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Organic Chemistry Chapter 22 What is organic chemistry? The study of carbon compounds Carbon can form stable bonds with itself so it can make long chains and ring structures Carbon atoms must make four bonds, oxygen must make two bonds, nitrogen must make three bonds and hydrogen must make one bond in organic compounds Characteristics of Organic Compounds Molecular formula – showing the actual number of atoms in each element in that compound Example – C2H4O2 for ethanoic acid Structural formula – showing how the atoms are arranged Example CH3CO2H for ethanoic acid meaning 3 H atoms are attached to 1 C atom, which is attached to another C which has 2 O atoms attached to it, one of which has an H attached to it Characteristics of Organic Compounds Graphical formula – showing how these atoms are arranged in space and the bonds between them. Lines represent covalent bonds (shared electrons) between the atoms Example – on board for ethanoic acid Characteristics of Organic Compounds Skeletal formula – showing an abbreviated form of the carbon chain, with each line segment understood to have a carbon atom at each end. May or may not show terminal carbons Example – on board for ethanoic acid Name – based on IUPAC nomenclature Example - CH3CO2H is called ethanoic acid Alkanes General formula – CnH2n+2 Examples Methane Butane Naming 1. Name the longest chain first Prefixes 1 C - meth 5 C – pent 6 C - hex 2 C - eth 7 C - hept 3 C - prop 8 C -oct 4 C – but ** If there are two longest chains of the same length give priority to the one with the most branching Alkanes Unreactive because the C-C bonds and C-H bonds are relatively strong Most common reaction is combustion (exothermic) Carbon compound + oxygen carbon dioxide + water Naming 2. Identify the type of bonding that is present Single bonds – -ane Double bonds – -ene Triple bonds – -yne Naming 3. Identify any functional groups Alkane – -e Alcohol – -ol Halogenalkane – chloro-, bromo- or iodo Aldehyde – -al Ketone – -one Naming 3. Identify any functional groups Carboxylic acid – -oic acid Ester – -oate Amide – -amide Amine – -amine Naming 4. Use numbers to indicate where groups or bonds are in a chain ** A side note – cycloalkanes Alkanes with ring-like structures Side chains are named by replacing -ane with –yl Naming **One more side note – aromatic compounds Called arenas Based off of benzene rings Use the smallest numbers possible Nomenclature Practice Name this compound CH3 H3C CH3 Cl H3C Nomenclature Practice Step #1: For a branched hydrocarbon, the longest continuous chain of carbon atoms gives the root name for the hydrocarbon Step #2: When alkane groups appear as substituents, they are named by dropping the -ane and adding -yl. Step #3: The positions of substituent groups are specified by numbering the longest chain of carbon atoms sequentially, starting at the end closest to the branching. Step #4: The location and name of each substituent are followed by the root alkane name. The substituents are listed in alphabetical order (irrespective of any prefix), and the prefixes di-, tri-, etc. are used to indicate multiple identical substituents. Alkenes and Alkynes Alkenes – have at least one double bond Alkynes – have at least one triple bond Both are said to be unsaturated Aromatic Hydrocarbons Cyclic unsaturated hydrocarbons with delocalized electrons The simplest aromatic hydrocarbon is benzene (C6H6) OR… Geometric Isomerism in Aromatics ortho (o-) = two adjacent substituents o-dichlorobenzene meta (m-) = one carbon between substituents m-dichlorobenzene para (p-) = two carbons between substituents p-dichlorobenzene Alcohols General formula – R-OH Naming Change ending of the main chain to -ol Amines General formula – RNH2 Naming Add amino- to the beginning or –amine to the ending of the name Aldehydes General formula Naming O R–C-H Change the ending of the main chain to -al Ketones General formula Naming O R–C-R Change the ending of the main chain to -one Carboxylic acids General formula O R – C - OH Naming Change the ending of the main chain to –oic acid Esters General formula Naming O R – C – O – R’ Name the