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Barne's Chemistry 130 final exam notes (chapters 20 & 23)

by: Christina Bouchillon

Barne's Chemistry 130 final exam notes (chapters 20 & 23) Chem 130

Marketplace > University of Tennessee - Knoxville > Chemistry > Chem 130 > Barne s Chemistry 130 final exam notes chapters 20 23
Christina Bouchillon
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These notes cover the last two chapters that were covered in class that will be on the final. You will need to study all past study guides to cover all the material that will be tested on the final...
Chemistry 130
Christiane Barnes
Study Guide
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This 15 page Study Guide was uploaded by Christina Bouchillon on Monday May 2, 2016. The Study Guide belongs to Chem 130 at University of Tennessee - Knoxville taught by Christiane Barnes in Spring 2016. Since its upload, it has received 119 views. For similar materials see Chemistry 130 in Chemistry at University of Tennessee - Knoxville.


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Date Created: 05/02/16
1 Chem 130 chapter 20 & 23 study guide Electrochemistry  Electrochemistry is the study of redox reactions that produce or require an electric current.  The conversion between chemical energy and electrical energy is carried out in an electrochemical cell .  Spontaneous redox reactions take place in a voltaic cell . o Also known as a galvanic cell  Nonspontaneous redox reactions can be made to occur in an electrolytic cell by the addition of electrical energy. Voltaic cells • Electrical current: The amount of electric charge that passes a point in a given period of time (either as electrons flowing through a wire or as ions flowing through a solution) • Redox reactions involve the movement of electrons from one substance to another (redox reactions have the potential to generate an electric current). • A spontaneous redox reaction does not require external energy to proceed, which means that’s ΔG for the reaction is negative. • Voltaic (galvanic) cells produce an electrical current from spontaneous redox reactions. – In order to use that current, we must separate the place where oxidation is occurring from the place where reduction is occurring. Electrochemical cells • Oxidation and reduction half-reactions are kept as separate in half-cells in an electrochemical cell. • To constitute an electrical circuit: • Electron flow through a wire along with • Ions (electrolyte) flowing through a solution via the salt bridge . • The flow of electrons require a conductive electrode to allow the transfer of electrons either through:  An external circuit or  Metal or graphite electrode • An electrochemical cell requires the exchange of ions between the two half-cells of the system via a salt bridge. • Anode • Electrode where oxidation occurs • Anions attracted to it 2 • Connected to positive end of battery in an electrolytic cell • Loses weight in electrolytic cell • Cathode • Electrode where reduction occurs • Cations attracted to it • Connected to negative end of battery in an electrolytic cell • Gains weight in electrolytic cell • Electrode where plating takes place in electroplating Voltage and Current Voltage is the difference in potential energy between the reactants and products. It is also called the potential difference. Unit = volt • 1 V = 1 J of energy per coulomb of charge – The voltage needed to drive electrons through the external circuit – The amount of force pushing the electrons through the wire is called the electromotive force, emf. Current is the number of electrons that flow through the system per second. Unit = ampere • 1 A of current = 1 coulomb of charge flowing each second 18 1 A = 6.242 × 10 electrons per second • Electrode surface area dictates the number of electrons that can flow. – Larger batteries produce larger currents. Cell potential • The difference in potential energy between the anode and the cathode in a voltaic cell is called the cell potential . 