General Chemistry Ch. 12 - Intermolecular Forces: Liquids, Solids, Phase changes
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| Phase | show 🗑
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| show | Bonding forces that exist WITHIN each molecule and influence the chemical properties of the substance
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| Intermolecular forces | show 🗑
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| Phase changes | show 🗑
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| Roughly, how do phase changes come about? | show 🗑
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| show | 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.
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| show | Conforms to shape and volume of container; high compressibility because molecules are so spread apart; low viscosity (flows and diffuses very easily)
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| show | 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.
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| Solid: shape and volume, compressibility, and viscosity | show 🗑
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| List all the types of phase changes | show 🗑
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| At what point do metals freeze? | show 🗑
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| show | Exothermic, heat is released as the particles become fixed and solid.
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| Melting and vaporizing: exothermic or endothermic? | show 🗑
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| Heat of vaporization and fusion | show 🗑
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| How to calculate deltaH_subl | show 🗑
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| show | A curve that shows the changes that occur when heat is added or removed from a particular sample of matter at a constant rate.
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| show | 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.
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| Change of heat during a phase change | show 🗑
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| Are phase changes reversible? Can they reach equilibrium? | show 🗑
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| show | When the rate of condensation is equal to the rate of vaporization. The pressure of the vapor is constant at that temperature.
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| show | Because although it appears static macroscopically, molecularly the particles are leaving and entering the liquid surface at equal rates.
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| show | Vapor pressure.
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| show | Temperature. RAISING THE TEMPERATURE INCREASED THE VAPOR PRESSURE.
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| show | 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.
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| Clausius-Clapeyron equation: what does it do? | show 🗑
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| show | ln(P2/P1) = [(-deltaH_vap)/R]*(1/T2 – 1/T1)
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| show | The temperature at which the vapor pressure equals the external pressure.
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| Boiling point’s relationship with elevation | show 🗑
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| Melting point | show 🗑
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| Why does sublimation occur? | show 🗑
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| show | A diagram which combines the liquid-gas, solid-liquid, and solid-gas curves
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| What do the lines between the regions in the phase diagrams represent? | show 🗑
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| How is the slope of line between solid and liquid phases different in H2O as compared to most other substances? | show 🗑
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| show | 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.
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| show | The pressure and temperature at which three phases are in equilibrium.
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| show | Intramolecular forces are relatively strong because they involve larger charges and are closer together than that of the relatively smaller intermolecular forces.
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| List and identify all distances in a molecule | show 🗑
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| Types of intermolecular forces (in order of strength, strongest-weakest) | show 🗑
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| Ion-dipole force | show 🗑
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| show | 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
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| H-bond | show 🗑
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| Polarizability | show 🗑
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| Polarizability trends | show 🗑
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| show | Dipoles caused by the distortion of molecules’ electron clouds by other particles.
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| The most universal intermolecular force | show 🗑
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| show | 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.
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| show | Instantaneous dipole-induced dipole forces.
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| show | Dispersion forces are dominant because they’re most prevalent. They make up most of the bonding forces of most substances.
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| show | The polarizability of the particles, which closely correlates with molar mass. I.e., as molar mass increases so do dispersion forces.
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| What else affects the relative strength of dispersion forces? | show 🗑
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| show | Liquid
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| Why does a liquid surface tend to have the smallest possible area? | show 🗑
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| show | The greater the strength of the attractions, the greater the surface tension.
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| How do surfactants work to decrease surface tension of water? | show 🗑
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| show | 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.
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| Differences in the meniscuses of H2O and Hg | show 🗑
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| show | Resistance to flow.
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| show | 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.
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| Relation of viscosity to molecular shape | show 🗑
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| show | Capillary action of the cotton fibers acting on the water on your skin.
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| Solvent properties of water | show 🗑
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| show | Aquatic animals could not survive without dissolved O2; aquatic plants couldn’t survive without dissolved CO, etc.
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| show | 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.
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| Thermal properties of water (2) | show 🗑
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| show | 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.
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| Density of solid and liquid water | show 🗑
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| Crystalline solids | show 🗑
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| show | Have poorly defined shapes because their particles lack long-range ordering throughout the sample.
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| show | A regular pattern of particles with identical surroundings. The arrangement of the points within the particles defines the lattice.
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| Unit cell | show 🗑
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| show | The coordination number of a particle in a crystal is the number of nearest neighbors surrounding it.
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| How many types of crystal systems are there? Which one will we be focusing on? | show 🗑
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| Which substances occur as cubic lattices? | show 🗑
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| show | (1) Simple cubic unit cell, (2) body-centered cubic unit cell, (3) face-centered cubic-unit cell
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| show | 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.
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| show | 6: 4 in its own layer, 1 in the layer above, and 1 in the layer below. Atom number: 1
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| Body-centered cubic unit cell | show 🗑
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| Coordination number of each particle in the body-centered cubic unit cell & atom number | show 🗑
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| show | Identical particles lie at each corner and in the center of each face but not in the center of the cube
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| Coordination number of each particle in the face-centered cubic unit cell & atom number | show 🗑
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| How to calculate atoms per unit cell | show 🗑
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| The higher the coordination number a crystal is … | show 🗑
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| show | 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.
