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UNC / Chemistry / CHEM 102 / Determine the factors that affect the strength of london dispersion fo

Determine the factors that affect the strength of london dispersion fo

Determine the factors that affect the strength of london dispersion fo

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Final Exam Concept Review IMF’s 1. Distinguish the relative strengths of different types of intermolecular forces. 2. Determine the factors that affect the strength of london dispersion forces. 3. Determine whether a molecule has a permanent dipole and the relative strengths of dipoles. 4. Be familiar with what compounds will hydrogen bond and how many hydrogen bonds can form. 5. Sort a set of compounds by increasing boiling points. ● IMFs are the various forces of attraction that may exist between the atoms and molecules of a substance due to electrostatic phenomena. ● LDFs (weakest) → Dipole-Dipole → Hydrogen Bonding (strongest) London Dispersion Forces: ● LDFs are present in all condensed phases, regardless of the nature of the atoms or molecules composing the substance. ● Because the electrons of an atom or molecule are in constant motion, at any moment in time, an atom or molecule can develop a temporary, instantaneous dipole if its electrons are distributed asymmetrically. The presence of this dipole can distort the electrons of a neighboring atom or molecule, producing an induced dipole. These two rapidly fluctuating, temporary dipoles thus result in a relatively weak electrostatic attraction between the species. ● Dispersion forces that develop between atoms in different molecules can attract the two molecules to each other. The forces are relatively weak, however and become significant only when the molecules are very close. ● Larger and heavier atoms and molecules exhibit stronger dispersion forces than do smaller and lighter atoms and molecules. ● The increase in melting and boiling points with increasing atomic/molecular size may be rationalized by considering how the strength of dispersion forces is affected by the electronic structure of the atoms or molecules in the substance. ● In a larger atom, the valence electrons are farther from the nuclei than in a smaller atom. This means they are less tightly held and can more easily form the temporary dipoles that produce the attraction. ● The measure of how easy or difficult it is for another electrostatic charge to distort a molecule’s charge distribution (its electron cloud) is known as p​ olarizability​. ● A molecule that has a charge cloud that is easily distorted is said to be very polarizable and will have large dispersion forces; one with a charge cloud that is difficult to distort is not very polarizable and will have small dispersion forces. ● LDFs depend on the size and shape of the molecules. An elongated shape has a greater surface area available for contact between molecules, resulting in stronger dispersion forces. A compact shape offers a smaller surface area available for intermolecular contact and therefore, has weaker dispersion forces. ● HINT: Think about strips of VELCRO- the greater the area of the strip’s contact, the stronger the connection.Dipole-Dipole ● The electrostatic force between the partially positive end of one polar molecule and the partially negative end of another. ● POLAR molecules Hydrogen Bonding ● A strong type of dipole-dipole attraction. ● Examples: H2O, HF, HOH, H3N ● Hydrogen bonded to either N, O or F (HN, OH, HF) ● This is the strongest type of intermolecular force In order to determine how many hydrogen bonds can be created: - Which atoms can hydrogen bond and how many of those atoms do you have? -How many hydrogen bonds can each atom make (for example, oxygen can hydrogen bond with two H’s)? When sorting a set of compounds by increasing boiling point: -If a molecule has hydrogen bonding, it will probably have the highest boiling point. Sometimes, it can be tricky → If you are given two molecules that both hydrogen bond, the one that has more than one hydrogen bond will have the highest boiling point. -Remember the relative strengths of the IMFs: LDFs (weakest) → Dipole-Dipole → Hydrogen bond ex) Sort these compounds by increasing boiling point: HCl, H2O, SiH4 SiH4<HCl<H2O https://www.youtube.com/watch?v=90q7xl3ndJ8 Kinetics 1. Define the rate of reaction in terms of concentrations of reactants/products over time. 2. Relate the rate of appearance of products to the rate of disappearance of reactants using stoichiometry. 