3 • The cell potential depends on the relative ease with which the oxidizing agent is reduced at the cathode and the reducing agent is oxidized at the anode. • The cell potential under standard conditions is called the standard emf, E° . cell – 25 °C, 1 atm for gases, 1 M concentration of solution – Sum of the cell potentials for the half-reactions Cell notation • Shorthand description of a voltaic cell is written as follows: electrode | electrolyte || electrolyte | electrode – Oxidation half-cell on the left; reduction half-cell on the right – Single | = phase barrier • If multiple electrolytes in same phase, a comma is used rather than | • Often use an inert electrode – Double line || = salt bridge – Example: Zn(s) | Zn (aq) || Cu (aq) | Cu(s) – Cathode = Cu(s) 2+ – Cu ions are reduced at the cathode. – Anode = Zn(s) – The anode is oxidized to Zn 2+ions. Electrodes • Typically, – the anode is made of the metal that is oxidized; and – the cathode is made of the same metal as is produced by the reduction. • If the redox reaction we are running involves the oxidation or reduction of an ion to a different oxidation state, or the oxidation or reduction of a gas, we may use an inert electrode. – An inert electrode is one that does not participate in the reaction but just provides a surface for the transfer of electrons to take place on. Standard reduction potential • The absolute tendency of a half-reaction standard reduction potential cannot be measured. 4 – Only the potentials relative to another half-reaction can be measured. • To overcome this limitation, a standard half-reaction for the reduction of H to H is se2ected and assigned a potential difference of 0 V. – Standard hydrogen electrode, SHE • A redox reaction will be spontaneous when there is a strong tendency for the oxidizing agent to be reduced and the reducing agent to be oxidized. – Higher on the table of standard reduction potentials = stronger tendency for the reactant to be reduced – Lower on the table of standard reduction potentials = stronger tendency for the product to be oxidized Half cell potential • SHE reduction potential is defined to be exactly 0 V. – Standard reduction potentials compare the tendency for a particular reduction half-reaction to occur relative to the + reduction of H to H . 2 – Half-reactions with a stronger tendency toward reduction than the SHE have a positive value for E° . red • Half-reactions with a stronger tendency toward oxidation than the SHE have a negative value for E° . red • For an oxidation half-reaction, E° oxidation° reduction • A half-reaction with a strong tendency to occur has a large positive half-cell potential. • When two half-cells are connected, the electrons will flow so that the half-reaction with the stronger tendency will occur. • Under standard conditions, zinc has a stronger tendency to oxidize than copper. • Electrons flow from anode to cathode. • Therefore, the electrons flow from zinc, making zinc the anode. Zn → Zn 2++ 2 e − E° = +0.76 Cu → Cu 2++ 2 e − E° = −0.34 Calculating cell potentials under standard conditions • Cell potentials are intensive properties of matter. – Because cell potentials are intensive physical properties, when determining the cell potential, do not multiply the half-cell E° values, even if you need to multiply the half- reactions to balance the redox equation. o • The cell potential of an electrochemical cell (E cellis the difference between the electrode potential of the cathode and that of the anode. 5 • Cell potentials can be determined using the following equation: E° = E° – E° cell cathode (reduction) anode (oxidation) EX- Use the standard electrode potentials to calculate the standard cell potential for the reaction occurring in an electrochemical cell at 25 °C. Pb=reduction/cathode Cr=oxidation/anode 3 Pb (aq) + 2 Cr(s) → 3 Pb(s) + 2 Cr (aq) 3+ [Cr3+(aq) → Cr(s), E° = –0.73 V; Pb2+ (aq) → Pb(s), E° = –0.13 V] E° cell E° cathode-anode = -.13- (-.73) = .60 V Spontaneous process** Predicting spontaneity of redox reactions • A spontaneous reaction will take place when a reduction half- reaction is paired with an oxidation half-reaction lower on the table. – If paired the other way, the reverse reaction is spontaneous. – When E cells positive, the redox reaction of the cell is spontaneous (ΔG will be negative). ° – When E cells negative, the redox reaction of the cell is nonspontaneous (ΔG will be positive). 