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Chemisty 1210, Week 2 Notes

by: Grace Campbell

Chemisty 1210, Week 2 Notes CHEM 1210

Marketplace > Ohio State University > CHEM 1210 > Chemisty 1210 Week 2 Notes
Grace Campbell
GPA 3.687

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These notes cover everything that was gone over in lecture and in the book from 8/29/16 to 9/2/16
General Chemistry I
Class Notes
General Chemistry, Chemistry
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This 12 page Class Notes was uploaded by Grace Campbell on Sunday September 4, 2016. The Class Notes belongs to CHEM 1210 at Ohio State University taught by in Fall 2016. Since its upload, it has received 3 views.


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Date Created: 09/04/16
WEEK  2  NOTES  8/29/16  –  9/2/16   Chapter  2:  Atoms,  Molecules,  and  Ions     A.   The  Atomic  Theory  of  Matter   a.   Dalton’s  Atomic  Theory   i.   Each  element  is  composed  of  extremely  small  particles  called  atoms   ii.   All  atoms  of  a  given  element  are  identical,  but  the  atoms  of  one  element   are  different  from  the  atoms  of  all  other  elements   iii.   Atoms  of  one  element  cannot  be  changed  into  atoms  of  a  different   element  by  chemical  reactions   1.   Atoms  are  neither  created  nor  destroyed  in  chemical  reactions   2.   Law  of  Conservation  of  Mass   iv.   Compounds  are  formed  when  atoms  of  more  than  one  element  combine   1.   A  given  compound  always  has  the  same  relative  number  and   kinds  of  atoms   a.   Law  of  Constant  Composition   2.   Atoms  combine  in  whole  numbers   a.   Law  of  Multiple  Proportions   b.   Dalton’s  Atomic  Theory  explains  the  law  of  constant  composition  and  the  law  of   conservation  of  mass,  and  deduced  the  law  of  multiple  proportions   B.   The  Discovery  of  Atomic  Structure   a.   Atoms   i.   Atomic  mass  unit  (amu)  =  1.6603  x  10^-­‐24  g     ii.   Mass  of  atoms:  1  –  260  amu   iii.   Radii   1.   Use  nm   iv.   1  Angstrom  =  10^-­‐10  m  =  10^-­‐8  cm   b.   Subatomic  particles   i.   Make  up  atoms     1   ii.   Electron     1.   Charge  =  -­‐1.6022  x  10^-­‐19  C;  relative  =  -­‐1   2.   m  =  9.1094  x  10^-­‐28  g  =  5.486  x  10^-­‐4  amu   3.   atoms  are  neutral  until  electrons  are  gained  or  lost   4.   number  of  electrons  in  an  atom  =  number  of  protons  in  the  atom   5.   electrons  are  in  constant  motion   6.   Can  be  traded  back  and  forth  between  atoms   iii.   Proton   1.   Charge  =  1.6022  x  10^-­‐19  C,  relative  =  +1   2.   m  =  1.6726  x  10^-­‐24  g  =  1.0073  amu   3.   make  up  about  ½  of  mass  of  an  atom   4.   number  of  protons  in  an  atom  of  an  element  is  that  element’s   atomic  number   5.   The  number  of  protons  do  not  change  in  a  normal  chemical   reaction   iv.   Neutron   1.   charge  =  0   2.   m  =  1.6749  x  10^-­‐24  g  =  1.0088  amu   3.   make  up  about  ½  of  mass  of  an  atom   v.   Particles  of  the  same  charge  repel  each  other;  particles  of  different   charges  attract  each  other   c.   Cathode  rays  and  Electrons   i.   The  radiation  produced  when  a  high  voltage  is  applied  to  electrodes  in  a   tube   ii.   Cause  certain  materials  to  give  off  light  (fluoresce)   iii.   J.J.  Thompson   1.   Discovered  the  electron  by  observing  cathode  rays  as  streams  of   negatively  charged  particles   2.   Measured  charge  to  mass  ratio  (coulombs)     2   iv.   Robert  Millikan-­‐  Oil  Drop  Experiment   1.   Determined  the  charge  of  an  electron:  1.6022  x  10^-­‐19   d.   Radioactivity-­‐  the  spontaneous  emission  of  radiation   i.   