3. Explain why the initial rate of reaction is the best indicator of rate defined by instantaneous rates. 4. Explain how physical state, temperature, concentrations and presence of a catalyst affect the rate of a reaction. 5. Be familiar with the effect of order of a reaction on the units on the rate constant, k. 6. Determine when it is useful to use an integrated rate law. 7. Explain when the initial concentration is and is not important in half-life reactions. 8. Explain the conditions that must be met molecularly for a reaction to occur (collisions, activation energy, geometry). 9. Explain why a termolecular reaction is very rare. 10. Explain the way a catalyst speeds up a reaction, thinking about this in terms of a reaction coordinate diagram. 11. Draw reaction coordinate diagrams for endothermic, exothermic, catalyzed and uncatalyzed reactions. 12. Determine which step is the rate limiting step and why.● The rate of reaction is the change in the amount of a reactant or product per unit time. ● Reaction rates are determined by measuring the time dependence of some property that can be related to reactant or product amounts. In order to relate the rate of appearance of products to the rate of disappearance of reactants using stoichiometry: -You have to have a BALANCED equation. For example: 2H​2​O​2​(​aq)⟶2H​ 2​O(​l)+O​2​(​g) The rate at which hydrogen peroxide decomposes can be expressed in terms of the rate of change of its concentration: Rate of decomposition of H​2​O​2​ = -Δ[H​2​O​2​]/ Δt ● Rate expression- mathematical representation of the change in species concentration over time. https://www.youtube.com/watch?v=wyQEPWTGHao ^Pay attention to this video- she solves a rate expression problem and tells you a shortcut to have the correct units on k! ● At any specific time, the rate at which a reaction is proceeding is known as its instantaneous rate. The instantaneous rate of a reaction at “time zero” when the reaction commences is its initial rate. ● The best instantaneous rate is the rate at t=0, the initial rate. ● All reactions slow down over time. Therefore, the best indicator of the rate of a reaction is the initial rate (instantaneous rate at the beginning). Factors Affecting Rate of Reaction ● Physical State- Except for substances in the gaseous state or in solution, reactions occur at the boundary between two phases. The rate of a reaction between two phases depends on the surface contact between them. A finely divided solid has more surface area available for reaction than does one large piece of the same substance. ● Temperature- Chemical reactions typically occur faster at higher temperatures. Food can spoil quickly when left on the kitchen counter but the lower temperature inside of a refrigerator slows the process down so that the same food can remain fresh for days. We use a burner or a hot plate in the laboratory to increase the speed of reactions that proceed slowly at ordinary temperatures. ● Concentration- The rates of many reactions depend on the concentrations of the reactants. Rates usually increase when the concentration of one or more of the reactants increases. ● Presence of a Catalyst- Hydrogen peroxide solutions foam when poured onto an open wound because substances in the exposed tissues act as catalysts, increasing the rate of hydrogen peroxide’s decomposition. In the bottle in the medicine cabinet (where there is no catalyst present), complete decomposition can take months. ● Catalyst- a substance that increases the rate of a chemical reaction by lowering the activation energy without itself being consumed by the reaction.● Activation energy is the minimum amount of energy required for a chemical reaction to proceed in the forward direction. ● A catalyst increases the reaction rate by providing an alternative pathway or mechanism for the reaction to follow. Units of k  Reaction Order  Units of k (m + n)  mol​1-(m+n)​ L​(m+n)-1​ s​-1 zero  mol/L/s first  s​-1 second  L/mol/s third  mol​-2​L​2​s​-1


- Which atoms can hydrogen bond and how many of those atoms do you have?



