… | the A and B portions form their own random coils. |
| ABA block copolymer | Has A chains linked at each end of a B chain (A-A-A-B-B-B-A-A-A) |
| Thermoplastic elastomers | Some of the block copolymers act as thermoplastic elastomers, materials shaped at high temperature that become elastomers at room temperature. Used a lot in footwear industry. |
| Nanotechnology | The science and engineering of nanoscale systems, whose sizes range from 1 to 100 nm. |
| Two main features of nanoscale engineering that occur routinely in nature: | Self-assembly and controlled orientation |
| Self-assembly | The ability of smaller, simpler parts to organize themselves into a larger, more complex whole. E.g. atoms or molecules aggregating through intermolecular forces. |
| Controlled orientation | The positioning of two molecules near each other long enough for intermolecular forces to result in a change. Enzymes work this way. |
| Quantum dots | Nanoparticles of semiconducting materials, such as gallium arsenide (GaAs) or gallium selenide (GaSe), that are smaller than 10 nm. Rather than having a band of energy levels, the dot has discrete energy levels, similar to a single atom. |
| Why are quantum dots useful? | They can be manipulated chemically. They can absorb and release photons and as a result emit different colors. Used in biology to target and attach to cells/proteins so they can be viewed spectroscopically. |
| Nanostructured materials | The construction of bulk materials from nanoscale building blocks, increases strength, ductility, plasticity, and many other properties. |
| High-surface area materials | E.g. carbon nanotubes, porous membranes, and multilayer films. Carbon nanotubes have surface areas up to 4500 m^2/g (about 4 football fields per gram) |
| Nanomachines | Viruses are biological machines. With regard to synthetic machines: nanovalves, nanopropellers, and even a nanocar have been developed. |
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