Ernest  Rutherford  found  3  types  of  radiation  with  his  gold  foil  experiment   ii.   Alpha   1.   Bent  by  an  electric  field   2.   Fast  moving  particles   3.   Have  positive  charge  of  +2;  attracted  to  a  negative  charge   iii.   Beta   1.   Bent  by  an  electric  field   2.   High  speed  electrons   3.   Radioactive  equivalent  of  cathode  rays   4.   Have  a  charge  of  -­‐1;  attracted  to  a  positive  charge   iv.   Gamma   1.   Unaffected  by  an  electric  field   e.   The  Nuclear  Model  of  the  Atom   i.   Atoms  contain  a  dense  nucleus   1.   Nucleus  composed  of  protons  and  neutrons   ii.   Nucleus  surrounded  by  electrons  in  mostly  empty  space   1.   Electrons  are  in  a  constant  motion   iii.   Atoms’  diameter  =  10^-­‐8  cm   C.   The  Modern  View  of  Atomic  Structure   a.   Atomic  Number  (Z)   i.   Number  of  protons  in  an  atom   ii.   Periodic  table  is  arranged  by  increasing  atomic  number   b.   Mass  Number  (A)   i.   A  =  number  of  protons  +  number  of  neutrons   ii.   Is  an  integer   c.   Elemental/Atomic  Symbol     3   i.   Represents  an  element   ii.   EX:  Zn  =  zinc   iii.   Number  in  top  left  corner  is  the  element’s  mass  number   iv.   Number  in  the  bottom  left  corner  is  the  element’s  atomic  number   d.   Isotopes   i.   Atoms  of  the  same  element  that  have  the  same  number  of  protons  but  a   different  number  of  neutrons   ii.   Have  the  same  atomic  number,  but  different  mass  number   iii.   Some  mass  gets  lost  in  isotopic  masses  because  it  gets  converted  to   energy  when  the  atoms  come  together   e.   Atomic  Weight   i.   Weighted  average  of  the  isotopes  of  an  atom   ii.   Given  in  amu   iii.   The  average  atomic  mass   1.   Doesn’t  give  you  the  actual  weight  of  the  atom,  but  gives  you  an   average   D.   Periodic  Table   a.   The  arrangement  of  elements  in  order  of  increasing  atomic  number   i.   Modern  periodic  law   b.   Places  elements  with  similar  properties  in  vertical  columns   c.   Groups  or  families   i.   Vertical  columns   ii.   Elements  within  a  group  share  similar  properties   iii.   Columns  labeled  at  the  top  by  either  numbers,  letters,  or  roman   numerals   iv.   Representative  Elements  (Main  group  elements)   1.   The  first  2  columns  and  the  last  6  columns   2.   1A-­‐  8A   3.   1A  =  alkali  metals  (except  for  hydrogen)     4   4.   2A  =  alkaline  earth  metals   5.   7A  =  halogens   6.   8A  =  noble  gases   v.   Transition  Metals   1.   The  middle  10  columns   2.   1B-­‐  8B   d.   Periods   i.   Horizontal  rows   ii.   The  2  long  rows  below  the  main  body  of  the  periodic  table  are  inner   transition  elements   1.   Lanthanides  –  part  of  row  6   2.   Actinides      -­‐  part  of  row  7   e.   The  stair  step  semimetals  are  transitional  metals   E.   Metals  vs  Nonmetals   a.   Metals   i.   Solids   1.   Except  Hg   ii.   Metallic  luster   iii.   Malleable  and  ductile   1.   Can  be  pounded  into  shapes  and  drawn  out  into  a  wire   iv.   Conduct  heat  and  electricity  well   v.   Oxides  (compounds):   1.   Nonvolatile  (ionic  compounds)   2.   High  melting   3.   MgO,  Na2O   b.   Nonmetals   i.   Gases  or  solids   1.   Except  Br   ii.   Variety  of  color  and  appearance     5   iii.   Solids  are  brittle   iv.   Not  good  conductors   1.   Insulators   v.   Oxides:   1.   Volatile   2.   Low  melting   3.   CO,  CO2,  SO2   F.   Molecules  and  Molecular  Compounds   a.   Molecular  formula   i.   Chemical  formula  that  indicates  the  actual  numbers  of  atoms  in  a   molecule   ii.   Subscripts  are  always  some  integer  multiple  of  the  subscripts  in  the   empirical  formula   b.   