2+ 2+ • Cu(s) + Zn (aq) → Cu (aq) + Zn(s) nonspontaneous • Cu (aq) + 2 e → Cu(s) E° red= +0.34 V E°cell= E° cathode- anode • Zn (aq) + 2 e → Zn(s) E° = −0.76 V E° =.34 V- (-.76 red cell V)=1.10 V EX- Are these redox2+eactions sp2+taneous under standard conditions? Zn(s) + Ni (aq) → Zn (aq) + Ni(s) -.76 -.23 Zn= oxidation/anode Ni= reduction/cathode E° cell -.23- (-.76) = .53 V à spontaneous 2+ 2+ • Zn(s) + Ca (aq) → Zn (aq) + Ca(s) -.76 -2.76 Zn= oxidation/anode Ca= reduction/cathode E° cell -2.76- (-.76) = -2.00 V à nonspontaneous Predicting whether metal will dissolve an acid • Metals dissolve in acids – if the reduction of the metal ion is easier than the reduction of + H (aq); – if their ion reduction reaction lies below H reduction on the table. • Almost all metals will dissolve in HNO . 3 – Having N reduced rather than H 6 – Au and Pt dissolve in HNO + HC3. E° , ΔG°, and K cell • For a spontaneous reaction, one that proceeds in the forward direction with the chemicals in their standard states, – ΔG° < 1 (negative) – E° > 1 (positive) – K > 1 – ΔG° = −RT ln K = −nFE° cell – n = the number of electrons – F = Faraday’s constant = 96,485 C/mol e − EX- Use tabulated electrode potentials to calculate ΔG° for the reaction: 2 Na(s) + 2 H O(l) → H (g) + 2 OH (aq) + 2 Na (aq)+ 2 2 Spontaneous? Na= oxidation/anode H2O= reduction/cathode V= J/C E°cell -.83 V – (-2.71 V) = 1.88 V à spontaneous ΔG° = −nFE° cell= -2 mol e- * 96485 C/ mol e- * 1.88 J/C = -362784 J =3.6 *10^5 J ΔG° is negative à reaction is spontaneous EX- Use the tabulated electrode potentials to calculate K for the oxidation of iron by H at 25 °C: + 3+ 2 Fe(s) + 6 H (aq) → 2 Fe (aq) + 3H (g) 2 ΔG° = −nFE° cell ΔG° = −RT ln K E°cell= .0592/ n * log K Fe= oxidation/anode H= reduction/cathode E°cell= 0 – (-.036) = .036 V E°cell= (.0592/n) *log K n= # of mols of e- transerred (.036 V) * 6/ .0529 V = log K 3 K= 10 ^3.65 à 4.5*10 Cell Potential when Ion Concentrations Are Not 1 M • There is a relationship between the reaction quotient, Q; the equilibrium constant, K; and the free energy change, ΔG°. • Changing the concentrations of the reactants and products so that they are not 1 M will affect the standard free energy change, ΔG°. • Because ΔG° determines the cell potential, E , thcellltage for the cell will be different when the ion concentrations are not 1 M. Nernst equation ΔG = ΔG° + RT ln Q 7 (–) nFE cell (–) nFE° cell RT ln Q At 25 °C (T), – Faraday’s constant (F) = 96,500 C/mol e n = the number of electrons Converting from ln to log (2.303), the Nernst equation becomes Ecell=° cell .0592 V/ n (log Q) EX- Determine the cell potential of an electrolytic cell based on the following two half reactions: (-.23) Oxidation: Ni(s) → Ni (aq) + 2 e - (1.00Reduction:VO (aq, 0.010 M) + 2 H (aq) (aq, 1.0 M) + e → - 2 2+ VO (aq, 2.0 M) +H O(l)2 E° cellE° cell.0592V/n (logQ) E° = E – E = 1.00- (-.23) = 1.23 V cell cathode anode *Multiply by 2 bc 2 e-* = 2VO2^+ + 4H+ + 2e- à 2VO^2+ + H2O Q= [VO^2+]^2/[VO2^+]^2[H+]^4 = (2.00)^2/(.010)^2(1.0)^4 =40000 E° cell 1.23- .0592V/2 * log 40000 = 1.09V Ex- In an electrochemical cell Q = 0.0010 M and K = 0.10. What can you conclude about E cellnd E cell • (a) E is positive and E is negative cell cell • (b) Ecells negative and E cells positive • (c) both E and E are positive cell cell • (d) both E cellnd E cellre negative E° cell (.0592V/n)* (logK) Log(<1) is negative, E° cells negativeà non-spontaneous E cellE° cell (.0592V/n)* (logQ) E is positiveà spontaneous under non standard conditions cell Concentration cells • It is possible to get a spontaneous reaction when the oxidation and reduction reactions are the same, as long as the electrolyte concentrations are different. • The difference in energy is due to the entropic difference in the solutions. – The more concentrated solution has lower entropy than the less concentrated solution. – Electrons will flow from the electrode in the less concentrated solution to the electrode in the more concentrated solution. – Oxidation of the electrode in the less concentrated solution will increase the ion concentration in the solution; the less concentrated solution has the anode. 