R’ chain first Name the R chain and change the ending to -oate Amides General formula Naming O R – C – NH2 Change ending to -amide Ethers General formula R – O – R’ Naming Name both R groups as side chains and add ether to the end Isomers Structural isomers – have the same molecular formula but a different structural formula; similar chemical properties Example – pentane Functional group isomers – same molecular formula but different functional groups; different chemical and physical properties Examples: Carboxylic acid and ester Aldehyde and ketone Electrochemistry Chapter 17 Oxidation and Reduction (Redox) Electrochemistry deals with nonspontaneous Redox reactions Electrons are transferred Spontaneous redox rxns can transfer energy Electrons (electricity) Heat Non-spontaneous redox rxns can be made to happen with electricity Electrochemistry Terminology Oxidation – A process in which an element attains a more positive oxidation state Na(s) Na+ + eReduction – A process in which an element attains a more negative oxidation state Cl2 + 2e- 2ClLEO says GIR Loss of electrons = oxidation Gain of electrons = reduction Electrochemistry Terminology Oxidizing agent The substance that is reduced is the oxidizing agent Reducing agent The substance that is oxidized is the reducing agent Redox reactions 0 0 +1 -1 2Na + Cl2 2NaCl Each sodium atom loses one electron: 0 +1 Na Na+ + esodium is oxidized Each chlorine atom gains one electron: 0 -1 Cl + e- ClChlorine is reduced Trends in Oxidation and Reduction Active metals: Lose electrons easily Are easily oxidized Are strong reducing agents Active nonmetals: Gain electrons easily Are easily reduced Are strong oxidizing agents Balancing a Redox reaction in Acidic Solutions Write separate equations for the oxidation and reduction half-reactions. For each half reaction 1. 2. a. b. c. d. 3. 4. Balance all the elements except H and O Balance O by using H2O Balance H by using H+ Balance the charge using e- If necessary, multiply one or both balanced half reactions to equalize the e- transferred. Add the half reactions canceling identical species. Examples – Acid Solution MnO4-(aq) + Fe2+(aq) Fe3+(aq) + Mn2+(aq) 5e- + 8H+(aq) + MnO4-(aq) Mn2+(aq) + 4H2O(l) 5Fe2+(aq) 5Fe3+(aq) + 5e- 5Fe2+(aq) + 8H+(aq) + MnO4-(aq) Mn2+(aq) + 5Fe3+(aq) + 4H2O(l) H+(aq) + Cr2O72-(aq) + C2H5OH(aq) Cr3+(aq) + CO2(g) + H2O(l) 4H2O(l) + C2H5OH(aq) 2CO2(g) + H2O(l) + 12H+(aq) + 12e12e- + 28H+(aq) + 2Cr2O72-(aq) 4Cr3+(aq) + 14H2O(l) 16H+(aq) + C2H5OH(aq) + 2Cr2O72-(aq) 2CO2(g) + 4Cr3+(aq) + 11H2O(l) Balancing a Redox Reaction in Basic Solution 1. 2. 3. Use the half reaction method like we did for acidic solutions To both sides of the equation obtained in #1, add a number of OH- ions that is equal to the number of H+ ions Form H2O on the side containing both H+ and OH- ions, and eliminate the number of H2O molecules that appear on both sides of the equation Example – Basic Solution Ag(s) + CN-(aq) + O2(g) Ag(CN)2-(aq) 4Ag(s) + 8CN-(aq) 4Ag(CN)2-(aq) + 4e4e- + 4H+ + O2(g) 2H2O(l) 4H2O(l) + O2(g) + 4Ag(s) + 8CN-(aq) 4Ag(CN)2-(aq) + 2H2O(l) + 4OH2H2O(l) + O2(g) + 4Ag(s) + 8CN-(aq) 4Ag(CN)2-(aq) + 4OH- Cu + Cr2O7-2 Cu+2 + Cr+3 3Cu 3Cu+2 + 6e6e- + 14H+ + Cr2O7-2 2Cr+3 + 7H2O 14H+ + 3Cu + Cr2O7-2 2Cr+3 + 7H2O + 3Cu+2 14H2O + 3Cu + Cr2O7-2 2Cr+3 + 7H2O + 3Cu+2 + 14OH - Electrode Potentials and Half Cells When a metal comes into contact with a solution containing its own ions, an equilibrium is set up Some reactive metals will lose electrons easily – equilibrium will lie to the left A less reactive metal will show less tendency to ionize – equilibrium will lie to the right Whenever an element is placed with a solution containing its own ions, an electrical charge will develop on the metal, or in the case of non-metals on the inert conductor placed in solution The charge is an electrode potential Mx+(aq) + xe- M(s) Size and sign of the potential depends on the ability to lose or gain electrons The system is a half cell Electrochemical Activity Series Elements at the top of the series gain electrons most readily, have more positive Eo values and are the best oxidizing agents Elements at the bottom lose electrons more