Chem 1040, Exam 1 study guide
Chem 1040, Exam 1 study guide CHEM 1040 - 003
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CHEM 1040 - 003
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This 10 page Study Guide was uploaded by Olivia Hammond on Monday February 8, 2016. The Study Guide belongs to CHEM 1040 - 003 at Auburn University taught by Ria Astrid Yngard in Spring 2016. Since its upload, it has received 94 views. For similar materials see Fundamental Chemistry II in Chemistry at Auburn University.
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Date Created: 02/08/16
CHEM 1040 YNGARD REVIEW EXAM 1 Highlight = Important Principle Highlight = Important Concept Highlight = Key Term Chapter 7.3 Intermolecular Forces Particles (atoms, molecules, or ions) in the condensed phases (solids and liquids) are held together by intermolecular forces, which are electrostatic attractions between opposite charges or partial charges. van der Waals forces : intermolecular forces acting between atoms or molecules in a pure substance; this includes dipoledipole interactions, hydrogen bonding, and London dispersion forces. London Dispersion forces (or just Dispersion forces): forces between instantaneous dipoles of molecules; act between all molecules, polar or non polar. DipoleDipole interactions : exist between polar molecules Hydrogen Bonding : an especially strong type of DipoleDipole interaction that occurs in molecules that have a Hatom attached to a highly electronegative atom such as Fluorine, Oxygen, or Nitrogen and a nonbonding electron pair on a Fluorine, Oxygen, or Nitrogen IronDipole Interactions: those that occur (in solutions) between ions and polar molecules (is prevalent when ionic bonding is present) The stronger the forces among the moles the more energy it will take to separate those molecules. Disperson < DipoleDipole < Hydrogen Bonding < IronDipole Forces Forces Forces Forces Chapter 12 Liquids and Solids 12.1 The Condensed Phases The magnitude of intermolecular forces in liquids and solids influences various properties including surface tension, vapor pressure, boiling point, and melting point. 12.2 Properties of Liquids Surface tension is the net pull inward on molecules at the surface of a liquid. Adhesion: the attraction of molecules of one kind for molecules of a different kind. [example: water “climbing” upward through thin glass tubes placed in water] Cohesion: the attraction of molecules for other molecules of the same kind. [example: water creates a dome like shape in a container] Capillary Action: in which liquid is drawn upward into a narrow tube. Viscosity is resistance to flow, reflecting how easily molecules more past one another. The stronger the intermolecular forces, the higher the viscosity The higher the temperature the lower the viscosity. Vapor Pressure: the pressure exerted when liquid and vapor states are in Dynamic Equilibrium Dynamic Equilibrium: when vaporization and condensation are occurring at the same rate A volatile substance has a high vapor pressure. The weaker your intermolecular forces, the greater your vapor pressure is. The ClausiusClapeyron Equation relates the vapor pressure of a substance to its absolute temperature P = vapor pressure R = gas constant T = absolute temperature C = a constant ∆H = molar heat of vaporization 12.5 Phase Changes The possible phase changes are melting or fusion [solid —> liquid], freezing [liquid—> solid], vaporization [liquid —> gas], condensation [gas —> liquid], sublimation [solid —> gas], and deposition [gas —> solid] - Molar Heat of Vaporization ( ∆Hvap): the amount of heat required to vaporize a mole of a substance at its boiling point - Molar Heat of Fusion ( ∆Hfus): the amount of energy required to melt a solid - Molar Heat of Sublimation ( ∆Hsub): is equal to the sum of the molar heat of fusion and vaporization ∆Hvap + ∆Hfus = ∆Hsub Heating Curve 12.6 Phase Diagrams A phase diagram indicates the phase of a substance under any combination of temperature and pressure. Lines between phases are called phase boundaries. The triple point is here all three phase boundaries meet. This is the temperature and pressure combination at which all three phases are in equilibrium. Critical temperature is the temperature above which a gas cannot be liquefied by applying temperature Critical pressure is the pressure necessary to liquefy a gas at its critical temperature Critical Point: substance at its critical temperature and Pressure. Above this point you cannot distinguish between the liquid and gas state. Chapter 13 Physical Properties of Solutions 13.1 Types of Solutions H omogeneous mixture is a mixture of two or more pure substances. The solute is uniformly dispersed throughout. A solution is made up of a solute that is dissolved in the solvent A solution that has water as its solvent is considered to be an aqueous solution Unsaturated Solution: contains less solute than the solvent has the chapatis to dissolve at specific temperature Saturated Solution: contains the maximum amount of solute that will dissolve in a solvent at a specific temperature Solubility:amount of solute