Chap11-Part2

# Chap11-Part2 - Combined Gas Law: ...

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Unformatted text preview: Combined Gas Law:  ‐ It's the Gas law that combines all the other gas laws: Boyles law   Charles Law and Gay‐Lussac's law  PV= nRT  PV/T = nR  PV/T = k (for a ﬁxed amount of gas)             P1V1/T1 = P2V2/T2  Review Session Problem  Solid potassium chlorate (KClO3) decomposes to produce solid  potassium chloride (KCl) and oxygen gas (O2) according to the balanced  chemical equaPon: 2KClO3 (s)  2KCl (s) + 3O2 (g)  What volume of gas, measure at 23.4oC and 75.4 kPa, will be produced  when 32.5 g of potassium chlorate is decomposed?  Given,  P = 75.4kPa   (75.4kPa) (1 atm)/ (101.325kPa) = 0.745 atm  T = 23.4oC = 273.1+23.4 = 296.5 K  32.5g of KClO3 decomposed  ? g of oxygen produced  2 moles of KClO3 produce 3 moles of Oxygen  2(122.55g) of KClO3 produce 3(32g) of oxygen  32.5g of KClO3 produce (3X32) (32.5)/ (2X 122.55) = 12.73g of Oxygen  n = 12.73/32 = 0.398 mol  V= (0.398) (0.08206) (296.5) / 0.745 = 13 L  1)  Ion‐dipole  2)  dipole – dipole  3)  H‐bonding  4)  London dispersion interacPon  The strengths of intermolecular interacPons are almost always less than the   strengths of covalent bonds. For example, less than 16 kJ/mol is needed to   overcome the intermolecular acracPons between HCl molecules, but 431 kJ/mol is   needed to break the covalent bond.  When a substance undergoes   a change from solid to liquid to gas, the covalent bonds usually remain intact.    Boiling and melPng points are largely determined by the strength of   intermolecular interacPons.  Above, the covalent bond is represented as a  solid line whereas the intermolecular interacPon is represented as a dashed line.  Three types of non‐covalent intermolecular interacPons are known to exist   between neutral molecules:  1) dipole‐dipole interacPons  2) London dispersion interacPons  3) hydrogen‐bonding  The three types of intermolecular acracPons are called as van der  Waals interacPons.  There is a fourth known interacPon important in soluPons, called the  4) ion‐dipole interacPon  Other non‐covalent interacPons have also been idenPﬁed and characterized,  ofen in the context of biochemistry, but these interacPons will not be   discussed here.  1) Dipole‐Dipole Interac>ons  Even though a molecule may not have a net charge, the molecule   may have a “permanent” dipole (one end posiPve and the other   negaPve).   In this case, there exists the possibility of molecular   acracPon if the two molecules arrange themselves as follows:  +  ‐  ‐  +  Molecules do not need to be charged to acract each other.  They  only have to be “polarized”.  1) Dipole‐Dipole Interac>ons  There are both acracPve and  repulsive interacPons.  Two molecules acracPng each  other spend more Pme together  than two molecules repelling  each other. The overall eﬀect   is a net acracPon.    For molecules of approximately  equal mass and size, the strengths  of intermolecular interacPons  increase with increasing polarity.  Even though the molecular weight and size are about equal,  the boiling point increases with increasingdipole moment.  This result implies that the intermolecular interacPons are  stronger for molecules with higher dipole moments.   The more polar the molecule, the higher is its boiling point.  2) London Dispersion Interac>ons  Even though some molecules are nonpolar on the average, they can sPll  have temporary asymmetries in their electron distribuPons and therefore  “instantaneous” dipole moments.  Two such atoms can then experience   acracPve interacPons as for molecules with permanent dipoles.   The temporary dipole on one atom can induce a temporary dipole on   the adjacent atom, causing the forces of acracPon. This interacPon called  as ‘London dispersion force’  2) London Dispersion Interac>ons  4 He    2   While the electrons in the 1s orbital of helium would repel  each other (and, therefore, tend to stay far away from  each other), it does happen that they occasionally wind up  on the same side of the atom.  2) London Dispersion Interac>ons   At that instant, then, the helium atom is polar, with an  excess of electrons on the lef side and a shortage on the  right side.  2) London Dispersion Interac>ons   Another helium nearby, then, would have a dipole  induced in it, as the electrons on the lef side of helium  atom 2 repel the electrons in the cloud on helium atom 1.  2) London Dispersion Interac>ons   London dispersion forces, or dispersion forces, are  acracPons between an instantaneous dipole and an  induced dipole.  2) London Dispersion Interac>ons  •  These forces are present in all molecules, whether they  are polar or nonpolar.  •  The tendency of an electron cloud to distort in this way is  called polarizability.  2) London Dispersion Interac>ons  Dispersion forces depend on the molecular polarizability, which describes  how easily the electron cloud can be rearranged.    More polarizable molecules have stronger dispersion forces.    Polarizabilites tend to increase with molecular size because there are more   electrons and they are farther from the nuclei.  2) London Dispersion Interac>ons  Molecular shapes also inﬂuence the strengths of dispersion forces.  