Study guide 4
Popular in General Chemistry 2
verified elite notetaker
Popular in Chemistry
This 11 page Study Guide was uploaded by Cassidy Zirko on Sunday April 24, 2016. The Study Guide belongs to Chem 143 at University of Montana taught by Dr. Cracolice in Spring 2016. Since its upload, it has received 266 views. For similar materials see General Chemistry 2 in Chemistry at University of Montana.
Reviews for Study guide 4
Report this Material
What is Karma?
Karma is the currency of StudySoup.
You can buy or earn more Karma at anytime and redeem it for class notes, study guides, flashcards, and more!
Date Created: 04/24/16
Study Guide 4 Chapter 66: Spontaneous process changes that are chemically or physically possible under specified conditions without an external energy source o occurs in a process that follows a logical time line (ie. Ice liquid) o depends on temperature, manipulates direction of spontaneity 2nd law of thermodynamics in a spontaneous process, energy changes from being localized to being spread out Entropy, S measure of how much energy is spread out and how widely spread out it becomes at a specified temperature o heat cannot be changed into an equivalent amount of work energy o some heat will almost always leave the system in a reaction Reversible change: 1) system must be at equilibrium with its surroundings, temperature and pressure must be closely balanced o 2) direction of reversible reaction change can be changed with a change in temperature or pressure q ∆S= rev q revH phasechange Entropy at a constant temperature: T but ∆ H phasechange o ∆ S= T (think about a phase change diagram) Entropy increases as a substance changes from solid Liquid, Liquid gas or solid gas o More ways to organize the energy in a liquid or gas Endothermic absorbing energy from surrounding: +∆H∧+∆E Exothermic releasing energy to the surroundings:−∆H∧−∆E Enthalpy, H the heat content of a system, at constant pressure, enthalpy is the heat transferred st 1 Law of thermodynamics the internal energy of a system is equal to the sum of the heat flow in or out of the system and the work done on or by the system ∆ E=q+w o o w work, w=−P∆V , an expanding system does work on the surroundings: w, a condensing system has work done on it by the surroundings: +w L∗.10133kJ o w=atm∙ 1atm∙ L o q= heat flow, q=mc∆T (use this one when there is a change in temp) or q=m∗∆ H fusion∨vaporiz tion o system gains heat: +q, system loses heat: q ∆ E the sum of total energy transferred between the system and surroundings o Energy passes from surroundings to system: +∆ E o Energy passes from system to surroundings: −∆ E ∆ H=∆E+P ∆V , because ∆ E=q+(−P∆V) ∆ H=q+ −P∆V +P∆V=q p where q pis the heat flow at a constant ressure ∆ H =q only when 1) the changes occur at a constant pressure and 2) the only work p being done on the system is pressurevolume work Volume changes are only significant in value when gases are invoved ∆ H=∆E+RT ∆n(gases) Chapter 67: Translational energy the energy from a particle moving around in its container, the energy of random motion, the energy levels are extremely close together Rotational energy the energy of a particle from rotating or tumbling through space, the energy levels are slightly farther apart than the translational energy Vibrational energy energy from the particles of a molecule vibrating within their bonds, greater spaces between the energy levels Electronic energy energy from the interactions (repulsions) of subatomic particles, the energy from the interactions of the electrons with the nuclei, huge spaces between the energy levels, so much so that most particles reside at ground state at room temperature Energy is always quantized, divided up in to defined energy levels Energy at the particulate level depends on the number of particles in each quantized energy level W the total number of quantum states that corresponds to the macroscopic energy of the system W final Boltzmann Equation: S=klnW ∆ S=kln this equation expresses (W)initial entropy at a particulate level rather than a macroscopic level. It requires us to know the amount of possible energy combinations that an atom can have, which is extremely extensive. The larger the amount of energy states (The higher W value) the higher the entropy A large positive entropy value (∆S ) value for a reaction means that it is a statistically favored process The second law of thermodynamics governs the direction that is favored in a physical or chemical change An increasing entropy, S, is an increase in molecular level energy dispersion and vice versa o Entropy increases when a solid/liquid is changed to a gas gas has higher motional energy and more spaces between particles allowing for more ways to disperse the energy of the particles o