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*Note that the units in the table can also be expressed in terms of molarity (M) instead of mol/L. Also, units of time other than the second (such as minutes, hours, days) may be used.* ● For a first order reaction, the half-life depends only on the ​rate constant​. ● For a second order reaction, the half-life depends on the inverse of the initial concentration. ● We use integrated rate laws and rate constants to relate concentrations and time. The rate law to use depends on the overall order of the reaction. To determine t, the time required for the initial concentration of a reactant to be reduced to some final value, we need to know the initial concentration. https://www.youtube.com/watch?v=WP9CHXm2_h0 ● Collision Theory: (What must happen for a reaction to occur?) 1. The rate of a reaction is proportional to the rate of reactant collisions. 2. The reacting species must collide in an orientation that allows contact between the atoms that will become bonded together in the product. 3. The collision must occur with adequate energy to permit mutual penetration of the reacting species’ valence shells so that the electrons can rearrange and form new bonds (and new chemical species). *ORIENTATION, ENERGY OF COLLISIONS*● Termolecular reactions involve three reacting molecules in one elementary step. Termolecular steps are relatively ​rare because they require the simultaneous collision of three molecules with sufficient energy in the correct orientation​. ● Unimolecular- single reactant ● Bimolecular- two reactants -It’s often the case that one step in a multistep reaction mechanism is significantly slower than the others. Because ​a reaction cannot proceed faster than its slowest step​, this step will limit the rate at which the overall reaction occurs. -The slowest step is called the rate-limiting step (or rate-determining step). https://www.youtube.com/watch?v=ShzW1LoQgoc https://www.youtube.com/watch?v=Lu_4SCqCLDs KNOW THIS CHART FORWARDS AND BACKWARDS!  Order  Rate Law  Integrated Rate Law Units of k  Straight Line Plot Zero  Rate= k  [A]= -kt + [A]​0 Mol L​-1​s​-1 [A] vs. t First  Rate= k[A]  ln[A]= -kt + ln[A]​0 s​-1 ln[A] vs. t Second  Rate= k[A]​2 Rate= k[A][B] 1/[A]= kt + 1/[A]​0 L s​-1​ mol​-1 1/[A] vs. t


-How many hydrogen bonds can each atom make (for example, oxygen can hydrogen bond with two H’s)?



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Reaction Coordinate Diagrams: Catalyzed vs. Uncatalyzed Reaction https://www.youtube.com/watch?v=48sLH9P8QK0● A catalyst speeds up the rate of a reaction by​ lowering the activation energy​. It also provides an alternate pathway or mechanism​ for the reaction to occur. Exothermic vs. Endothermic Reaction Coordinate Diagram Delta H→ Negative (exothermic- energy released), Delta H → Positive (endothermic-energy absorbed) Equilibrium 1. Explain what is happening molecularly in a system at equilibrium. 2. Determine whether an equilibrium is product or reactant favored given its equilibrium constant, K. 3. Explain the difference between the reaction quotient and equilibrium constant. 4. Determine whether a system will shift towards products or reactants given Q and K values. 5. Explain how changing the concentration of one species, changing the pressure of the system or changing the temperature will affect an equilibrium. 6. Be able to relate K to ice table- given equilibrium values find K, given K and initial, find equilibrium concentrations. 7. Explain how a catalyst affects equilibrium. ● In a chemical equilibrium, the forward and reverse reactions occur at equal rates, and the concentrations of products and reactants remain constant. ● At the molecular level, equilibrium is a dynamic state of equality, where the molecules continue to react at equal but opposite rates. ● Reaction quotient- Q; when evaluated using concentrations, it is called Q​c​. ● The reaction quotient is equal to the molar concentrations of the products of the chemical equation (multiplied together) over the reactants (also multiplied together), with each concentration raised to the power of the coefficient of that substance in the balanced chemical equation.Ex) 2NO​2​ (​g) ⇋ N​2​O​4​ (​g) Q​c​= [N​2​O​4​]/[ NO​2​]​2 ● As the reaction proceeds, the value of Q​c​ increases as the concentrations of the products increase and the concentrations of the reactants simultaneously decrease. ● When the reaction reaches equilibrium, the value of the reaction quotient no longer changes because the concentrations