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Dr. Streit Week 2 Notes

by: Rachel Ferrell

Dr. Streit Week 2 Notes CHEM 1030 - 003

Marketplace > Auburn University > Chemistry > CHEM 1030 - 003 > Dr Streit Week 2 Notes
Rachel Ferrell
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covers chapter 2 and part of chapter 3
Fundamentals Chemistry I
John D Gorden
Class Notes
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This 4 page Class Notes was uploaded by Rachel Ferrell on Thursday January 28, 2016. The Class Notes belongs to CHEM 1030 - 003 at Auburn University taught by John D Gorden in Fall 2015. Since its upload, it has received 73 views. For similar materials see Fundamentals Chemistry I in Chemistry at Auburn University.


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Date Created: 01/28/16
Rachel  Ferrell   CHEM  1030   1/26/16     Chapter  2:  Atoms  and  the  Periodic  Table     Atoms  First:   • Atom=  the  smallest  quantity  of  matter  that  still  retains  property  of  matter   • Element=  substance  that  cant  be  broken  down  into  2  or  more  simpler  substances  by  any  means     o Ex.  Gold,  oxygen,  helium   • John  Dalton→said  atoms(which  make  up  all  matter)  are  tiny,  indivisible  particles   o He  was  right  in  that  they  are  tiny,  but  atom  can  be  broken  down  into  subatomic  particles   o The  nature  (charge),  number,  and  arrangement  of  subatomic  particles  determine   properties  of  the  atom  and  whatever  thing  the  atoms  make  up  together   Subatomic  Particles  and  Atomic  Structure:   • Discovery  of  the  Electron:   • Radiation=the  emission  and  transmission  of  energy  in  the  form  of  waves   • Cathode  ray  tube=  is  a  glass  tube  with  two  metal  plates  inside  labeled  (+)  and  (-­‐)   o When  exposed  to  voltage,  (-­‐)  plate  emits  a  radiation  called  a  cathode  ray,  which  moves   toward  the  (+)  plate   o Cathode  ray  is  not  visible  to  the  eye,  but  it  emits  a  light  when  it  moves  across  the  tube   • Columb’s  Law=  like  charges  repel  each  other,  and  opposite  charges  attract  each  other   • Thomson   o Used  the  cathode  ray  tube  to  suggest  that  the  cathode  rays  were  actually  (-­‐)  charged   particles  called  electrons  (since  the  rays  moved  towards  the  (+)  plate)   o By  varying  the  electric  field  and  measuring  the  degree  of  deflection  of  the  cathode  rays,  he   ???? determined  the  charge-­‐to-­‐mass  ratio  of  electrons=  1.76  X  ????????  C/g   § C=  Coulomb=  SI  unit  foe  electric  charge   • Millikan   o Determined  the  charge  of  an  electron  by  examining  the  motion  of  tiny  oil  drops     o Charge  of  1  electron=  -­‐1.6022  X  10 -­‐1C   o Knowing  this  charge,  he  could  use  Thomson’s  charge-­‐to-­‐mass  ratio  to  determine  the  mass   of  an  electron   ▯▯▯▯▯▯ ▯▯.▯▯▯▯  ▯  ▯▯▯▯▯ ▯ -­‐28   o Mass  of  an  electron= ▯▯▯▯▯▯/▯▯▯▯  ▯▯.▯▯  ▯  ▯▯ ▯/▯ =  9.10  X  10g   • Radiation   • Rontgen   o Discovered  X-­‐Rays  (type  of  radioactive  emission)   o X-­‐rays  are  not  made  from  charged  particles  because  they  were  not  deflected  from  electric   fields   • Becquerel     o Discovered  radioactivity=spontaneous  emission  of  radiation     o Radioactive  substances  (like  uranium)  can  produce  3  types  of  radiation:   § Alpha(????)  Rays=  (+)  charge,  deflect  away  from  positively  charged  plate   § Beta  (????)  