C5H12  CH3 CH3 H3C CH3 H3C CH3 N‐pentane has stronger dispersion forces and a higher boiling point  because its more linear shape allows for greater intermolecular interacPons.  Which Have a Greater Eﬀect?  Dipole‐Dipole InteracPons or Dispersion Forces  •  All molecules experience dispersion forces, whether they are  polar or not.  For polar molecules, the strength of the  dispersion interacPons is usually much greater than the  strength of the dipole‐dipole interacPons.  •  If two molecules are of comparable size and shape, dipole‐ dipole interacPons will likely the dominaPng force.  •  If one molecule is much larger than another, dispersion forces  will likely determine its physical properPes.  Eﬀect of London Dispersion Forces & Polarizability  on Physical Proper>es  Fritz London (1900‐1954)  ‐ TheorePcal Physicist  ‐  London Dispersion forces  ‐  ElectromagnePc properPes of Superconductors  ‐ His studies supported Superﬂuidity concept    (images from Google)  3) Hydrogen Bonding  Boiling points generally increase   with molecular size because of   the higher London dispersion   forces.  Water deviates considerably from   this trend.  The deviaPon is explained by the   high propensity of water to   hydrogen bond with itself.  Hydrogen bonds are stronger  than dipole‐dipole or dispersion  forces but weaker than covalent  bonds.   •  The nonpolar series (SnH4 to CH4) follow the expected trend.  •  The polar series follows the trend from H2Te through H2S, but water is quite an  anomaly.  3) Hydrogen Bonding  Hydrogen bonding is a special type of dipole‐dipole   interacPon.  These bonds occur between the hydrogen atom in  a   polar bond (H‐F, H‐O, H‐N, etc.) and an unshared   electron pair on a nearby electronegaPve atom   (F, O, N, etc.).  The posiPve side of the bond dipole consists of the   nearly bare proton of the hydrogen nucleus. The   negaPve side of the bond dipole consists of the   electronegaPve atom in the other molecule.  Because hydrogen is small, it can come close to its   molecular partner and increase the interacPon  strength.  Water & Ammonia:  Solid lines show  covalent bonds and dashed lines  show hydrogen bonds.  3) Hydrogen Bonding  Unlike for most substances, solid water is less dense   than liquid water.  This phenomenon can be explained   by hydrogen bonding.  In ice, the molecules maximize their hydrogen bonding,  leading to an open, ordered arrangement. In water,  the arrangement collapses, although hydrogen bonding  is sPll very prominent.   Eﬀect of Hydrogen Bonding on Physical Proper>es  Importance of H‐Bonding:  ‐ Stabilizing the structure of Proteins  ‐ Holds the double stranded DNA   ‐ Carrying GenePc InformaPon with DNA  ‐ Drug delivery systems  Books on Hydrogen Bonding  4) Ion‐Dipole Forces  These forces exist between a charged ion and the parPal charge on the end  of a polar molecule.  They are parPcularly important for ions dissolved in  polar liquids; e.g., salt water.  4) Ion‐Dipole Forces  Iden>fying Intermolecular Forces  London dispersion forces are present in all cases.  The strengths of the  other interacPons generally increase from lef to right.  What kind of intermolecular acracPve force is illustrated in the following examples?  What kind of intermolecular acracPve force is illustrated in the following examples?  hydrogen bond  What kind of intermolecular acracPve force is illustrated in the following examples?  hydrogen bond  London dispersion  What kind of intermolecular acracPve force is illustrated in the following examples?  hydrogen bond  London dispersion  ion‐dipole  What kind of intermolecular acracPve force is illustrated in the following examples?  hydrogen bond  London dispersion  ion‐dipole  dipole‐dipole  Which member of the following pairs has stronger dispersion forces?  (a) CH3CH2SH or CH3CH2CH2SH  (b) CH3CH2CH2Cl or (CH3)2CHCl?  Which member of the following pairs has stronger dispersion forces?  (a) CH3CH2SH or CH3CH2CH2SH  (b) CH3CH2CH2Cl or (CH3)2CHCl?  (larger)  Which member of the following pairs has stronger dispersion forces?  (a) CH3CH2SH or CH3CH2CH2SH  (larger)  (b) CH3CH2CH2Cl or (CH3)2CHCl?  (less spherical)  substance                             structure      boiling point  ethylene glycol                     HOCH2CH2OH         198oC  ethylene glycol dimethyl ether        CH3OCH2CH2OCH3    83oC  ethyl alcohol    ethyl methyl ether                         CH3CH2OH                     CH3CH2OCH3            78oC           11oC  Explain why replacing the hydrogen atoms on the oxygen atoms with CH3  results in a lower boiling point.  substance                             structure      boiling point  ethylene glycol                     HOCH2CH2OH         198oC  ethylene glycol dimethyl ether        CH3OCH2CH2OCH3    83oC  ethyl alcohol    ethyl methyl ether                         CH3CH2OH                     CH3CH2OCH3            78oC           11oC  Explain why replacing the hydrogen atoms on the oxygen atoms with CH3  results in a lower boiling point.  The eﬀect cannot be one of molecular size leading to higher dispersion forces,  which would have the opposite eﬀect.    The eﬀect is probably due to decreased hydrogen bonding.  ...
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