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| show | 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.
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| Packing efficiency: Hexagonal and face-centered cubic unit cells | show 🗑
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| show | In either hexagonal closest packing or cubic closest packing
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| Five most important types of solids | show 🗑
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| Atomic solids | show 🗑
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| How do atomic solids crystallize? | show 🗑
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| Molecular solids | show 🗑
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| Example of molecular solid | show 🗑
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| show | 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
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| show | 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).
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| show | Several different structures, but many use cubic closest packing.
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| Ionic solid: Sodium chloride structure: prevalence? Ratio? | show 🗑
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| Ionic solid: Sodium chloride structure: how does it arise? | show 🗑
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| show | 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.
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| show | 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.
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| show | 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.
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| The properties of ionic solids are a direct consequence of … | show 🗑
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| show | 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.
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| show | Because they have the same structure but twice as many delocalized valence electrons.
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| show | 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
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| show | 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.
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| Diamond | show 🗑
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| show | 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.
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| Amorphous solids | show 🗑
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| show | 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.
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| show | A model of metallic bonding developed in the field of quantum mechanics which is more quantitative and more useful than the electron sea model.
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| How does band theory work? | show 🗑
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| In metals, the valence and conduction bands are contiguous: what does that mean? | show 🗑
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| show | 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.
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| show | 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.
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| show | 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.
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| Insulators (nonmetals) | show 🗑
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| Superconductivity | show 🗑
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| Crystal defects | show 🗑
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| show | 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.
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| n-type semiconductor | show 🗑
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| p-type semiconductor | show 🗑
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| p-n junction. Describe when current flows and when it doesn’t | show 🗑
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| Rectifier | show 🗑
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| show | 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.
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| Liquid crystals | show 🗑
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| Anisotropic | show 🗑
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| show | 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.
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| Orientation of liquid crystals | show 🗑
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| Liquid crystals can arise in two general phases: | show 🗑
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| show | Develops as a result of change in temperature.
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| show | Occurs in a solution as the result of changes in concentration, but the conditions for forming such a phase vary for different substances.
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| show | Nematic, cholesteric and smectic
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| Nematic phase | show 🗑
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| Cholesteric phase | show 🗑
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| Because of the helical (corkscrew) arrangement of the cholesteric phase, what is it often called? | show 🗑
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| show | The molecules lie parallel to each other with their ends aligned. The layers are stacked directly over each other.
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| show | Nonmetallic, nonpolymeric solids that are HARDENED BY HEATING to high temperatures. Clay ceramics consist of silicate microcrystals suspended in a glassy cementing medium.
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| Why do ceramics become harder when heated? | show 🗑
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| Why are clay ceramics (e.g. bricks, porcelain, and glazes) useful? | show 🗑
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| show | An extremely large molecule (macromolecule) consisting of smaller molecules called monomers.
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| show | Repeat unit of a polymer
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| Examples of synthetic polymers | show 🗑
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| show | 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
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| Number-average molar mass of polymers | show 🗑
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| show | The backbone
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| show | The number of repeat units (degree of polymerization) * the length of each repeat unit l_o. I.e. n * l_o
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| show | 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.
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| show | It’s expressed by its radius of gyration, R_g
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| Radius of Gyration (R_g) | show 🗑
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| Polymer crystallinity | show 🗑
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| Flow behavior of polymers | show 🗑
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| How does molar mass of the polymer correlate with viscosity of solution? | show 🗑
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| show | 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.
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| show | Refers to a material that, when deformed, it returns to its original shape.
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| show | 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.
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| The more branches… | show 🗑
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| show | 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.
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| show | Branches that link one chain to another. The extent of the crosslinking can result in remarkable differences in properties.
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| Thermoplastic polymer | show 🗑
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| show | 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.
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| show | Above their glass transition temperatures, many thermosets become elastomers, polymers that can be stretched and immediately spring back to their initial shapes when released.
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| Examples of elastomers | show 🗑
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| show | Consists of one type of monomer (A-A-A-A-A- …)
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| Copolymer | show 🗑
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| The simplest copolymer | show 🗑
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| show | the A and B portions form their own random coils.
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| ABA block copolymer | show 🗑
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| Thermoplastic elastomers | show 🗑
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| show | The science and engineering of nanoscale systems, whose sizes range from 1 to 100 nm.
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| show | Self-assembly and controlled orientation
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| show | The ability of smaller, simpler parts to organize themselves into a larger, more complex whole. E.g. atoms or molecules aggregating through intermolecular forces.
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| show | The positioning of two molecules near each other long enough for intermolecular forces to result in a change. Enzymes work this way.
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| Quantum dots | show 🗑
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| show | 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.
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| Nanostructured materials | show 🗑
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| show | 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)
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| show | Viruses are biological machines. With regard to synthetic machines: nanovalves, nanopropellers, and even a nanocar have been developed.
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