no longer change. ● When a mixture of reactants and products of a reaction reaches equilibrium at a given temperature, its reaction quotient always has the same value called the equilibrium constant (K) at that temperature. ● At equilibrium, the reaction quotient (Q​c​) is equal to K​c ● When a gas is involved- can also use the same process, but the concentrations will be pressure instead and it is Q​p​ and K​p​. K>1: products are favored K=1: neither reactants nor products are favored K>1: reactants are favored Q >K: reactants are favored Q=K: reaction is at equilibrium Q<K: products are favored Hint: If you remember what one of the quotients is, the opposite will be true for the other. For example, if you know that K being greater than 1 means products are favored, then Q being greater than K means reactants are favored. https://www.youtube.com/watch?v=g5wNg_dKsYY https://www.youtube.com/watch?v=DP-vWN1yXrY Catalysts do not affect the equilibrium state of a reaction. Calculating Equilibrium Constants: http://www.chem.purdue.edu/gchelp/howtosolveit/Equilibrium/Calculating_Equilibrium_Constants .htm#Kone Calculating the Reaction Quotient, Q: http://www.chem.purdue.edu/gchelp/howtosolveit/Equilibrium/Reaction_Quotient.htm Determining Equilibrium Quantities from Initial Quantities and K: http://www.chem.purdue.edu/gchelp/howtosolveit/Equilibrium/Reaction_Quotient.htm Acid Base 1. Define Lewis acids and bases.2. Define Bronsted-Lowry acids and bases. 3. Define the general equilibria for aqueous solutions of weak acids or weak bases. 4. Define the equilibrium for auto ionization in water. 5. List the strong acids and strong bases. 6. Define an amphiprotic species. 7. Relate pH, concentration of hydronium ion, Kw, and concentration of hydroxide ion. 8. Relate the relative strengths of weak acids based on their Ka values. 9. Define percent ionization and when this equation is useful. 10. Explain how you could find the Ka given pH and initial concentration of weak acid in solution. 11. Explain why the reaction of equal moles of strong acid and strong base yields a neutral solution with pH=7. 12. Explain why the reaction of equal moles of weak acid and strong base does not yield a neutral solution and describe what the expected pH range of this solution would be. 13. Define buffer capacity in terms of acid base ration and in terms of pH. 14. How would you determine the new pH of a buffer solution after adding a strong acid? A strong base? 15. Draw a titration curve and label the different regions: before addition, after some addition, at equivalence point, at end point. 16. How could you tell from looking at a titration curve whether it is a strong acid/strong base titration, a weak acid/strong base titration, or a weak base/strong acid titration? 17. Explain how you would determine the pH at the following points on a titration curve for a weak acid titrated with strong base: before adding any strong base, when some small amount of strong base has been added, at the equivalence point. 18. Determine what is present in solution at different points along a titration curve. ● Bronsted-Lowry acid: a compound that donates a proton to another compound ● Bronsted-Lowry base: a compound that accepts a proton ● Lewis acid: accepts an electron pair. ● Lewis base: donates an electron pair. ● Amphoteric/amphiprotic species: can act either as an acid or base. ● A solution is neutral if it contains equal concentrations of hydronium and hydroxide ions; acidic if it contains a greater concentration of hydronium ions than hydroxide ions; and basic if it contains a lesser concentration of hydronium ions than hydroxide ions. ● The pH of a solution is defined as: pH= -log [H​3​O​+​] ● [H​3​O​+​]= 10​-pH ● ^The same goes for OH- except you will use the concentration of hydroxide instead of hydronium. K​w​= [H​3​O​+​][OH​-​]; the ​value of K​w​ is 1.0 x 10​-14​​and so: 14.00 = pH + pOH The Auto-Ionization of Water, K​w​: http://www.chem.purdue.edu/gchelp/howtosolveit/Equilibrium/Autoionization_of_Water.htmIdentifying Conjugate Acid Base Pairs: ​https://www.youtube.com/watch?v=7qBRIWSA3Yc Strong Acids and Bases: Know these! Acids Bases HCl- hydrochloric acid HBr- hydrobromic acid HI- hydroiodic acid HNO​3​- nitric acid HClO​3​- chloric acid HClO​4​- perchloric acid H​2​SO​4​- sulfuric acid LiOH- lithium hydroxide NaOH- sodium hydroxide KOH- potassium hydroxide Ca(OH)​2​- calcium hydroxide Sr(OH)​2​- strontium hydroxide Ba(OH)​2​- barium hydroxide


● Collision Theory: (What must happen for a reaction to occur?