8 – Reduction of the solution ions at the electrode in the more concentrated solution reduces the ion concentration; the more concentrated solution has the cathode. • When the cell concentrations are equal, there is no difference in energy between the half-cells, and no electrons flow. • When the cell concentrations are different, electrons flow from the side with the less concentrated solution (anode) to the side with the more concentrated solution (cathode). Electrochemical cell overview • In all electrochemical cells, oxidation occurs at the anode and reduction occurs at the cathode. ° • In voltaic cells (spontaneous reactions; E cells positive), – the anode is the source of electrons and has a (−) charge; – the cathode draws electrons and has a (+) charge. ° – In electrolytic cells (non spontaneous reactions; E cells negative), – electrons are drawn off the anode, so there must be a place to release the electrons—the positive terminal of the battery; – electrons are forced toward the anode, so there must be a source of electrons—the negative terminal of the battery. Electrolysis • Electrolysis is the process of using electrical energy to break a compound apart. • Electrolysis is done in an electrolytic cell. • Electrolytic cells can be used to separate elements from their compounds. • In electrolysis we use electrical energy to overcome the energy barrier of a nonspontaneous reaction, allowing it to occur. 9 • The reaction that takes place is the opposite of the spontaneous process. 2 H 2g) + O (g2 → 2 H O(l)2spontaneous 2 H 2(l) → 2 H (g2 + O (g) 2 electrolysis • Some applications of electrolysis are the following: (1) Metal extraction from minerals and purification (2) Production of H f2r fuel cells (3) Metal plating Electrolytic cells • The electrical energy is supplied by a direct current power supply. – AC alternates the flow of electrons so the reaction won’t be able to proceed. • Some electrolysis reactions require more voltage than E cellredicts. This is called the overvoltage. Electrolysis of aqueous solutions • Possible cathode reactions: – Reduction of cation to metal – Reduction of water to H 2 • 2 H 2 + 2 e →H + 2 2H − E° = −0.83 V at standard conditions E° = −0.41 V at pH 7 • Possible anode reactions: – Oxidation of anion to element – Oxidation of H O2to O 2 • 2 H 2 → O + 2 e + 4H− + E° = −1.23 V at standard conditions E° = −0.82 V at pH 7 – Oxidation of electrode: • Particularly Cu • Graphite doesn’t oxidize • Half-reactions that lead to least negative E cellill occur. – Unless overvoltage changes the conditions Electrolysis of pure compounds • The compound must be in molten (liquid) state. • Electrodes are normally graphite. • Cations are reduced at the cathode to metal element. • Anions are oxidized at the anode to nonmetal element. Ex- Predict the half reactions occurring at the anode and the cathode for the electrolysis of aqueous Na SO 2 4 10 Anode = 2 H2O (l) à O2 (g) + 4 H+ + e- E.= 1.23V Cathode= 2 H2O + 2e- à H2 (g) + 2OH- E.= -.83 V E.cell= -.83 V – 1.23 V = -.206 V Anode= SO4 2- + 4H+ à H2SO3 + H2O + 2e- E.= .20 V Cathode= Na+ + e- à Na E.= -2.71 V E.cell= -2.71-.20 V= -2.91 V Electroplating • In electroplating, the work piece is the cathode. – Cations are reduced at cathode and plate to the surface of the work piece. – The anode is made of the plate metal. The anode oxidizes and replaces the metal cations in the solution. Mixtures of ions and electrolysis • When more than one cation is present, the cation that is easiest to reduce will be reduced first at the cathode. – Least negative or most positive E° red – When more than one anion is present, the anion that is easiest to oxidize will be oxidized first at the anode. – Least negative or most positive E° ox Stoichiometry of Electrolysis • In an