readily, have more negative Eo values and are the best reducing agents The more positive the Eo1/2, the more it tends to occur. Electrochemical Cells An apparatus for generating electrical energy from a spontaneous REDOX reaction Connect two half cells with different electrode potentials A salt bridge connects the two halves Allows the transfer of ions between the two solutions and completes the circuit Electrons flow through the wire toward the more positive half cell Anode: The electrode where oxidation occurs Cathode: The electrode where reduction occurs Red Cat: Reduction at the Cathode Line Notation An abbreviated representation of an electrochemical cell Zn(s) | Zn2+(aq)|| Cu2+(aq) | Cu(s) Anode Anode Cathode Cathode material solution solution material Standard Electrode Potential Eo is the variable Measured relative to a standard hydrogen electrode (Eo = 0.00V) Under standard conditions: 25oC, 1atm and 1M solutions For the Example: From a table of reduction potentials: Zn2+ + 2e- Zn E = -0.76V Cu2+ + 2e- Cu E = +0.34V Zn-Cu Galvanic Cell The less positive, or more negative reduction potential becomes the oxidation… Zn Zn2+ + 2eCu2+ + 2e- Cu Zn + Cu2+ Zn2+ + Cu E = +0.76V E = +0.34V E0 = + 1.10 V Galvanic (Electrochemical) Cells Spontaneous redox processes have: A positive cell potential, E0 A negative free energy change, (-G) A nonspontaneous redox process have: A negative cell potential A positive free energy change, (+G) Calculating G0 for a Cell G0 = -nFE0 n = moles of electrons in balanced redox equation F = Faraday constant = 96,485 coulombs/mol e1 V = 1 J/coulomb Zn + Cu2+ Zn2+ + Cu E0 = + 1.10 V G0 = -(2mol e-)(96485 coulombs)(1.10 J mol e- = -212267 J = -212 kJ coulomb ) Gibb’s Free Energy and Ecell K Eo Go Conclusion >1 Positive Negative Spontaneous cell reaction =1 0 0 At equilibrium <1 Negative Positive Non-spontaneous cell reaction; the reaction is spontaneous in the reverse direction The Nernst Equation Standard potentials assume a concentration of 1 M. The Nernst equation allows us to calculate potential when the two cells are not at standard conditions. E = Eo – RT ln(Q) nF R = 8.31 J/(molK) T = Temperature in K n = moles of electrons in balanced redox equation F = Faraday constant = 96,485 coulombs/mol eEo = the voltage at standard conditions Q = the reaction quotient The Nernst Equation Example At 25˚C, the middle term can be simplified to 0.0591 If the reaction Zn(s) + Cu2+(aq) Cu(s) + Zn2+(aq) is carried out with 5.00M Zn2+ and 0.300M Cu2+ at 298K, what is the actual cell voltage? E = 1.10 V – 0.0591 log (5.0 M) = 1.064 V 2 (0.3 M) Calculating an Equilibrium Constant from a Cell Potential Zn + Cu2+ Zn2+ + Cu E0 = + 1.10 V 0 volts = 1.10 – 0.0591 log(K) = (1.10)(2) = log(K) 2 0.0591 1037.2 = K = 1.58 x 1037 At equilibrium, forward and reverse reactions occur at equal rates, therefore: The battery is “dead” ; The cell potential, E, is zero volts Modifying the Nernst Equation (at 25 C): 0 volts = E0 – 0.0591 log(K) n Concentration Cell Both sides have the same components but at different concentrations. Step 1: Determine which side undergoes oxidation, and which side undergoes reduction. We’re going to use Zn instead of Ag Concentration Cell Both sides have the same components but at different concentrations. The 1.0 M Zn2+ must decrease in concentration, and the 0.10 M Zn2+ must increase in concentration Zn2+ (1.0M) + 2e- Zn (reduction) Zn Zn2+ (0.10M) + 2e- (oxidation) Zn2+ (1.0M) Zn2+ (0.10M) Concentration Cell Step 2: Calculate cell potential using the Nernst Equation (assuming 25 C). Zn2+ (1.0M) Zn2+ (0.10M) E = Eo – 0.0591 log(Q) n Eo = 0.0 volts n = 2 Q = (0.10)/(1.0) E = 0.0 – 0.0591 log(0.10) = 0.030 volts 2 1.0 Electrolysis Process in which electrical energy is used to cause non-spontaneous REDOX reactions Opposite of an electrochemical cell Uses and electrolytic cell Electrolytic processes are NOT spontaneous. They have: A negative cell potential, (-E0) A positive free energy change, (+G) Electroplating of Silver Anode reaction: Ag Ag+ + eCathode reaction: Ag+ + e- Ag Electroplating requirements: 1. Solution of the plating metal 3. Cathode with the object to be plated 2. Anode made of the plating metal 4. Source of current Solving an Electroplating Problem How many seconds will it take to plate out 5.0 grams of silver from a solution of AgNO3 using a 20.0 Ampere current? 