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Strong Acids and Strong Bases: Examples of Calculating the pH https://www.youtube.com/watch?v=gsu4gjrFApA pH, pOH, pKa, and pKb: http://www.chem.purdue.edu/gchelp/howtosolveit/Equilibrium/Calculating_pHandpOH.htm Ka values: -Low Ka/Kb values mean that the acid/base (respectively) does NOT dissociate well and it is a weak acid/weak base. Ka to pH and Percent Ionization: https://www.youtube.com/watch?v=PGJ6Svmkfnc How to Calculate Percent Ionization: [H​3​O​+​]​eq​/[HNO​2​]​0​ x 100 ● pH of strong acid-strong base → =7 (neutral), because both strong acids and strong bases dissociate 100%, H+=OH- ● pH of strong acid-weak base → <7 (acidic) ● pH of weak acid- strong base → >7 (basic) Buffer- a mixture of a weak acid and its conjugate base (or a mixture of a weak base and its conjugate acid). Video on Buffers: ​https://www.youtube.com/watch?v=8Fdt5WnYn1k Buffer capacity- the amount of acid or base that can be added to a given volume of a buffer solution before the pH changes significantly.How to Select Appropriate Buffers: 1. A good buffer mixture should have about equal concentrations of both of its components. 2. Weak acids and their salts are better as buffers for pHs less than 7; weak bases and their salts are better as buffers for pHs greater than 7. The Henderson Hasselbach Equation: pH = pKa + log [base]/[acid] https://www.youtube.com/watch?v=Ub7eLn6ddvg Buffer Solutions: http://www.chem.purdue.edu/gchelp/howtosolveit/Equilibrium/Buffers.htm Titration Curves: https://www.youtube.com/watch?v=JCAsOJYkn-s ^ I HIGHLY recommend watching this video,especially if you struggle with titrations. It is long (a little over an hour and a half), but it is a comprehensive video on titration curves and calculating pH from the titration curves. It provides a basic introduction to titrations. It shows you how to calculate the unknown concentration of an acid solution and how to determine the volume of base added to completely neutralize the acid and to reach the equivalence point. It also explains how to calculate the pH of acid base titration experiment before, at and beyond the equivalence point. It has plenty of examples, equations, formulas and practice problems. :) There is a list of topics in the description of the video. Notice the pH at the equivalence points for both of these graphs. They both have a pH of 7 at the equivalence point because the left graph is a strong acid-strong base titration curve and the right graph is a strong base-strong acid titration curve.Things to Remember Concerning Titration Curves: -The steepest region of the curve is the equivalence point! -The weaker the acid (smaller Ka or large pKa), the smaller the equivalence point region.^This screenshot shows a standard titration curve with the regions marked. Thermodynamics 1. Explain what entropy is and how the entropy of a system and its surroundings relate to the entropy of the universe. 2. List an example of a process that is spontaneous at one temperature range and non spontaneous at others. 3. Explains what gibbs free energy is and what is energetically favorable (positive or negative delta G). 4. Describe the relationship between gibbs free energy, entropy, and enthalpy. 5. Determine whether a process will be spontaneous at all temperatures, above/below a certain temperature, or nonspontaneous at all temperatures based on the signs of delta S and delta H. 6. Explain how gibbs free energy relates to equilibrium given a gibbs free energy reaction diagram. ● Entropy (S)- state function that is a measure of the matter and/or energy dispersal within a system, determined by the number of system microstates often described as a measure of the disorder of the system. → Microstate (W): possible configuration or arrangement of matter and energy within a system. ● The entropy change of the universe is the total entropy change of the system and its surroundings: DeltaS(universe)= DeltaS(surroundings) + DeltaS(system) ● In the solid phase, the atoms are restricted to nearly fixed positions with respect to each other and are capable of only modest oscillations about these positions. In the liquid phase, the atoms are free to move over and around each other, though they remain in relatively close proximity to one another. S(liquid) > S(solid) and the process of converting a substance from solid to liquid (melting) is characterized by an increase in entropy, Delta S > 0. By the same logic, the reciprocal process (freezing → liquid to solid) exhibits a decrease in entropy, Delta S <0. ● Entropy for any substance increases as the temperature increases. DeltaS(surroundings)= DeltaH(system)/T https://www.youtube.com/watch?v=ZsY4WcQOrfk ● Spontaneous process- one that occurs naturally under certain conditions → Ice melting at room temperature ● Nonspontaneous process- will not take place unless it is “driven” by the continual input of energy from an external source → Water freezes at room temperature ● A process that is spontaneous in one direction under a particular set of conditions is nonspontaneous in the reverse direction. ● Gibbs free energy change (G): G= H-TS ● Free energy is a state function and at constant temperature and pressure, the standard free energy change (DeltaG°): DeltaG: DeltaH-TDeltaS DeltaS(universe)>0 Delta G<0 Spontaneous DeltaS(universe)<0 Delta G>0 Nonspontaneous DeltaS(universe)=0 Delta G =0 Reversible (at equilibrium)