Rays=  (-­‐)  charge,  electrons,  deflect  away  from  negatively  charged  plate   § Gamma  (????)  Rays=  neutral  charge,  like  X-­‐rays  because  no  charge   • Nuclear  Model  of  Atom   • Rutherford   o Before  Rutherford,  Thomson  proposed  plum-­‐pudding  model=  electrons  are  just  embedded   with  the  positive  protons  in  the  nucleus→not  true   o Rutherford  used  ????  particles  to  prove  the  structure  of  the  atom  using  gold  foil   o Majority  of  particles  penetrated  gold  foil  undeflected   o Sometimes,  ????  particles  bounced  back   § This  happened  because  sometimes  the  gold  foil  bounced  off  of  the  nucleus,  which  is   positively  charged  like  the  ????  particles     § Most  of  the  time  though  it  would  hit  the  electron  or  electron  shell  and  would  go   right  through  it   o Proposed  nuclear  model:   § Positive  charge  is  concentrated  in  the  nucleus   § The  rest  is  mostly  empty  space   § Nucleus  is  most  of  the  atoms  mass  and  it  is  an  extremely  dense  core  of  an  atom   § Protons=  (+)   § Electrons=  (-­‐)   § Concluded  that  atom  was  made  of  a  nucleus  that  accounted  for  most  of  its  mass  but   occupied  only  a  tiny  amount  of  its  volume   • Chadwick   o Discovered  neutrons=  electrically  neutral;  mass  slightly  greater  than  protons   Atomic  Number,  Mass  Number,  and  Isotopes:   • All  atoms  can  be  identified  by  number  of  protons,  neutrons     • Atomic  number=  number  of  protons;  defines  an  element   o Since  atoms  are  neutral,  this  is  also  number  of  electrons  (-­‐)   o ▯????????→Z  is  the  atomic  number   • Mass  n▯mber=  total  number  of  protons  and  neutrons   o ▯ ????→  A  is  the  mass  number   o protons  +  neutrons  =  mass  number   • Most  elements  have  more  than  2  isotopes=  atoms  with  the  same  atomic  number  but  different   mass  numbers  (therefore  different  number  of  neutrons)   o Isotopes  of  same  element→usually  have  similar  chemical  properties   o Ex.  Same  types  of  compounds  and  similar  reactivities   Nuclear  Stability:   • Since  the  density  of  an  atom  is  so  high,  some  force  is  needed  to  hold  the  particles  together  so   tightly   • Since  all  protons  in  a  nucleus  are  (+)  charge,  they  repel  each  other  →  repulsion  factor   • However,  the  addition  of  neutrons  can  add  short-­‐range  attraction  between  particles  to   compensate  for  the  repulsion  factor   • Atomic  Stability=  Columb  repulsion  factor  –  short  range  attraction   o Higher  the  density  of  the  atom/element→stronger  repulsion  forces   o Therefore,  atoms  with  a  higher  density  need  more  neutrons  to  stabilize  the  atom   • Neutron-­‐proton  ratio   o Shows  how  stable  an  atom  is   o Elements  with  a  higher  atomic  number  have  a  higher  ratio  (because  of  their  higher  density)   o Patterns  of  Stability   § More  stable  nucleus  with  2,8,20,50,82,126  protons/neutrons   § More  stable  with  even  numbers   § All  atoms  with  atomic  number  >83  are  radioactive   • This  means  the  nucleus  is  unstable  and  will  spontaneously  decay  by  losing   protons  and  neutrons  through  ????  