1 Ampere = 1 coulomb/second Ag+ + e- Ag 5.0g Ag x 1mol Ag x 1mol e- x 96485C x 1s 107.87g 1mol Ag 1mol e- 20.0C = 2.2 x 102s The Nucleus: A Chemist’s View Chapter 18 Nuclear Symbols Mass number, (A) + (p + o n) 235 92 Atomic number, (Z) (number of p+) U Element symbol Radioactivity Spontaneous decay of certain atoms with the evolution of a, b and g particles Comes from the nucleus Radioactivity Alpha (α) Beta (β) Gamma (γ) Nature Helium Nucleus Negative particle with no mass High energy & frequency radiation Charge +2 -1 0 Mass 4 amu 1/1840 amu 0 Movement in To negative electronic field To positive None Penetrating Power Intermediate Strong Weak Balancing Nuclear Equations SAreactants = SAproducts 235 + 1 235 1 = U+ n 92 0 92 + 0 = 142 142 + 91 + 91 3(1) 1 Ba + Kr + 3n 56 56 36 + 36 + SZreactants = SZproducts 0 3(0) Balancing Nuclear Equations #2 1. 226 Ra 88 2. U+n 4 a 2 +? 1 + 2n + ? Alpha Decay Alpha production (a): an alpha particle is a helium nucleus Alpha decay is limited to heavy, radioactive nuclei Limited to VERY large nuclei. Can be shown by α or He symbols as products Beta Decay Beta production (b): A beta particle is an electron ejected from the nucleus Beta emission converts a neutron to a proton Converts a neutron into a proton. Gamma Decay Gamma ray production (g): Gamma rays are high energy photons produced in association with other forms of decay. Gamma rays are mass-less and do not, by themselves, change the nucleus Other Types of Emission Positron Emission Positrons are the anti-particle of the electron Positron emission converts a proton to a neutron Electron Capture Inner-orbital electron is captured by the nucleus Electron capture converts a proton to a neutron Half Life The time taken for half of the atoms of a sample to decay Three methods to determine half life Graphically Use of the expression: ln(N/No) = - kt Use of the expression: fraction of remaining activity = 1/2n Half Life: ln(N/No) = - kt Decay occurs by first order kinetics (the rate of decay is proportional to the number of nuclides present) ln N = -kt t1/2 = ln(2) = 0.693 N0 k k N = number of nuclides remaining at time t N0 = number of nuclides present initially k = rate constant t = elapsed time Half Life: n 1/2 Fraction of the remaining activity = 1/2n n = the number of half lives Examples: After 21 days a radioactive isotope has only one eighth of its original activity. What is the half life of the radioactive isotope? The half life of 14C is 5585 years. A sample of carbon gave a count rate of 2.02. Another sample of carbon gave a count rate of 1.71. Calculate the age of the second sample. Nuclear Stability & Decay Decay will occur in such a way as to return a nucleus to the band (line) of stability. The most stable nuclide is Iron-56 If Z > 83, the nuclide is radioactive Stable (non-radioactive) nuclei tend to have a neutron: proton ratio of 1:1 A radioactive nucleus reaches a stable state by a series of steps Energy and Mass Nuclear changes occur with small but measurable losses of mass. The lost mass is called the mass defect, and is converted to energy according to Einstein’s equation: E = mc2 m = mass defect E = change in energy c = speed of light Because c2 is so large, even small amounts of mass are converted to enormous amount of energy. Nuclear Fission and Fusion Fusion: Combining two light nuclei to form a heavier, more stable nucleus. 3 1 4 0 2 He + 1 H 2 He + 1e Fission: Splitting a heavy nucleus into two nuclei with smaller mass numbers. 1 235 142 91 1 0 n + 92 U 56 Ba + 36 Kr + 3 0 n Fission Processes A self-sustaining fission process is called a chain reaction. Neutrons causing Event fission Result subcritical <1 reaction stops Critical =1 sustained reaction Supercritical >1 violent explosion A Fission Reaction Use of Radiation Medicine Isotopic dating Thickness control in engineering Radioactive source is placed on one side and the detector on the other Leak detection Nuclear fission – power and atomic bomb