electrolytic cell, the amount of product made is related to the number of electrons transferred. – Essentially, the electrons are a reactant. • The number of moles of electrons that flow through the electrolytic cell depends on the current and length of time. – 1 amp = 1 coulomb of charge/second − – 1 mole of e = 96,485 coulombs of charge • Faraday’s constant Conceptual plan: time (in seconds) → coulombs → moles of electrons → moles of metal → grams of metal Ex- Silver can be plated out of a solution containing Ag+ according to the half reaction: • Ag+(aq) + e- → Ag(s) • How much time in min does it take to plate 12 g silver using a current of 3.0 A? A= C/S 12.0 g Ag * 1 mol/ 107.87 g * 1 mol e-/ 1 mol Ag * 96485 C/1 mol * 1 sec/ 300 C * 1 min/ 60 sec = 59.6 minutes à 60 min/1 hr 11 Corrosion: Nondesirable Redox Reaction • Corrosion is the spontaneous oxidation of a metal by chemicals in the environment. – Mainly O 2 • Because many materials used are active metals, corrosion can be a very big problem. – Metals are often used for their strength and malleability, but these properties are lost when the metal corrodes. – For many metals, the product of corrosion does not adhere to the metal, and as it flakes off more metal can corrode. Reduction of O 2 • O 2s very easy to reduce in moist conditions. − − O 2g) + 2 H O2l) + 4 e → 2 OH (aq) E° = 0.40 V • O 2s even easier to reduce under acidic conditions. O 2g) + 4 H+ + 4 e → 2 H O(l) 2 E° = 1.23 V • Because the reduction of most metal ions lies below O on the 2 table of standard reduction potentials, the oxidation of those metals by O 2s spontaneous. Rusting • At the anodic regions, Fe(s) is oxidized to Fe . 2+ • The electrons travel through the metal to a cathodic region where O 2s reduced. – In acidic2+olution from gases dissolved in the moisture – The Fe ions migrate through the moisture to the cathodic region, where they are further oxidized to Fe ,which+ combines with the oxygen and water to form rust. – Rust is hydrated iron(III) oxide, Fe O ·2nH3O. 2 • The exact composition depends on the conditions. – Moisture must be present. • Water is a reactant. • It is required for ion flow between cathodic and anodic regions. • Electrolytes promote rusting. – They enhance current flow. – Acids promote rusting. – Lowering pH will lower E° redf O 2 Preventing corrosion 12 • One way to reduce or slow corrosion is to coat the metal surface to keep it from contacting corrosive chemicals in the environment. – Paint – Some metals, such as Al, form an oxide that strongly attaches to the metal surface, preventing the rest from corroding. – Another method to protect one metal is to attach it to a more reactive metal that is cheap. – Sacrificial electrode • Galvanized nails Sacrificial anode • If a metal more active than iron, such as magnesium or aluminum, is • in electrical contact with iron, the metal rather than the iron will be • oxidized. This principle underlies the use of sacrificial electrodes to • prevent the corrosion of iron. Transition metals • The properties of the transition metals are similar to each other. – And very different from the properties of the main group metals – High melting points, high densities, moderate to very hard, and very good electrical conductors • The similarities in properties come from similarities in valence electron configuration; they generally have two valence electrons. 2 x • For first and second t2ansxtion series, ns (n 2 1xd – First = [Ar]4s 3d ; second = [Kr]5s 4d • For third and fourth transition series, ns (n − 2)f (n − 1)d x • Some irregularities, some electron configurations must be found experimentally. Ex- Write ground state electron configurations of Nb 2+ and Os. Os= #76 =[Xe] 6s2 4f14 5d6 Nb 2+ = [Kr] 5s2 4d3 **lose highest occupied orbital =[Kr] 4d3 The atomic radius remains relatively constant because electrons are added to an (n-1) orbital Properties • Atomic size: of all transition metals are very similar, second and third row are very similar because of the added 14 f electrons (Lanthanide contraction) 13 • Ionization Energy: increases slowly across