Entropy decreases when a gas is dissolved in water particles have less space between each other which limits the freedom of energy o Entropy increases with increasing mass heavier mass means closer spaced energy levels because of the electron behavior, heavier particles have more energy levels compared to a lighter molecule o Entropy increases with more particles in an otherwise similar compound more particles means more ways that the compound can move ad organize its energy o Entropy increase when a solid/liquid is dissolved in water energy levels become more spread out with mixing o Entropy increases as temperature increases with an increase in temperature, translational energy changes from being distributed over a small area to being distributed over a large area Third Law of Thermodynamics the entropy, S, of a pure crystalline solid of a pure substance is zero at an absolute temperature of zero Entropy is usually independent of temperature (Except at low temperature) but is pressure dependent V Entropy for changing volume: ∆ S=Rln( 2) , ∆ S>0 V 1 The spontaneous irreversible constant temperature expansion of an ideal gas is equal to the increase in entropy of the system and surroundings o ∆ S>0 for the system and surroundings, T ∆ S>q irreversible Reversible constant temperature expansion of an ideal gas results in no change in entropy of its system and surroundings ∆ S=0 T ∆ S=q o for system and surroundings, reversible Entropy is a state function, it only depends on the final and initial values of the system Chapter 68 Driving forces of Spontaneous changes: 1) minimization of heat energy, 2) maximization of entropy Gibbs Free Energy: net effect of those 2 tendencies exposed in a thermodynamic state functions, relates temperature, entropy and enthalpy Reaction Potential: the progress of the reaction in relation to the amount of products that are being created Spontaneous direction is always downward or negative o Ex. spontaneous reaction like a ball rolling down a hill no extra energy is needed to occur o Non spontaneous reaction like a ball rolling up the hill extra energy is needs for reaction to occur Extent of reaction: represents amount of products formed o Beginning of reaction, no product is formed, End of reaction only products are formed o When amount of products and reactants are equal, system is at equilibrium, Gibbs free energy is at a minimum The reaction chemical potential is zero at the point of minimum Gibbs free energy ∆ G=0 o At equilibrium, , there are no chemical forces being done on the reaction the spontaneous reactions direction is toward equilibrium concentrations when extent of reaction is less than the equilibrium extent of reaction o –slope=positive forces on the system, pushing reaction toward products at equilibrium ∆ G<0 o , forces >0, spontaneous reaction in the forward direction At points to the right of equilibrium, the reaction will have negative chemical potential with the spontaneous direction being towards the formation of more reactants o Positive slope=negative forces on system, pushing reaction towards reactants at equilibrium o ∆ G>0 , Forces <0, spontaneous reaction in the reverse direction qrev ∆ S= ∆H ∆ H=q ∆ S= T , at a constant pressure, T because rev this focuses on the system ∆ S = −∆ H surroundingsT because the entropy of the surroundings must be equal but opposite of the entropy of the system Gibbs Free energy Equation for reactants and products in standard states: ∆ G =∆ H −T ∆ S ° r r r , only valid at a constant pressure Gibbs Free energy of formation is the change in free energies for the reaction in which n∗ΔG° ¿ f reactants pure stable elements react to form 1 mole of product: n∗ΔG° ¿ f products ∆rG°=Σ¿ ° Nonidealized free energy reaction: ∆ r=∆ G +RrlnQ where RTlnQ is the adjustment factor for a reaction that the species are not at one molar concentration ° ∆ r>∆ G r Increasing temperature and vice versa ° Because at equilibrium ∆ r =0 , at equilibrium you can replace Q with K p 0=∆ G=∆ G +RTln K ° ∆ G=−RTlnK o r r p r p ∆n(gase) RT ¿ K =ppressure equilibrium constant: K pK ¿ c Small K /Kp facors the reverse reaction and the reactants side Large K /Kp facors the forward reaction and the products side ∆ G K /p can be used to confirm r and predict the correctly favorable direction Remember that both Q and K are the concentrations of the products raise to their respective coefficients (if applicable) over the concentrations of the reactants raise the their respective coefficients (if applicable) and that solid/liquids/gases ARE NOT included in K or Q ∆ G ∆ G =∆ H −T ∆ S ° To predict the sign/value of r refer back to r r r ° ° ∆ r ∧∆ S r o The sign will vary based on the relationship of ° ° ° o If ∆ r ∧∆ S r are opposite signs, then of ∆ r is temperature independent and spontaneous at all temperatures ° ° What varies is the DIRECTION of spontaneity: + ∆ r ∧−∆ S r ° causes + ∆ r and the reaction is spontaneous in the REVERSE direction ∆ r ∧+∆ S r ° ∆ r ° causes and the reaction is spontaneous in the FORDWARD direction ∆ H ∧∆ S ° ∆ G ° o If r r have the same signs, then of r is temperature dependent and spontaneity depends on the temperature ° ° ° If both ∆ r ∧∆ S r are positive at high temperatures ∆ r will be negative and spontaneous in the FORWARD direction ° ° ° If both ∆ r ∧∆ S r are positive at low temperatures ∆ r will be positive and spontaneous in the REVERSE direction ∆ H ∧∆ S ° ∆ G ° If both r r are negative at high temperatures r will be positive and spontaneous in the REVERSE direction ∆ H ∧∆ S ° ∆ G ° If both r r are negative at low temperatures r will be negative and spontaneous in the FORWARD direction Chapter 69 Voltaic Cell a system that spontaneously generates a flow of electrons as a result of a chemical change Salt bridge a solution of electrolytes that are not involved in the net chemical change Oxidation loss of electrons and increase in oxidation number Reduction gain in electrons and decrease in oxidation number Half reactions the 2 halves of a complete electron transfer reaction which shows the path of electrons and how they are transferred Both oxidation and reduction reactions must occur because you cant have free floating electrons, they need to have somewhere to go The number of electrons produced by the oxidation reaction must all be used in the reduction reaction Electron Bookkeeping: Assigning Oxidation Numbers 1) Oxidation number of any elemental substance is zero 2) Oxidation number of a monatomic ion is the same as the charge of the ion 3) Oxidation number of oxygen that is in a molecule is 2, unless in peroxides ( 1) or OF 2(+2) 4) Oxidation number of combined hydrogen is +1, except when used as a monatomic hydride ion, H (1) 5) In any molecular or ionic species, the sum of oxidation numbers of all the atoms present in any formula will equal the overall charge on the unit One element MUST decrease in oxidation number and another element MUST increase in oxidation number Hints and Tips: o An element in its elemental state must change oxidation number, element on one side has an oxidation number of zero so it will have an oxidation number that is not zero on the other side o Not in elemental form, oxygen=2 and hydrogen= +1, these usually don’t change unless on one side of the equation they are in elemental form o Group 1A/1 and 2A/2 elements has 1 oxidation state other than zero, it does not change unless in elemental form Volt the potential difference between 1 joule of energy being required or released in moving one coulomb of charge from one point to another in a circuit, V=J/C Voltage electromotive force, emf Anode location in battery where oxidation occurs, has a negative charge Cathode location in battery where reduction occurs, has a positive charge Oxidizing agent species that accepts electrons and is reduced in the process Reducing agent species that donates electrons and is oxidized in the process Chapter 70 1. Balance Oxidation numbers a. Identify the elements that are being oxidized and reduced base on the equation. Then write half reactions for the oxidation and reduction reactions b. Balance element that is being oxidized or reduced. Make sure there are the same number on each side before you do anything with oxidation numbers c. Balance the elements other than hydrogen or oxygen. These elements are the ones that don’t have a variation on oxidation number. This keeps the equation balanced d. Determine how the oxidation number changes. This tell you how many electrons are being produced in the half reaction. Remember if you have multiple of the element that is being oxidized or reduced that will affect the number of electrons that are being added to or produced e. Add the electrons so that the net oxidation number is the same on both sides of the equation 2. Balance Charges a. Add hydrogen ions (H ) when the reaction is occurring in an acidic solution. These will be added to the oxidation reaction and you are looking to increase the charge to make it equal. Ex. if the charge on one side is 2 and the charge on the other side is +5 then you will want to add 7H to the side of the equation with 2. This will make both sides of the equation have the same net charge. Also, when you are balancing your oxygens, the number of hydrogens will come based on the number of waters are needed in the equation b. Add hydroxide ions (OH) when the reaction occurs in a basic solution. This is the same procedure as 2a but you are working to making the net charge of the reaction more negative 3. Balance