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Chem 101 Exam #1 Study Guide

by: Josh Trac

Chem 101 Exam #1 Study Guide Chem 1010

Marketplace > Clemson University > Chemistry > Chem 1010 > Chem 101 Exam 1 Study Guide
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Notes for Exam #1 in Chemistry 101.
Chemistry 101
William Wallace
Study Guide
Chemistry, exam, midterm, Study Guide, study, guide, 101, notes, final, test, definitions, Chem
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This 10 page Study Guide was uploaded by Josh Trac on Tuesday January 5, 2016. The Study Guide belongs to Chem 1010 at Clemson University taught by William Wallace in Winter 2016. Since its upload, it has received 80 views. For similar materials see Chemistry 101 in Chemistry at Clemson University.


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Date Created: 01/05/16
Fall  2015_CH1010_Dr.  Kreider-­‐Mueller   CH1010  Exam  #1  Study  Guide   For  reference  see  “Chemistry:  An  Atoms -­‐focused  Approach”  by  Gilbert,  Kirss,  and  Foster     Chapter  1:  Matter  and  Energy,  An  Atomic  Perspective   • Know  the  following  SI  base  units  of  measure  and  their  abbreviations  (Table  1.2  in  text):   *Mass  (kg)       *Temperature  (K)     *Time  (s)     *Length  (m)       *Amount  of  substance  (mol)   • You  need  to  know  (memorize)  the  following  prefixes  for  multiples  of  SI  Units     o Mega,  kilo,  deci,  centi,  milli,  micro,  and  nano  (Table  1.1  in  your  textbook).     o Be  able  to  do  the  following:   § Know  the  abbreviations  for  each  of  the  prefixes   3 § Know  the  conversion  factors  (ex.  1  g  =  10  mg)   § Be  able  to  perform  these  types  of  conversions  in  word  problems   • Know  the  equations  for  converting  between  Kelvin  &  Celsius  (we  will  not  provide  you   with  these  equations  on  the  exam!)   K  =  ˚C  +  273.15˚       ˚C  =  K  –  273.15˚   • Know  what  the  difference  is  between  SI  base  units  and  derived  units  and  how  to   perform  conversions  with  derived  units  (ex:  convert  3  m  to  3  cm )   3 o Know  the  following  conversions  for  volume:   1  dm  =  1  L         1  cm  =  1  mL   o Know  the  equation  for  density  (g/mL)  (we  will  not  provide  you  with  this   equation!)  and  how  to  use  it  in  a  word  problem.  Practice  solving  for  density  given   a  mass  and  volume.  Practice  solving  for  mass  given  a  density  and  volume.  Practice   solving  for  volume  given  a  density  and  mass.  Recognize  how  changes  in  mass   and/or  volume  affect  density.   o Unit  for  energy  is  Joules  (J).  1  J  =  (kg•m )/s   • Know  how  to  work  with  scientific  notation.  Given  a  number,  be  able  to  represent  it  in   scientific  notation.  Given  a  number  in  scientific  notation,  be  able  to  represent  the  number   in  its  standard  form.  (See  Appendix  1  in  your  textbook)   • KNOW  THE  RULES  OF  SIGNIFICANT  FIGURES:   1) All  nonzero  digits  are  significant   2)  Zeros  in  the  middle  of  a  #  are  significant   § Ex:  2.9905  has  5  significant  figures                3)  Zeros  at  the  beginning  of  a  #  are  not  significant   § Ex:  0.025  has  2  significant  figures                4)  Zeros  at  the  end  of  a  #  and  after  a  decimal  point  are  significant   § Ex:  245.0  has  4  significant  figures   5)  Zeros  at  the  end  of  a  #  and  before  a  decimal  point  may  or  may  not  be  significant   § Ex:  34,200  may  have  3,  4,  or  5  significant  figures   § In  this  case  it  is  best  to  use  scientific  notation  to  indicate  the  proper  number             of  significant  figures     1  of  10   Fall  2015_CH1010_Dr.  Kreider-­‐Mueller   • Know  the  rules  for  significant  figures  when  performing  addition  &  subtraction,  as  well   as  multiplication  &  division.     • What  is  the  difference  between  accuracy  and  precision  in  measurements?   • Know  how  to  round  numbers:     o If  the  1st  digit  you  remove  is  less  than  5  round  down  by  dropping  it  and  all   following  digits.     o It  the  1st  digit  you  remove  is  5  or  greater,  round  up  by  adding  1  to  digit  on  the  left.     • PRACTICE  CONVERTING  BETWEEN  ONE  UNIT  AND  ANOTHER.  We  have  done   countless  examples  in  class,  on  quizzes,  in  the  lecture  slides,  on  HMWK  assignments  etc.     You  need  to  be  comfortable  with  conversion  problems.  You  may  need  to  perform  several   conversions  in  any  one  problem.     • Be  able  to  distinguish  between  solids,  liquids,  and  gases  at  all  3  levels  of  representation:   symbolic  (H O),2  particulate  (one  molecule  of  H O)2  and  macroscopic  (1  g  of  2 O)   o Be  able  to  recognize  graphical  representations  of  molecules   o The  properties  of  matter  are  related  to  its  molecular  level  structure   • Understand  that  chemistry  is  an  experimental  science  and  depends  on  accurate,  precise,   and  reproducible  measurements     o Scientific  Method:  an  approach  to  acquiring  knowledge  based  on  observation  of   phenomena,  development  of  a  testable  hypothesis,  and  additional  experiments   that  test  the  validity  of  the  hypothesis.   o Hypothesis:  a  tentative  and  testable  explanation  for  an  observation  or  a  series  of   observations   o Scientific  Theory:  A  general  explanation  of  widely  observed  phenomena  that  has   been  extensively  tested.   o Scientific  Law:  a  concise  and  generally  applicable  statement  of  a  fundamental   scientific  principle   • Matter  can  be  divided  into  two  principle  classes  (Figure  1.5  in  textbook):   o Pure  substances:  Matter  that  cannot  be  separated  into  simpler  matter  by  a   physical  process   § Element:  a  pure  substance  that  cannot  be  separated  into  simpler   substances  by  an  chemical  process   • Atom:  the  smallest  particle  of  an  element  that  retains  the  chemical   characteristics  of  the  element   § Compound:  a  pure  substance  that  is  composed  of  2  or  more  elements   linked  together  in  fixed  proportions  and  that  can  be  broken  down  into   those  elements  by  some  chemical  process.  (A  sample  of  a  given  compound   may  contain  1  molecule  or  several  molecules  of  that  compound.)   • Molecule:  a  collection  of  atoms  chemically  bonded  together   o Mixtures:  a  combination  of  pure  substances  in  variable  proportions  in  which  the   individual  substances  retain  their  chemical  identities  and  can  be  separated  from   one  another  by  a  physical  process     2  of  10   Fall  2015_CH1010_Dr.  Kreider-­‐Mueller   § Homogeneous  Mixture:  a  mixture  in  which  the  components  are   distributed  uniformly  throughout  and  have  no  visible  boundaries  or  regions   § Heterogeneous  Mixture:  a  mixture  in  which  the  components  are  not   distributed  uniformly,  so  that  the  mixture  contains  distinct  regions  of   different  compositions     Introduction  to  Energy   • Matter  exists  in  3  phases  or  physical  states:  solid,  liquid  gas   • Matter  can  be  transformed  from  one  physical  state  to  another  as  its  temperature  is  raised   or  lowered:   o solid  è  liquid  (melting;  endothermic)   o liquid  è  gas  (vaporization;  endothermic)   o solid  è  gas  (sublimation;  endothermic)   o gas  è  liquid  (condensation;  exothermic)   o liquid  è  solid  (freezing;  exothermic)   o gas  è  solid  (deposition;  exothermic)   • Understand  the  difference  between  an  exothermic  and  endothermic  process   o Exothermic  process:  energy  flows  from  a  system  into  its  surroundings   o Endothermic  process:  energy  flows  from  the  surroundings  into  the  system   • Be  able  to  identify  the  system  and  the  surroundings  for  phase  changes.  Where  does  the   energy  (heat)  flow?  From  the  system  to  the  surroundings?  Or  from  the  surroundings  to   the  system?   • Understand  the  difference  between  heat  and  temperature:   o Heat:  a  flow  of  energy  from  one  object  or  place  to  another  due  to  differences  in  the   temperatures  of  the  objects  or  places   o Thermal  energy:  the  portion  of  the  total  internal  energy  of  a  system  that  is   proportional  to  its  absolute  temperature   o Temperature:  a  measure  of  thermal  energy   • What  is  kinetic  energy?       KE  =  ½mv                  (where  m  =  mass,  v  =  velocity)   • What  is  potential  energy?   • Why  is  thermal  energy  classified  as  kinetic  energy  rather  than  potential  energy?     Chapter  2:  Atom,  Ions,  and  Molecules   • Know  the  following  Laws:   o Law  of  Mass  Conservation:  Mass  is  neither  created  nor  destroyed  in  chemical   reactions   o Law  of  Definite  Proportions:  Different  samples  of  a  pure  chemical  compound   always  contain  the  same  proportion  of  elements  by  mass   o Law  of  Multiple  Proportions:  Elements  can  combine  in  different  ways  to  form   different  chemical  compounds,  whose  mass  ratios  are  simple  whole-­‐number   multiples  of  each  other.     3  of  10   Fall  2015_CH1010_Dr.  Kreider-­‐Mueller   • Know  the  various  points  in  Dalton’s  Atomic  Theory:   1) Matter  is  composed  of  small  particles  called  atoms.   2) Atoms  of  the  same  element  are  identical  in  shape  and  mass,  but  differ  from  the   atoms  of  other  elements.   3) Atoms  of  an  element  cannot  be  changed  into  atoms  of  a  different  element  by   chemical  reactions.  Atoms  cannot  be  created  or  destroyed  in  chemical   reactions.   