a series • Electronegativity: increases slowly across a series (exception is the last element) • Oxidation states: many transition metals exhibit multiple oxidation states, with the highest oxidation state being equal to the group number Complex ions • When a monatomic cation combines with multiple monatomic anions or neutral molecules it makes a complex ion. • The attached anions or neutral molecules are called ligands. • The charge on the complex ion can then be positive, negative or neutral, depending on the numbers and types of ligands attached. Coordination compounds • When a complex ion combines with counterions to make a neutral compound, it is called a coordination compound. • The primary valence is the oxidation number of the metal. • The secondary valence is the number of ligands bonded to the metal. – coordination number – Coordination numbers range from 2 to 12, with the most common being 6 and 4. CoCl 3 6H O2= [Co(H O) 2Cl6 3 Oxidation numbers • Knowing the charge on a complex ion and the charge on each ligand, one can determine the oxidation number for the metal. • Or knowing the oxidation number on the metal and the charges on the ligands, one can calculate the charge on the complex ion Inside bracket= ligand Outside bracket= complex ion EX- What is the charge of the complex formed by a platinum(II) metal ion surrounded by two ammonia molecules and two bromide ions? H- H- à N à Pt2+ à BR, Br, NH3 H- +2 +2(0) +2(-1)= 0 Complex ion formation • Complex ion formation is a type of Lewis acid–base reaction. 14 • A bond that forms when the pair of electrons is donated by one atom is called a coordinate covalent bond. Polydentate ligands • Some ligands have 2 or more donor atoms • These are called polydentate ligands or chelating agents • Ethylene diamine, NH CH 2H N2 , 2fte2 referred to as “en” has 2 donor atoms (both N atoms) – We call ethylene diamine a bidentate ligand • Ethylene diamine tetraacetate, abbreviated EDTA, has 6 donor atoms, it is hexadentate. • Chelating agents generally form more stable complexes than do monodentate ligands. • They can render metal ions inactive without actually removing them from solution. Coordination number • The atom of the ligand that supplies the nonbonding electrons for the metal-ligand bond is the donor atom • The number of these atoms is the coordination number • Some metals such as chromium(III) and cobalt(III) consistently have the same coordination number (6 in case of these two metals • The most commonly encountered numbers are 4 and 6 Geometries • There are two common geometries for metals with a coordination number of four: – __square planar____ – __tetrahedral_____ • When the coordination number is six, the geometry most commonly encountered is octahedral Isomers • _structural_ isomers are molecules that have the same number and type of atoms, but they are attached in a different order. • __stereoisomers are molecules that have the same number and type of atoms, and that are attached in the same order, but the atoms or groups of atoms point in a different spatial direction. • __linkage__ isomers are structural isomers that have ligands attached to the central cation through different ends of the ligand structure. • __geometric isomers are stereoisomers that differ in the spatial orientation of ligands. 15 • Cis–trans isomerism in square-planar complexes MA B 2 2 • In cin-trans isomerism, two identical ligands are either adjacent to each other (cis) or opposite to each other (trans) in the structure. • Cis–trans isomerism in octahedral complexes MA B 4 2 • __optical___ isomers are stereoisomers that are nonsuperimposable mirror images of each other. • Just like a right hand won’t fit into a left glove, two enantiomers cannot be superimposed on each other Enatiomers • A molecule or ion that exists as a pair of enantiomers is said to be chiral • The physical properties of chiral molecules are the same except in instances where the spatial placement of an atom matters • They will rotate plane-polarized light into opposite directions


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