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CHMA11 ~ Textbook Notes.docx

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Chemistry
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CHMA11H3
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CHMA11 ~ TEXTBOOK NOTES 10.2 VESPR Theory: The Five Basic Shapes (400-403)  valence shell electron pair repulsion (VESPR) theory: allows prediction of shapes of molecules based on idea that electrons- either as lone pairs or as bonding pairs- repel one another  based on idea that electron groups repel one another through coulombic forces  electron groups: lone pairs, single bonds, multiple bonds, or lone electrons in a molecule  repulsions between electron groups on interior atoms of a molecule determine geometry of molecule  preferred geometry of a molecule is one where electron groups have max separation (min energy) possible  for molecules with one interior (central) atom, molecular geometry depends on  number of electron groups around central atom  how many of those electron groups are bonding and how many are lone pairs Two Electron Groups: Linear Geometry  linear geometry: molecular geometry of 3 atoms with 180° bond angle due to repulsion of 2 electron groups  molecules that form only 2 single bonds, with no lone pairs, rare because they don’t follow octet rule  same geometry observed in all molecules that have 2 electron groups and no lone pairs  Three Electron Groups: Trigonal Planar Geometry  trigonal planar geometry: molecular geometry of four atoms with 120° bond angles in a plane  double bond contains more electron density than single bond, exerts slightly greater repulsion on single bonds  Four Electron Groups: Tetrahedral Geometry  tetrahedral geometry: molecular geometry of 5 atoms with 109.5° bond angles  Five Electron Groups: Trigonal Bipyramidal Geometry  trigonal bipyramidal geometry: molecular geometry of 6 atoms with 120° bond angles between 3 equatorial electron groups and 90° bond angles between the two axial electron groups and the trigonal plane  Six Electron Groups: Octahedral Geometry  octahedral geometry: molecular geometry of 7 atoms with 90° bond angles  10.8 Molecular Orbital Theory: Electron Delocalization (432-)  11.2, 11.3, 11.4, 11.5, 11.6, 11.7, 11.8, 11.11, 11.12, 12.2, 12.3, 12.4, 12.5, 12.6, 12.7 CHAPTER 13 13.2 The Rate of a Chemical Reaction (564-569)  if chemical reaction has slow rate, only relatively small fraction of molecules react to form products in given period of time; fast rate means large fraction of molecules react  rate of chemical reaction measured as change in amounts of reactants or products (usually in concentration units) divided by change in time  reaction rate: negative of change in concentration of a reactant divided by change in time  negative sign usually part of definition when reaction rate is defined in terms of a reactant because reactant concentrations decrease a reaction proceeds; change in concentration of a reactant is negative  negative sign makes overall rate positive  change in concentration of a product positive; no negative sign when rate defined in respect to product  definition of rate with respect to each reactant and product must reflect stoichiometric coefficients of reaction The Average Rate of the Reaction  calculate average rate of reaction using: –Δ *A+/Δt (M/s)  rate is average rate within given time interval Instantaneous Rate of the Reaction  rate at any one point in time; represented by instantaneous slope of curve at that point  can be determined from slope of tangent to curve at point of interest  rate is same whether we use one of reactants or product for calculation  generic reaction: aA + bB → cC + dD rate of reaction:  example on page 567 Measuring Reaction Rates  polarimetry: measuring degree of lighting passing thorough a reacting solution  spectroscopy: most common; light of specific wavelength passed through sample, intensity of transmitted light (depends on how much light absorbed by sample) is measured and recorded  reactions in which number of moles of gaseous reactants and products changes as reaction proceeds can be readily monitored by measuring changes in pressure 13.3 The Rate Law: The Effect of Concentration on Reaction Rate (569-572) n  rate law: relationship between rate of reaction and concentration of reactants; rate = k [A]  k= rate constant: constant of proportionality  n= reaction order: determines how rate depends on concentration of reactants  if n=0, reaction is zero order, rate independent of concentration of A  if n=1, reaction is first order, rate directly proportional to concentration of A  if n=2, reaction is second order, rate proportional to square of concentration of A  other orders possible, including non-integers, but these most common Zero-Order Reaction  rate = k[A] = k  rate constant because reaction doesn’t slow down as concentration of A decreases  rate same at any concentration of A  occur under conditions where amount of reactant available for reaction unaffected by changes in overall quantity of reactant First-Order Reaction 1  rate = k[A]  rate slows down as reaction proceeds because concentration of reactant decreases  rate directly proportional to concentrations Second-Order Reaction 2  rate = k[A]  rate more sensitive to reactant concentration Determining the Order of Reaction  order of reaction can be determined only by experiment  method of initial rates: initial rate measured by running reaction several times with different initial reactant concentrations to determine effect of concentration on rate  can determine value of k by solving rate law for k and substituting concentration and initial rate from any one of the measurements st  for 0 order reaction concentration doubles as initial rate stays constant, for 1 order reaction, concentration doubles as initial rate doubles, for 2 order concentration doubles as initial rate quadruples  -1 st -1 nd -1 -1  rate constant for 0 order reaction is Ms , 1 order reaction is s , 2 order reaction is M s  example on page 570 Reaction Order for Multiple Reactants  consider reaction aA + bB → cC + dD m n  as long as reverse reaction negligibly slow, rate = k [A] [B]  overall order: sum of orders of all reactants in a chemical reaction  rate law for any reaction must be determined experimentally  example of rate law and rate order on page 572 13.4 The Integrated Rate Law: The Dependence of Concentration on Time (573-580)  integrated rate law: relationship between concentrations of the reactants in a chemical reaction and time First-Order