Final Exam Review great review! covers midterm material aswell!

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CHEM 123
Rebecka Peterborough

CHEM 123 Final Review Intermolecular forces (Van der Waals) London Dispersion: caused by oscillations in the electron density in non-polar molecules or atoms; instantaneous non-permanent dipoles; size and shape determine polarizability -> strength of interaction; larger atom-weaker forces, smaller atom stronger forces Dipole-Dipole: partial changes attract each other; permanent dipole; depends on magnitude = u u /r 12 22 6 Hydrogen Bonds: form between H attached to an electronegative atom and the electronegative atom and the electronegative atom in another molecule; common ones are O F and N The stronger the forces the higher the melting and boiling points and vice versa; harder to break forces in longer molecules rather than a circular shaped molecule The intermolecular forces that occur between molecules are collectively known as van der Waals forces. The most common intermolecular forces of attraction are those between instantaneous and induced dipoles (dispersion forces, or London forces). The magnitudes of dispersion forces depend on how easily electron displacements within molecules cause a temporary imbalance of electron charge distribution, that it, on the polarizability of the molecule. In polar substances, there are also dipole-dipole forces. Some hydrogen-containing substances exhibit significant intermolecular attractions called hydrogen bonds, in which H atoms bonded to highly electronegative atoms N, O, or F in a molecule are simultaneously attracted to the other highly electronegative atoms in the same molecule or in different molecules. Hydrogen bonding has a profound effect on physical properties, such as boiling points and is a vital intermolecular force in living systems Phases of a Substance Solid: ordered arrays of molecules. Kinetic energy of the particles is much less than the intermolecular forces holding them together Liquid: disordered groups of molecules that are close together. Kinetic energy of the particles is still less than the intermolecular forces, but not as low as in the solid form Gas: diffused arrangements of molecules. Kinetic energy of the particles is greater than the intermolecular forces that hold them together Surface tension, the energy required to extend the surface of a liquid, and viscosity, a liquids resistance to flow, are properties related to intermolecular forces. Familiar phenomenon such as drop shape, meniscus formation, and capillary action depend on surface tension. Specifically, these phenomena are influenced by the balance between cohesive forces, intermolecular forces between molecules in a liquid and adhesive forces, intermolecular forces between liquid molecules and a surface. Vapour pressure, the pressure exerted by a vapour in equilibrium with a liquid, is a measure of the volatility of a liquid and is related to the strength of the intermolecular forces. The conversion of a liquid to a vapour is called vaporization; the reverse process is called condensation. The dependence of vapour pressure on temperature is represented by a vapour pressure curve and can be expressed in the Clausius-Clapeyron equation. When the pressure exerted by the escaping molecules from the surface of the liquid equals the pressure exerted by the molecules in the atmosphere, boiling is said to occur. The normal boiling point is the condition of temperature and pressure at which liquid and its vapour become indistinguishable. When crystalline solids are heated, a temperature is reached where the solid state is converted to a liquid melting occurs. The temperature at which this occurs is the melting point. When liquids are cooled, the crystalline material will form during the process of freezing, and the temperature at which this occurs is the freezing point. Under certain conditions solids can directly convert into vapour by the process of sublimation. The reverse process is called deposition. Among the properties of a solid affected by intermolecular forces are its sublimation (vapour) pressure and its melting point Phase Transitions sublimation solid to gas deposition gas to solid4 vaporization liquid to gas condensation gas to liquid melting solid to liquid solidify liquid to solid A phase diagram is a graphical plot of conditions under which solids, liquids, and gases (vapours) exist, as single phases or as two or more phases in equilibrium with one another. Significant points on a phase diagram are the triple point (where all three phases coexist), melting point, boiling point, and critical point, beyond which a supercritical fluid is possible. Some substances can exist in different forms in the solid start; such behaviour is called polymorphism. Generalizations for Phase Diagrams From low to high temperatures on a constant-pressure line enthalpy increases (heat absorbed) From low to high pressure on a constant-temperature line volume decreases (high pressure=high density) Critical Point conditions are reached where liquid and vapour become indistinguishable; highest point on vapour pressure curve; highest temperature at which a liquid can exist Viscosity resistance to flow. The stronger the interactions between particles the more resistant the liquid is to flow Vapour Pressure pressure exerted by a specific amount of vapour; when system is closed condensation = vaporization (equilibrium); always positive; increases as temperature increases Surface Tension the energy or work required to increase the surface area of a liquid. The stronger the interaction, the higher the surface tension. (Hydrogen bonds) Types of Packing Crystalline regular packing of atoms or molecules Amorphous irregular packing of atoms or molecules In network covalent solids, chemical bonds extend throughout a crystalline structure. For theses substances, the chemical bonds are themselves intermolecular forces. Lattice energy is the energy released when separated gaseous ions come together to form one mole of an ionic solid. Structures of Crystalline Structures Simple Cubic (SC) 1 sphere per unit cell; coordination number = 6 Body Centered Cubic (BCC) 2 spheres per unit cell; coordination number = 8 Face Centered Cubic (FCC) 4 spheres per unit cell; coordination number = 12 Some crystal structures can be described in terms of the packing of spheres. Depending on the way in which the spheres are packed, different unit cells are obtained. The hexagonal unit cell is obtained with hexagonal closest packed (hcp) spheres; a face centred cubic (fcc) unit cell is obtained with cubic closest packed spheres. A body centered cubic (bcc) unit cell is found in some cases where spheres are not packed as closely as in the hcp and fcc structures. The dimensions of the unit cell can be determined by X-ray diffraction, and these dimensions can be used in calculating atomic radii and densities. An important consideration with ionic crystals in that the ions are not all of the same size or charge. Ionic crystals can often be viewed as an array of anions with the cations fitting in the holes within the anion array. Summary of Hole Sizes Tetrahedral r= .225R Octahedral r = .414R Cubic r = .732R Lattice energies of ionic crystals can be related to certain atomic and thermodynamic properties by means of the Born-Fajans-Haber cycle. Born Haber Cycle 1. Sublime one mole of solid cation(Na) 2. Dissociate half a mole of anion (Cl 2 into one mole of anion(Cl) + 3. Ionize one mole of cation(Na ) 4. Convert one mole of anion(Cl) to charged anion(Cl ) - + - 5. Allow cation(Na ) and anion(Cl) to form one mole of product(NaCl) 5 Things that influence Reaction Rates 1. Chemical Nature stable compounds react slowly even if the products of the reaction are at a lower energy 2. Physical State molecules react when they bump into each other 3. Concentration molecules react by bumping into each other. The more molecules or atoms that are present, the more often they will bump into each other; the more often a reaction will occur 4. Temperature when a reaction occurs bonds break and new bonds form. The energy to break the bonds comes from the kinetic energy of the molecules as they collide. As the temperature increases the kinetic energy of the molecules increase. More collisions have sufficient energy to break the bonds in the reactant, and a great number of collisions will lead to product formation.
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