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https://www.youtube.com/watch?v=zslB7dr-eZA ^Another long one (54 minutes), but comprehensive and well worth it! Delta S and Delta H Relationship: KNOW THESE. 1. Both Delta H and Delta S are positive- This condition describes an endothermic process that involves an increase in system entropy. In this case, ​ Δ​G will be negative if the magnitude of the ​TΔ​S term is greater than Δ​H. If the ​TΔ​S term is less than Δ​H, the free energy change will be positive. Such a process is ​spontaneous at high temperatures and nonspontaneous at low temperatures. 2. Both Δ​H and Δ​S are negative. This condition describes an exothermic process that involves a decrease in system entropy. In this case, Δ​G will be negative if the magnitude of the ​TΔ​S term is less than Δ​H. If the ​TΔ​S term’s magnitude is greater than Δ​H, the free energy change will be positive. Such a process is ​spontaneous at low temperatures and nonspontaneous at high temperatures. 3. Δ​H is positive and Δ​S is negative. This condition describes an endothermic process that involves a decrease in system entropy. In this case, ΔG​ will be positive regardless of the temperature. Such a process is ​nonspontaneous at all temperatures. 4. Δ​H is negative and Δ​S is positive. This condition describes an exothermic process that involves an increase in system entropy. In this case, Δ​G will be negative regardless of the temperature. Such a process is ​spontaneous at all temperatures. Equilibrium and Delta G: Delta G= DeltaG° + RT lnQ, where R is the gas constant (8.314 J/K mol) and T is absolute temperature (Kelvin). Q is the reaction quotient. You can use this equation to predict the spontaneity for a process under any given set of conditions. Relations between Standard Free Energy Changes and Equilibrium Constants K DeltaG° What does this mean? >1 <0 Products are more abundant at equilibrium <1 >0 Reactants are more abundant at equilibrium =1 =0 Reactants and products are equally abundant at equilibrium

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Gases 1. Explain how temperature and energy are related. 2. Be familiar with conversion factors for pressure units. 3. Determine whether there is an inverse or direct relationship between the following variables: pressure, volume, moles, temperature. 4. Describe the 5 postulates of the kinetic molecular theory (KMT). 5. Explain the relationship between partial pressure, total pressure, and mole fraction. 6. Define diffusion and effusion. 7. Relate molar mass to rate of effusion of a gas particle. 8. Explain the differences between ideal gases and real gases and what the two van der waals constants correct for in the ideal gas law. ● As the temperature increases, energy increases. Temperature and energy have a direct relationship. As the temperature decreases, energy decreases as well. ● 1 torr= 1/760 atm, 1 torr= 1 mmHg ● Temperature and pressure have a direct relationship: as temperature increases, pressure increases as well. P1/T1= P2/T2 ● Volume and temperature have a direct relationship: (Charles’ Law) → The volume of a given amount of gas is directly proportional to its temperature on the kelvin scale when the pressure is held constant. ● Volume and pressure have an inverse relationship (Boyle’s Law) → As the volume increases, the pressure decreases. As the pressure increases, volume decreases. ● The ideal gas law: PV= nRT, where R is the ideal gas constant. ● Combined gas law: P1V1/T1= P2V2/T2 ● STP is 273.15 K and 1 atm. At STP, an ideal gas has a volume of about 22.4 L- standard molar volume. ● KMT Postulates 1. Gases are composed of molecules that are in continuous motion, travelling in straight lines and changing direction only when they collide with other molecules or with the walls of a container. 2. The molecules composing the gas are negligibly small compared to the distances between them. 3. The pressure exerted by a gas in a container results from collisions between the gas molecules and the container walls. 4. Gas molecules exert no attractive or repulsive forces on each other or the container walls; therefore, their collisions are ​elastic (do not involve a loss of energy). 5. The average kinetic energy of the gas molecules is proportional to the kelvin temperature of the gas.