Integrated Rate Law  in A → products, rate directly proportional to concentration  rate = k [A]  ; aka differential rate law  first order integrated rate law: ln[A] t -kt + ln[A] 0r  [A] = concentration of A at any time, k= rate constant, [A] = initial concentration of A st t 0  for 1 order reaction, plot of natural log of reactant concentration as function of time yields straight like with slope of –k and y-intercept of ln[A] 0  example on page 575-576 Second-Order Integrated Rate Law  in A → produ2ts, rate proportional to square of concentration of A  rate = k [A]   second order integrated rate law:  for straight line, plot inverse of concentration of reactant as function of time; slope of k, intercept of 1/[A]0  example on page 577 Zero Order Integrated Rate Law  rate proportional to constant  rate = k [A] = k   zero order integrated rate law: [A] = -kt + [A] t 0  for straight line, plot concentration of reactant as function of time; slope of k, intercept of [A] 0 The Half-Life of a Reaction  half-life (1/2: time required for concentration of reactant or amount of radioactive isotope to fall to one-half of its initial value  half-life expression: defines dependence of half-life on rate constant and initial concentration; different for different reactions First-Order Reaction Half-Life  integrated rate law:  half-life of first-order reaction:  t1/2independent of initial reaction  constant half-life unique to first order reactions  example on page 579 Second-Order Reaction Half-Life  integrated rate law:  half-life of second-order reaction:  half-life depends on initial concentration  half-life continues to get longer as concentration decreases Zero-Order Reaction Half-Life  integrated rate law: [A] t -kt + [A] 0  half-life of first order reaction:  half-life depends on initial concentration 13.5 The Effect of Temperature on Reaction Rate (581-588)  temperature dependence of reaction rate contained in k  k only constant when temperature constant  increase in temperature generally results in increase in k, resulting in faster rates  Arrhenius equation: relates rate constant of reaction to temperature, activation energy, and frequency factor  ; k= rate constant, T=temperature (K)  R= gas constant= 8.34 J/molK  frequency factor (A): aka pre-exponential factor; number of times that reactants approach activation energy per unit time  activation energy (E a: energy barrier in chemical reaction that must be overcome for reactants to be converted into products The Activation Energy  activated complex (transition state): high-energy intermediate state between reactant and product  activation energy is energy required to reach activated complex  higher Ea, slower reaction rate (at given temperature) The Frequency Factor  approaching activation barrier not equivalent to surmounting it  most approaches don’t have enough total energy to make it over activation barrier The Exponential Factor  exponential factor: number from 0-1 that represents fraction of molecules that have enough energy to pass activation barrier on given approach  fraction of approaches that are successful and result in product   as temperature increases, number of molecules having enough thermal energy to pass activation barrier increases Arrhenius Plots: Experimental Measurements of the Frequency Factor and the Activation Energy  ( )  Arrhenius plot: plot of natural log of rate constant (ln k) versus inverse of temperature in kelvins (1/T) that yields a straight line with slope oa –E /R and y-intercept of ln A  example on page 584-585  when either data are limited or plotting capabilities absent, can calculate activation energy if we know rate constant at 2 different temperatures  2-point form of Arrhenius equation: ( )  example on page 585-586 The Collision Model: A Closer Look at the Frequency Factor  consider 2 gas phase reactants: A(g)B (g)roducts  collision model: model of chemical reactions in which reaction occurs after a sufficiently energetic collision between two reactant molecules  each approach to activation barrier is a collision between the reactant molecules  frequency factors of most gas-phase chemical reactions tend to be smaller than number of collisions that occur per second   orientation factor (p): fraction of sufficiently energetic collisions  collision frequency (z): number of collisions that occur per unit time; can be calculated for a gas-phase reaction from pressure of gases and temperature of reaction mixture  under typical conditions, single molecule undergoes on the order of 10 collisions every second  small orientation factor indicates that orientational requirements for this reaction very stringent  reactions between individual atoms usually have orientation factor of ~1 because atoms spherically symmetric  if orientation factor greater than 1, collisions aren’t needed for reaction  harpoon mechanism: positive charge on potassium and negative charge on bromine cause 2 species to attract each other and form a bond without a collision 13.6 Reaction Mechanisms (588-591)  overall equation doesn’t show intermediate steps  reaction mechanism: series of individual chemical steps by which an overall chemical reaction occurs  elementary step: an individual step in a reaction mechanism; can’t be broken down into simpler steps  individual steps in mechanism add to overall reaction  reaction intermediates: species that are formed in step of a reaction mechanism and consumed in another  can piece together a reaction mechanism by measuring kinetics of overall reaction and working backward to write a mechanism consistent with measured kinetics Rate Laws for Elementary Steps  molecularity: number of reactant particles involved in elementary step; unimolecular and bimolecular most common  unimolecular: describes a reaction that involves only one particle that goes on to form products  bimolecular: an elementary step in a reaction that involves two particles, either the same species or different, that collide and go on to from products  termolecular: an elementary step of a reaction in which three particles collide and go on to form products; rare because probability of three particles colliding simultaneously is small  rate law for overall chemical reaction can’t be deduced from balanced chemical equation  rate law proportional to product of concentration of particles in elementary step Rate-Determining Steps and Overall Reaction Rate Laws  rate-determining step: step in a reaction mechanism that occurs much slower than any of the other steps’ limits overa
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