dissolved in a given volume of a saturated solution at a a specific temperature Supersaturated Solution: contains more dissolved solute than is present in a saturated solution and generally unstable 13.2 A Molecular View of The Solution Process Solvation : solute molecules are separated from one another and surrounded by solvent molecules salvation depends on the relative strengths of these interaction between particles: solutesolute, solventsolvent, or solutesolvent interactions A system will have a tendency to loose energy to reach its most stable form. If something requires energy it is not in its most stable form. This idea is the driving force behind whether or not something will occur spontaneously. Solution formation may be endothermic or exothermic overall. An increase in entropy is the driving force for solution formation. Entropy : a measure of how dispersed or spread out energy is There is a natural tendency for entropy to increase or for the energy of a system to become more dispersed or spread out (unless something, like a barrier, is preventing that dispersal) Substances with similar intermolecular forces tend to be soluble in one another. “Like dissolves like.”: molecules that have similar intermolecular forces will mix well; however, different intermolecular forces do not mix well [example: oil and water Polar substances tend to dissolve in polar substances Nonpolar substances tend to dissolve in nonpolar substances Miscible : two liquids are soluble in each other 13.3 Concentration Unites In addition to molarity (M) and mole fraction (X), molality (m) and percent by mass are used to express the concentrations of solutions. Molality (m or mol/kg) = moles of solute mass of solvent (kg) Molarity is temperature dependent; therefore, the molarity changes as temperature changes. Molality is temperature independent; therefore it remains constant as temperature changes. Percent by Mass = mass of solute total mass of solution Percent by Mass , Parts per Million, and Parts per Billion are all temperature independent 13.4 Factors That Affect Solubility Increasing he temperature increases the solubility of most solids in water and decreases the solubility of most gases in water. Added temperature causes added energy ; therefore, gases are more likely to escape from the liquid into the gas phase Increasing the pressure increases the solubility of gases in water but does not affect the solubility of solids. Henry ’s Law: the solubility of a gas in a liquid is proportional to the pressure of the gas over the solution C = KP C = molar concentration K = Henry’s Law constant (mol/L atm) P = pressure of gas over solution (atm) 13.5 Colligative Properties Colligative Properties : properties that depend on the number of solute particles in a solution; depend on the concentration of solute particles regardless of whether those particles are atoms, molecules, or ions. A volatile substance is one that has a measurable vapor pressure. A nonvolatile substance is one that does not have a measurable vapor pressure. 1. Vapor Pressure Lowering: when a nonvolatile solute is dissolved in a liquid, the vapor pressure exerted y the liquid decreases Raoult ’s Law: the partial pressure of a solvent over a solution, P1, is given by the vapor pressure of the pure solvent, P1 , times the mole fraction of the solvent in the solution X1 X1 = mole fraction P1 = partial pressure of a solvent over a solution P1 = vapor pressure of the pure solvent The decrease in vapor pressure, ∆P, is directly proportional to the solute concentration expressed as a mole fraction. The larger the amount of solute means there is a smaller X(solvent) which equals to a lesser vapor pressure The larger the amount of solvent is means there is a larger X(solvent) which equals to a greater vapor pressure The smaller difference in entropy between the solution and gas phases, results in a decreased tendency for solvent molecules to enter the gas phase. This lowers the vapor pressure. The solvent in a solution will always exert a lower vapor pressure than the pure solvent. 2. BoilingPoint Elevation: The boiling point of the solution is greater that the boiling point of a pure solvent; A higher temperature is needed to make the solvent’s vapor pressure equal to atmospheric pressure. ∆Tb = Tb T b = (Kb)(m) m = molality Kb = the molal boiling point elevation constant of the solvent 3. Freezing Point Depression: the solution freezes at a lower temperature than does the pure solvent; this occurs regardless of the solute’s volatility. 