Empirical  formula   i.   Chemical  formula  that  gives  only  the  relative  number  of  atoms  of  each   type  in  a  molecule   ii.   The  smallest  whole  number  ratio  of  atoms   c.   Molecular  elements   i.   Diatomic  molecule-­‐  a  molecule  made  up  of  2  atoms   1.   EX:  H2…  has  2  H  atoms  bonded  together   2.   All  of  the  halogens  are  diatomic   ii.   Polyatomics   1.   A  molecule  made  up  of  more  than  2  atoms   2.   EX:  P4,  S8,  O3   d.   Molecular  Compounds   i.   Compounds  composed  of  molecules  containing  more  than  one  type  of  an   atom  (contain  atoms  of  different  elements)   ii.   EX:  CO2,  H2O   e.   Structural  formula     6   i.   Shows  which  atoms  are  attached  to  which   ii.   Demonstrates  structure   G.   Ions  and  Ionic  Compounds   a.   Ion-­‐  a  charged  particle  that  is  formed  when  electrons  are  added  or  removed   from  an  atom   i.   Cation-­‐  positively  charged  ion   1.   Metals  tend  to  lose  electrons  and  become  cations   ii.   Anion-­‐  negatively  charged  ion   1.   Nonmetals  tend  to  gain  electrons  and  become  anions   b.   Polyatomic  ion-­‐  atoms  joined  as  a  molecule  but  have  a  net  charge   i.   NH^4+,  SO4^2-­‐   c.   Predicting  ionic  charges   i.   Many  atoms  gain/lose  electrons  to  end  up  with  the  same  number  of   electrons  as  the  noble  gas  closest  to  them  in  the  periodic  table   1.   Representative  elements  (1A-­‐  8A)   ii.   Elements  in  the  same  isoelectronic  series  have  the  same  electron   configuration   iii.   Periodic  Table   1.   Cation  groups   a.   Charge  =  group  number   i.   Except  for  H   b.   Special  cations:   i.   Al  =  +3   ii.   Zn  =  +2   iii.   Ag  =  +1   2.   Anion  groups   a.   Charge  =  group  number  –  8   3.   5A  =  -­‐3,  6A  =  -­‐2,  7A  =  -­‐1   iv.   Ionic  Compounds     7   1.   Compound  made  of  cations  and  anions   a.   Oppositely  charged  ions  are  held  together  by  electrostatic   attractions   2.   Compounds  of  metals  and  nonmetals   a.   Electrically  neurtral   3.   Crystalline  solids  (salts)   4.   EX:  Ca^2+  and  CO3^2-­‐  form  CaCO3  (calcium  carbonate)   5.   Compounds  are  NOT  molecules   a.   They  are  formula  units   6.   Empirical  formula-­‐  chemical  formula  that  shows  the  simplest  ratio   of  ions  in  a  compound   H.   Naming  Inorganic  Compounds   a.   Organic  compounds  contain  carbon  and  hydrogen,  all  other  compounds  are   considered  inorganic   b.   Chemical  nomenclature-­‐  system  of  naming  substances   c.   Monatomic  ions   i.   Cations   1.   Formed  from  metal  atoms:  name  of  the  elements  followed  by   “ion”   2.   EX:  K^+  is  potassium  ion   ii.   Anions   1.   Add  “-­‐ide  ion”  to  the  end  of  the  element’s  name   2.   EX:  Br^-­‐  is  bromide  ion   d.   Stock  System   i.   Metals  that  can  have  more  than  1  possible  charge   1.   Transition  and  representative  metals   ii.   Use  roman  numerals  in  parenthesis  after  the  element’s  name  to   represent  the  ion’s  charge   iii.   EX:  Fe  ^3+  ….  Iron  (III);  Sn  ^4+  ….  Tin  (IV)     8   e.   Polyatomic  Ions   i.   Group  of  chemically  bonded  atoms  with  an  overall  charge   ii.   Cations   1.   EX:  NH4  ^+      ammonium  ion   2.   H3O^+            hydronium  ion   3.   Hg2^2+        mercury  (I)  ion   iii.   Anions   1.   Ending  in  “-­‐ide”   a.   EX:  OH^-­‐          hydroxide  ion   2.   Others   a.   EX:  C2H3O2^-­‐          acetate  ion   iv.   Oxyanions-­‐  polyatomic  anions  that  contain  oxygen   1.   Suffixes   a.    “-­‐ate”  (most  common)     b.   “-­‐ite”  (same  charge  but  1  fewer  O  atom)   2.   Prefixes   a.   “per-­‐“   i.   1  more  O  atom  than  “-­‐ate”   b.   “hypo-­‐“   i.   1  less  O  atom  than  “-­‐ite”   3.   EX:   a.   ClO4^-­‐                perchlorate   b.   ClO3^-­‐                              chlorate   c.   ClO2^-­‐                                chlorite   d.   ClO^-­‐                      hypochlorite   e.   Cl^-­‐                                                chloride     f.   ***Overall  charge  remains  the  same   4.   Addition  of  H^+  to  a  -­‐2  or  -­‐3  oxyanion   a.   Results  still  charged  (anions)     9   b.   EX:   i.   HCO3^-­‐                bicarbonate/  hydrogen  carbonate   ii.   HPO4^2-­‐            monohydrogen  phosphate     1.   Mono-­‐  is  usually  dropped   iii.   H2PO4^-­‐              dihydrogen  phosphate   5.   Acids   a.   Acid-­‐  H^+  produces  a  neutral  compound  when  combined   with  an  anion   b.   Not  ionic   i.   But  ionize  in  H2O  to  produce  H^+  (H3O^+)   c.   EX:  HCl  (g)        H2O                H^+  (aq)    +    Cl^-­‐  (aq)   d.   Binary  (2  element)  acids                   i.   Hydrogen  +  nonmetal   1.   –ide      à  hydro-­‐        -­‐ic  acid   a.   EX:  HF  (aq)            hydrofluoric  acid   2.   Per-­‐      -­‐ate    à    per-­‐        -­‐ic  acid   3.   –ate    à        -­‐ic  acid   4.   –ite      à        -­‐ous  acid   5.   hypo-­‐      -­‐ite    à      hypo-­‐      -­‐ous  acid   I.   Binary  Molecular  Compounds   a.   Compounds  contain  2  different  elements   i.   Nonmetals  or  nonmetals  and  semimetals   b.   Element  further  left  in  the  periodic  table  (closest  to  the  metals)  is  usually  listed   first   i.   Exception  when  the  compound  contains  oxygen  and  any  halogen  (except   for  fluorine),  then  oxygen  is  always  written  last   c.   If  both  elements  are  in  the  same  group,  the  lower  one  is  listed  first  (less   electronegative)   d.   The  name  of  the  second  element  is  given  an  –ide  ending       10   e.   Greek  prefixes  are  used  to  indicate  the  number  of  atoms  of  each  element   i.   Mono-­‐  is  never  used  with  the  first  element   1.   When  the  prefix  ends  with  an  a  or  an  o  and  the  second  element   begins  with  a  vowel,  the  prefix  is  dropped   ii.   Prefixes:   1.   Mono-­‐   2.   Di-­‐   3.   Tri-­‐   4.   Tetra-­‐   5.   Penta-­‐   6.   Hexa-­‐   7.   Hepta-­‐   8.   Octa-­‐   9.   Nona-­‐   10.  Deca-­‐   iii.   EX:  N2O4              Dinitrogen  tetroxide   iv.                SO3                    Sulfur  trioxide   J.   Simple  Organic  Compounds   a.   Hydrocarbons-­‐  compounds  that  contain  only  carbon  and  hydrogen   b.   Alkanes-­‐  simplest  class  of  hydrocarbons   i.   Each  carbon  is  bonded  to  4  other  atoms   ii.   Methane      CH4   iii.   Ethane              C2H6   iv.   Propane        C3H8   c.   Functional  groups   i.   Hydrogens  are  replaced  with  functional  groups-­‐  specific  groups  of  atoms   ii.   Alkenes   1.   Ethene        CH2    (double  bond)  CH2   iii.   Alkynes     11   1.   Ethyne          CH  (triple  bond)  CH   d.   Alcohols   i.   Replace  hydrogen  atom  with  an  –OH  group   ii.   Methanol,  ethanol,  propanol   1.   Are  structural  isomers  of  each  other  (have  same  chemical  formula   but  different  arrangements  of  atoms)   e.   Ethers   i.   Dimethyl  ether,  diethyl  ether   ii.   Functional  group  isomers  with  alcohols   f.   Aldehydes   i.   C  (double  bond)  O  group  …  carbonyl   ii.   EX:  HCHO          formaldehyde   g.   Ketones   i.   EX:  CH3COCH3            acetone   h.   Carboxylic  acids   i.   Carboxyl  group   ii.   EX:  H-­‐CO2H        formic  acid   iii.                CH3CO2H          acetic  acid   i.   Amines   i.   organic  analogues  of  NH3  (ammonia)   ii.   EX:  CH3NH2        methyl  amine                 12  


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