radiation     Average  Atomic  Mass:   • Atomic  mass=  mass  of  an  atom  in  atomic  mass  units  (amu)   o 1  amu  =  1/12  mass  of  Carbon-­‐12  atom   • Average  atomic  mass=  mass  listed  on  the  periodic  table;  represents  average  mass    of  naturally   occurring  mixture  of  the  isotopes     o How  to  calculate  average  atomic  mass  for  multiple  isotopes:   § (???????????????????????????????????? ▯▯▯▯▯▯▯  ▯ )(???????????????? ▯▯▯▯▯▯▯  ▯)  +  (???????????????????????????????????? ▯▯▯▯▯▯▯  ▯ )(???????????????? ▯▯▯▯▯▯▯  ▯ )  =  Av.  Mass   The  Periodic  Table:   • =a  chart  in  which  elements  with  similar  chemical  and  physical  properties  are  grouped  together   • periods=horizontal  rows;  increasing  atomic  number   • metals=  good  conductors  of  heat/electricity(most  known  elements  are  metals)   • nonmetals=poor  conductors  of  heat  and  electricity   • metalloids=intermediate  properties(fewest  number  of  elements  are  metalloids)   • group=  vertical  column  (also  referred  to  as  families)   o  Group  1A→  Alkali  Metals   § Li,  Na,  K,  Rb,  Cs,  Fr   o Group  2A→Alkaline  Earth  Metals   § Be,  Mg,  Ca,  Sr,  Ba,  Ra   o Group  6A→Chalcogens   § O,  S,  Se,  Te,  Po   o Group  7A→Halogens   § F,  Cl,  Br,  I,  At   o Group  8A→Noble  Gases   § He,  Ne,  Ar,  Kr,  Xe,  Rn   o Elements  in  middle→transition  metals   The  Mole  and  the  Molar  Mass:   • Mole=  the  amount  of  substance  that  contains  as  many  atoms  as  there  are  atoms  in  12  grams  of   Carbon-­‐12   • Avagadro’s  Number  (???? )=  6.▯22  X  10   23 o Similar  measurement  as  1  dozen=  12   • Ex.  Problem:   ▯▯ o 30  mol  Ca  X  ▯▯  ▯  ▯▯ ▯▯▯▯▯  ▯▯  =  1.807  X10 atoms  Ca         ▯  ▯▯▯   o 1.00  X  10  X   ▯  ▯▯▯  =  1.66  X  10 mol  Ca   ▯.▯▯▯  ▯  ▯▯▯▯▯▯▯▯▯  ▯▯ • molar  mass=  the  amount  of  mass  in  grams  of  1  mol  of  the  substance   o mass  of  1  mol→same  as  atomic  mass  for  each  element   o usually  expressed  as  g/mol   o Ex.  Problem   § 25  g  C  x    =  2.082  mol  C   ▯▯.▯▯  ▯  ▯ • Interconverting  Mass,  moles,  and  number  of  atoms   o Molar  mass  is  a  conversion  factor  from  mass  to  moles,  and  vice  versa   o Ex.  Problem   ▯  ▯▯▯  ▯ ▯.▯▯▯  ▯  ▯▯▯▯▯▯▯▯▯  ▯ 22   § 0.515  g  C ▯▯.▯▯  ▯  ▯  x   ▯  ▯▯▯  ▯  =  2.58  x  10 atoms  C     Chapter  3:  Quantum  Theory  and  Electronic  Structure  of  Atoms:     Units  of  Energy:   • Joule=  the  SI  unit  for  energy   o the  amount  of  energy  possessed  by  a  2  kg  mass  moving  at  the  speed  of  1  meter/second   § 1  kg  x  m s  =  1  J   o can  also  be  defined  as  the  amount  of  energy  exerted  when  a  force  of  1  Newton(N)    is   applied  over  the  distance  of  1  meter   § 1  Nxm  =  1  J   o also  is  often  is  expressed  in  kJ  (1  kJ  =  1000J)   Nature  of  Light:   • electromagnetic  spectrum=  continuum  of  radiant  energy   o visible  light  is  only  a  small  portion   • Properties  of  Waves:   o All  forms  of  electromagnetic  radiation  travel  in  waves   o Characterized  by:   § Wavelength(????)=  distance  between  identical  points  on  succeeding  waves   § Frequency  (v)  =  number  of  waves  that  pass  through  a  point  in  1  second   § Amplitude=  vertical  distance  from  midline  of  a  wave  to  the  top  peak  of  bottom  of  a   trough     8   8 o Speed  of  Light(c)=  2.99792458  x  10 m/s  OR  3.00  x  10  m/s   o Speed,  wavelength,  and  frequency  are  related  in  the  equation:     § C=  ????  ×  ????      


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