● Rate 1/Rate 2= Square root of MM1/MM2 (MM= molar mass) ● ^Example of Graham’s Law (rate of effusion): https://www.youtube.com/watch?v=YHKhxvdNyAs http://www.chemteam.info/GasLaw/Gas-Graham.html ● Lower molecular mass molecules diffuse faster than higher molecular mass molecules. ● Partial Pressure, Total Pressure and Mole Fraction: https://www.youtube.com/watch?v=agNtq38TA6o ● Real vs. Ideal Gas: -An ideal gas is made up of particles that are single points with no volume. This means that the atoms or molecules take up absolutely no room. Real gases are made up of atoms or molecules that actually take up some space, no matter how small. -An ideal gas is made up of particles that do not attract or repel one another. Real gases are made up of atoms or molecules that may attract one another strongly. -Real gases behave like ideal gases at “ordinary” temperatures and pressures, but if you heat them up and compress them to high pressure, then their behavior departs from ideal. Ideal vs. Real Gases: ​https://www.youtube.com/watch?v=Y3ymgxsuVfQ Liquids and Solids 1. Define the following properties of liquids: viscosity, surface tension, adhesion, cohesion, and capillary action. 2. Describe how to calculate the enthalpy change of the transition from a solid to a liquid to a gas. 3. Define vapor pressure. 4. Explain how intermolecular forces affect vapor pressure. 5. Differentiate the following types of solids based on their particle types, attractive forces and properties: ionic, metallic, covalent network, molecular. (I’ll provide examples of each) Properties of Liquids Defined: https://www.youtube.com/watch?v=BqQJPCdmIp8 -Viscosity: measure of a liquid’s resistance to flow -Surface tension: energy required to increase the area, or length, of a liquid surface by a given amount -Adhesion: force of attraction between molecules of different chemical identities -Cohesion: force of attraction between identical molecules-Capillary action: flow of liquid within a porous material due to the attraction of the liquid molecules to the surface of the material and to other liquid molecules https://www.youtube.com/watch?v=zz4KbvF_X-0 ^Specific heat, heat of fusion and vaporization: This video shows you how to calculate how much heat is needed to convert 200g of ice at -10 degrees C to 110 degrees C steam. *highly recommend* -Vapor pressure: (also equilibrium vapor pressure) pressure exerted by a vapor in equilibrium with a solid or a liquid at a given temperature → ​If the intermolecular forces are weak, then molecules can break out of the solid or liquid phase more easily into the gas phase​. ● Ionic Solids: -ex. Sodium chloride and nickel oxide -Composed of positive and negative ions held together by electrostatic attractions -High melting points due to the strong attractions between ions -Metal and nonmetal ● Metallic Solids: -ex. Crystals of copper, aluminum, iron -Formed by metal atoms -High thermal and electrical conductivity, luster and highly malleable -Hard and strong ● Covalent Network Solid -ex. Crystals of diamond, silicone, some other nonmetals and some covalent compounds such as silicon dioxide (sand) and silicon carbide (abrasive sandpaper) -Minerals -Held together by a network of covalent bonds -Hard, strong, high melting points ● Molecular Solid -ex. Ice, sucrose (table sugar), iodineTypes of Crystalline Solids and Their Properties Type of Solid Type of Particles Type of Attractions Properties Ionic Ions Ionic Bonds Hard, brittle, conducts electricity as a liquid but not as a solid, high to very high melting points Metallic Atoms of electropositive elements Metallic bonds Shiny, malleable, ductile, conducts heat and electricity well, variable hardness and melting temperature Covalent Network Atoms of electronegative elements Covalent bonds Very hard, not conductive, very high melting points Molecular Molecules (or atoms) IMFs Variable hardness, variable brittleness, not conductive, low melting points

Solutions 1. Explain the thermodynamics involved in making a solution (spontaneity, change in enthalpy). 2. Define a saturated solution, unsaturated solution, and supersaturated solution. 3. How does solubility of solids in solution change with temperature? Gases? 4. Describe what is meant by “like dissolves like.” 5. Define molality and the four colligative properties that it is used to define. 6. Explain how molality effects vapor pressure, boiling point, freezing point, and osmotic pressure. 7. Define isotonic, hypertonic and hypotonic. 