oxygen and hydrogens a. Add water molecules to balance the oxygens. This works because water has a net charge of zero so it doesn’t manipulate all the work you just did to make the charges balanced. When accounting for the number of oxygens, be sure to account for oxygens that are in other compounds in the reaction b. Check the number of hydrogens on each side. If you have done all the previous steps right then your amount of hydogens should already be equal 4. Check a. Final step is to check both the atom and the charge balance, both should be equal on both sides of the equation Chapter 71 Strong oxidizing agent strong attraction for electrons Weak oxidizing agent attracts electrons only slightly Strong reducing agent releases electrons readily Weak reducing agent holds on to its electrons Flourine is the best oxidizing agent meaning that it is easily reduced The activity series can be used to help with comparing oxidizing and reducing ability o The first element in the series, Lithium, is the best reducing agent so it is the most easily oxidized o The last element in the series is the best oxidizing agent so it is the most easily reduced The table of electron potential can also be used to determine the relative strengths of oxidizing and reducing agents o If the reduction potential is negative then the reaction would rather be preformed in the reverse direction causing the species to be oxidized (acting as a reducing agent) o If the reduction potential is positive then the reaction would like to occur the way it was written in the forward direction causing the species to be reduced (acting as a oxidizing agent) o These conditions are true because the table is of reduction potentials so a high reduction potential means that the species wants to be reduced In aqueous redox reaction, a strong oxidizing agent takes electrons from a strong reducing agent creating weaker oxidizing and reducing agent as the products The reaction favored in direction of the weaker oxidizing and reducing agent Acids have an ability to release hydrogen gas on a reaction with certain metals metals that will release hydrogen gas are reducers that are weaker than hydrogen Voltage/potential depends on the temperature and ion activity ℃ Standard state condition 25 , pure elemental electrodes, if gas at 1 bar pressure, concentrations at 1.00 Molar Standard electrode potential the voltage produced when the anode and cathode are in a system of the standard states Standard hydrogen electrode the staondard at which standard state potentials are calculated relative to Standard reduction potential the reduction potential between standard hydrogent electrode and the second electrode If the reduction potential is positive electrode functions as the cathode (location of the reduction half reaction), occurs spontaneously If the reduction potential is negative electrode functions as the anode (location of the oxidization half reaction), occurs spontaneously in the reverse direction Sometimes you have to think about the relative strength of both electrodes if both are positive/negative. The values that are closer to zero between the two electrodes will have the opposite reaction that is written in the reduction potentials table. This causes the reduction potential sign to switch Voltages are a measurement per electron you do not multiply the voltages when multiplying the half reactions, the reduction potential will remain constant for any variation on the half reaction Electron Transfer vs. Proton Transfer o Acid/base reactions are proton transfer, redox reaction is transfer of electrons o Special names: acidproton donor, baseproton acceptor, reducing agent electron donor, oxidizing agent electron acceptor o Behavior: species behaving as proton donor (acid) in one reaction and a proton acceptor (Base) in another, a species behaving as an electron donor (reducing agent) can also act as an electron acceptor (oxidizing agent) in another reaction o Classification: acids/bases classified base on ability to donate or accept protons, strengths of oxidizing or reducing agents classified based on ability to attract or release electrons o Equilibrium: favored side of acid/base equilibrium can be predicted based on strength, favored side of redox equilibrium can be determined by relative strength Chapter 72 9.65∗10 C(J) w eleE° where n=moles of electrons, F faraday constant moleof electrons(V) and E standard reduction potential in volts w maxw =ele° The maximum work is the negative free energy charge so ∆ G°=−nFE° r Free energy is in kJ/mol so you will need to convert from joules Remember that a spontaneous reaction in the forward direction has a negative free energy and spontaneous in the reverse direction has a positive free energy To