0 ∆Tf = T f Tf = (Kf)(m) m = molality Kf = the molal freezing point depression constant of the solvent 4. Osmotic Pressure Osmosis : the selective passage of solvent molecules through a semipermeable membrane from a more dilute solution to a more concentrated one. Osmotic Pressure of a solution ( ∏): the pressure required to stop osmosis ∏= MRT ∏ = osmotic pressure M = molarity R = ideal gas constant T = absolute temperature Two solutions of equal concentration have the same osmotic pressure are said to be isotonic to each other. Hypotonic refers t a solution with a lower osmotic pressure Hypertonic refers to a solution with a higher osmotic pressure An example of osmotic pressure is red blood cells. In electrolyte solutions, the number of dissolved particles is increased by dissociation or ionization. The magnitudes of colligative properties are thus increased by the van’t Hoff factor (i), which indicates the degree of dissociation or ionization. When i is measured it is usually a little less than when it is calculated because of ion pair formation (figure 13.3 in book) Iron pair formation is when ion collide with each other and are held together by electrostatic forces for a brief period in time the number of particles in the solution is reduced, thus reducing the observed colligative properties The more diluted the solution is the closer you will be to the calculated value of i because there is less iron pair formation and less collisions because there is more space to spread out The more concentrated the solution is the more collisions there will be and, therefore, a greater chance for iron pair formation. 13.5 Calculations Using Colligative Properties Experimentally determined colligative properties can be used to calculate the molar mass of a nonelectrolyte or he percent dissociation (or percent ionization) of a weak electrolyte. 13.6 Colloids Colloids : a dispersion of particles of one substance throughout another substance Colloids can be distinguished from true solutions by the Tyndall effect, which is the scattering of visible light by colloidal particles An example of this is how light from cars is scattered in fog. Colloids are either hydrophilic (water loving) or hydrophobic (water fearing) Chapter 14 Entropy and Free Energy 14.1 Spontaneous Processes Spontaneous Processes: a process that does occur under a specific set of conditions without ongoing outside intervention. Nonspontaneous Process: a process that does not occur under a specific set of conditions 14.2 Entropy Entropy (S) : a measure of how dispersed the system’s energy is. Molecules exhibit several types of motion: Translational : movement of the entire molecule from one place to another Rotational : rotation of the molecule about on its axis or sigma bonds Vibrational : periodic motion of atoms within a molecule k: the Boltzman Constant = 1.33 X 10 J/K23 W: the number of possible arrangements The most probable state is the one with the largest number of possible arrangements. 14.3 Entropy Changes in a System Entropy change of a process can be calculated using standard entropy values or can be predicted qualitatively based on factors such as temperature, phase, and number of molecules. Whether or not a process is spontaneous depends on the change in enthalpy and the change in entropy of a system Standard Entropy: the absolute entropy of a substance at 1 atm Entropy trends: Temperature changes: as temperature increases, enthalpy increases Volume changes: as volume increases, enthalpy increases Phase Changes: entropy increases from solid to liquid to the gas phase [solid < liquid < gas]; entropy increases during melting, vaporization, and sublimation. Molar Mass/Complexity: molar mass/complexity increases, as entropy increases Chemical Reaction: reaction resulting in a greater number of gas molecules creates a greater entropy Dissolution of a Solute: the process of dissolving substance often leads to an increase in entropy; but not always the case with ionic solutes 14.4. Entropy Changes in The Universe The system is the part of the universe we are investigation [i.e. reaction]. The surroundings are everything else. The system and the surroundings make up the universe. Solving for ∆S surroundings: Second Law of Thermodynamics: ∆S universe must be positive (in the forward direction). The system may undergo a decrease in entropy as long as the surroundings undergoes a larger increase in entropy and vice versa. A process where in ∆S universe is negative is not spontaneous as written. An equilibrium process is one that does not occur spontaneously in either the net forward or net reverse direction ∆S > 0 for a spontaneous process ∆S < 0 for a nonspontaneous process ∆S = 0 for an equilibrium process According to the Third Law of Thermodynamics, the entropy of a perfectly crystalline substance at 0 K is zero. 14.5 Predicting Spontaneity The Gibbs Free energy (G) of a system is the energy available to do work. ∆G < 0 —> the forward reaction is spontaneous (∆S > 0) ∆G > 0 —> the forward reaction is nonspontaneous (∆S < 0) ∆G = 0 —> the system is at equilibrium Chapter 14.5-14.6 The standard free energy of reaction (∆Grxn ) is free energy change for a reaction when it o occurs under standard-state conditions (1atm, 1M, pure solid/liquid, Gf = 0 fro elements in their most stable form. n.m = coefﬁcients ∆Gf = standard free energy of formation In living systems, thermodynamically favorably reactions provide the free energy needed to drive necessary but thermodynamically unfavorable reactions. By linking these two processes together, creates one overall spontaneous process although individually the ﬁrst reaction is nonspontaneous.
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