8. Explain what the vant hoff factor is and what it would be for a non electrolyte, weak acid, and a strong electrolyte. ● Formation of a solution is an example of a spontaneous process ● Two criteria that favor (but do not guarantee) the spontaneous formation of a sol’n: 1. A decrease in the internal energy of the system (an exothermic change) 2. An increase in the disorder in the system (which indicates an increase in the entropy of the system) Solubility- extent to which a solute may be dissolved in water, or any solvent● For many solids dissolved in liquid water, the solubility increases with temperature. The increase in kinetic energy that comes with higher temperatures allows the solvent molecules to more effectively break apart the solute molecules that are held together by intermolecular attractions. Unsaturated- solute concentration is less than its solubility Saturated- solute concentration is equal to its solubility Supersaturated- solute concentration exceeds its solubility “Like dissolves like” → Polar dissolves polar, nonpolar substances dissolve nonpolar substances. Molality: mol solute/ kg solvent; do NOT get this confused with molaRity. Colligative Properties: properties of a solution that depends only on the CONCENTRATION of a solute species 1. Boiling point elevation- elevation of the boiling point of a liquid by addition of a solute 2. Freezing point depression- lowering of the freezing point of a liquid by addition of a solute 3. Osmotic Pressure- opposing pressure required to prevent bulk transfer of solvent molecules through a semipermable membrane 4. Vapor Pressure Depression- dissolving a nonvolatile substance in a volatile liquid results in a lowering of the liquid’s vapor pressure Volatile- easily evaporated at normal temperatures Nonvolatile- not easily evaporated at normal temperatures https://www.youtube.com/watch?v=WrIg-rUmtfk ^Colligative Properties Boiling Point Elevation: ​https://www.youtube.com/watch?v=Jo4ocC1r5-Y Freezing Point Depression: h​ ttps://www.youtube.com/watch?v=06Buf6N2Yp4 Vapor Pressure Lowering: ​https://www.youtube.com/watch?v=wum0RNuJalQ Osmotic Pressure: ​https://www.youtube.com/watch?v=8WoyxPN1sSo Boiling Point & Molality: The higher the concentration (molality), the higher the boiling point. Vapor Pressure Lowering: The higher the concentration (molality), the lower the vapor pressure. Freezing Point Depression: Directly proportional to the molal concentration; DeltaT​f​= K​f​m, where m is the molal concentration of the solute in the solvent, K​f​ is the freezing point depression constant. Osmotic Pressure: Directly proportional to molality● Isotonic- of equal osmotic pressure ● Hypotonic- of less osmotic pressure ● Hypertonic- of greater osmotic pressure What is the vant hoff factor? - The number of particles produced - Mol particles in solution/ mol solute dissolved https://www.youtube.com/watch?v=7o86KBssCOA Lewis Acid-Base 1. Define a lewis acid and base and provide one example. Lewis Acid- a compound that accepts an electron pair (BF​3​) Lewis Base- a compound that donates an electron pair (NH​3​) https://www.youtube.com/watch?v=aJQV5LqZJ6Q Electrochemistry 1. Define oxidation and reduction. 2. Outline the steps in the half reaction method in an acidic solution and a basic solution. 3. Outline the processes that occur at the anode and cathode. 4. Which electrode do the electrons flow towards? 5. Which electrode do cations and anions flow towards? 6. Explain how to determine which species will be reduced and which will be oxidized if given their reduction potentials. ● Oxidation- loss of electrons; Reduction- gain of electrons → To help you remember which is which, use the mnemonic LEO GER (Loss of Electrons- oxidation; Gain of Electrons reduction). Steps in the half reaction method in an acidic solution: https://www.youtube.com/watch?v=pESvgnP64ME Steps in the half reaction method in a basic solution:​https://www.youtube.com/watch?v=v5sDNmYCaqo *May want to watch this if you don’t understand how to determine oxidation states: https://www.youtube.com/watch?v=Sa1QQKmq8Ds ● Electrons flow from the anode to the cathode ● Oxidation occurs at the anode ● Reduction occurs at the cathode ● Anions in the salt bridge flow toward the anode and cations in the salt bridge flow toward the cathodeCalculating the cell potential at standard and non-standard conditions: https://www.youtube.com/watch?v=62rEVSBBJMw ● Galvanic cells have positive cell potentials. ● The reaction at the anode will be the half-reaction with the smaller (or more negative) standard reduction potential https://www.chem.wisc.edu/deptfiles/genchem/netorial/rottosen/tutorial/modules/electrochemistry/0 5potential/18_52.htm ● The standard reduction potential can be determined by subtracting the standard reduction potential for the reaction occurring at the anode from the standard reduction potential for the reaction occurring at the cathode. The minus sign is necessary because oxidation is the reverse of reduction. ● Reverse the potential that will give you the highest positive value. How to solve : http://www.chem.purdue.edu/gchelp/howtosolveit/Electrochem/Electrochemical_Cell_Potentials.ht m
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