find n its simply the number of electrons that are transferred in a redox reaction If don’t know how many electrons are being transferred based on the net reaction, write out the half reactions ∆ G=−RTlnK+RTlnQ Remember: RT RT Nernest equation: E= lnK− lnQ general equation nF nF E= .0257 lnQ For reactions at standard conditions n The reduction potential predicts spontaneous forward reaction if it is positive, a large K value shows that the reactants will almost completely turn into products The reduction potential predicts a spontaneous reverse reaction if it is negative, a small K value show shows the reverse reaction is favored The equilibrium constant can be anything from solubility, to acid/base to just regular concentration equilibrium E=E°− .0257 lnQ n is only for reactions at nonstandard conditions because it is taking the reduction potential at standard conditions and allowing an adjustment based .0257 on the concentration from n lnQ Remember that Q and K only have aqueous species in solution −∆G ,+E Forward change is spontaneous: Q<K, Q=K ,∆G=0,E=0 At equilibrium: Reverse change is spontaneous: Q>K ,+∆G ,−E Equilibrium, gibbs free energy and standard electron potential all have a relationship to one another E=E°− .0257lnQ Sometime you will need to do n for individual half reactions if all the concentrations of each of the ions vary from 1 molar Chapter 73: Primary Batteries based on reactions that are difficult to reverse, it is dead and done Secondary batteries based on reactions that are easily reversed, rechargeable How does an Alkaline Battery Work? o Electrolytes is potassium hydroxide creating a basic solution How does a Calculator or Watch Battery Work? o Button battery alkaline battery, being small has advantages o Anode has oxidation of powdered zinc o Electrons reduces silver oxide rather than the magnesium oxide How Does an Automobile Battery work? o Needs to produce a very large amount of energy and needs to be rechargeable o Most cost efficient – 12V lead acid battery o Has 6 cells, lead is oxidized, electrons reduce lead(IV) at cathode How Does a Lithium Ion Battery Work? o Cell phones, laptops computers, digital cameras o High energy to weight ratio o Can be recharged while holding a full charge o Anode is a form of carbon, cathode a metal ion o Electrolyte is lithium ion How Does a Fuel Cell Work? o Fuel Cell electrochemical cell that converts energy of a fuel into electricity o Both fuel cell and oxidizing agent are continuously supplied to cell o Hydrogen is fuel and oxygen is oxidizing agent o Key point: proton exchange membrane, thin, semipermeable membrane that allows passage of hydrogen ions but not electrons ( proton exchange membrane) o Hydrogen gas moves into cell on left, into anode o Electrons pass through electric current powers electronic devices o Major advantages: high energy efficiency, lack greenhouse gas emissions o Hydrogen acts as an energy storage, not an energy source 73.2 Why do Metals Corrode and How can it be Prevented? Corrosion oxidization of a metal by a substance in the environment producing an unwanted product Electron transfer reaction forming rust varies with environment How can Corrosion be Prevented? o Corrosion protection can be protective coatings o Pain prevents direct contact between steel o Galvanizing practice that adds zinc coating to steel or iron o Zinc is more reactive iron acts as anode in electrochemical cell 73.3 How Does an Electrolyte Cell Operate? Electrolyte ionic solution in where electrode are immersed into Electrodes are connected by wires Electronic current flows through metallic parts of battery Anode has oxidation, has a negative charge Cathode has reduction, has positive charge Basic units of energy, ampere, volt, ohm, watt, coulomb, joule Coulomb quantity of electric charge, symbolize by the letter C, Ampere is rate of flow of charge, measured in coulombs/sec When multiplied by time it changes just to coulombs Faraday, F quantity of charge current by one mole of electrons How Does Quantity of Electric Charge Relate to the Mass of Metal Deposited in an Electrolytic Cell? o Faraday’s Law of Electrolysis quantity (mass) of an elemental substance released or deposited in electrolysis is proportional to the quantity of electrical charge that has passed through the system How Is Mass of Metal Deposited for an Electrolytic Process? o Half reactions can be used in stoichiometry problems o Quantity of charge amperes*seconds converted to moles of electrons with 9.65*10 A*S
Are you sure you want to buy this material for
You're already Subscribed!
Looks like you've already subscribed to StudySoup, you won't need to purchase another subscription to get this material. To access this material simply click 'View Full Document'