4) Atoms  of  different  elements  may  combine  with  other  atoms  in  fixed,  simple,   whole  number  ratios  to  form  compounds.  A  given  compound  always  has  the   same  relative  number  and  kind  of  atoms.   ∗ Note:  we  now  know  that  atoms  can  be  further  subdivided  into  protons,   neutrons,  and  electrons  (but  not  by  chemical  processes).    We  also  know  that   atoms  of  some  elements  vary  in  their  masses  and  densities  (isotopes).  Dalton’s   2  postulate  stated  above  is  not  correct!   • What  were  the  main  conclusions  from  these  experiments  (What  did  we  learn  from   them?  Section  2.1  in  textbook)   o Thomson’s  cathode  ray  experiment     o Millikan’s  Oil  drop  experiment     o Rutherford’s  Gold  Foil  experiment   • Understand  our  current  model  of  the  atom:   o Nucleus:  the  positively  charged  center  of  an  atom  that  contains  nearly  all  the   atom’s  mass   o Proton:  a  subatomic  particle,  present  in  the  nucleus  of  an  atom,  that  has  a   relative  charge  of  1+  and  a  mass  number  of  1   o Neutron:  An  electrically  neutral  (uncharged)  subatomic  particle  with  a  mass   number  of  1   o Electron:  a  subatomic  particle  that  has  a  relative  charge  of  1−  and  essentially   zero  mass   • Be  able  to  relate  the  symbols  and  names  for  the  elements  in  the  first  four  rows  of  the   periodic  table  (elements  hydrogen  through  krypton)   • Dimitri  Mendeleev  published  a  table  that  is  considered  the  forerunner  of  the  modern   periodic  table  of  elements.     • The  modern  periodic  table  arranges  the  elements  in  order  of  their  atomic  numbers   • What  is  a  period  of  elements?  All  of  the  elements  in  a  row  of  the  periodic  table.   • What  is  a  group  or  family  of  elements?  All  of  the  elements  in  a  column  of  the  periodic   table.  Elements  in  a  group  have  similar  properties!   • Know  how  to  use  the  periodic  table  to  look  up  the  atomic  number  (Z)  and  Mass   Number  (A)  of  different  elements.     • Know  how  to  read/write  a  chemical  symbol  for  a  given  element.  Where  to  we  put  the   mass  number?  Where  do  we  put  the  atomic  number?   o Atomic  number  (Z):  the  number  of  protons  in  the  nucleus  of  an  atom     4  of  10   Fall  2015_CH1010_Dr.  Kreider-­‐Mueller   o Mass  Number  (A):  the  number  of  nucleons  in  an  atom   § Nucleon:  a  proton  or  neutron  in  a  nucleus   • Know  how  to  calculate  the  mass  number  (A)   A  =  Number  of  protons  (Z)  +  number  of  neutrons   • Understand  the  difference  between  Mass  Number  and  Average  Atomic  Mass.   o Average  Atomic  Mass:  the  weighted  average  of  masses  of  all  isotopes  of  an   element,  calculated  by  multiplying  the  natural  abundance  of  each  isotope  by  its   mass  in  atomic  mass  units  and  then  summing  the  products.  The  average  atomic   mass  is  not  a  whole  number!   o Isotope:  atoms  of  an  element  containing  the  same  number  of  protons  but   different  numbers  of  neutrons.   o Be  able  to  identify  the  information  given  by  the  symbol  for  the  isotope  of  an   element   o Be  able  to  calculate  atomic  masses  from  relative  abundances  of  isotopes   and  vice  versa.     • Know  where  the  following  groups  are  located  in  the  periodic  table,  and  be  familiar   with  a  few  of  their  properties  (recognize  that  there  are  multiple  ways  to  label  groups   in  the  periodic  table):   o Transition  Metals     o Inner  Transition  Metals   § Lanthanides   § Actinides   o Main  Group  Elements   § Main  Group  nonmetals   • Halogens  (Group  7A  or  Group  17)   • Noble  gases  (Group  8A  or  Group  18)   § Main  Group  Metals   • Alkali  Metals  (Group  1A  or  Group  1)   • Alkaline-­‐Earth  Metals  (Group  2A  or  Group  2)   • How  do  metals,  semimetals,  and  nonmetals  differ?  Where  are  they  located  in  the   periodic  table?  What  are  the  basic  properties  of  each?  Which  ones  are  good   conductors  of  electricity?   • Be  able  to  recognize  the  difference  between  ionic  solids  and  discrete  molecules   o Ion:  An  atom  or  molecule  that  has  a  positive  or  negative  charge   § Cation:  A  positively  charged  ion   § Anion:  A  negatively  charge  ion   o Ionic  solid:  a  solid  consisting  of  monatomic  or  polyatomic  ions  held  together   by  ionic  bonds   o Be  able  to  assign  charges  to  the  ion  derived  from  MAIN  GROUP  elements   (Figure  2.10  in  textbook)     5  of  10   Fall  2015_CH1010_Dr.  Kreider-­‐Mueller     • Interpret  chemical  formulas  &  calculate  the  formula  mass  of  a  compound.   • Understand  that  in  our  macroscopic  world,  quantities  of  substances  contain  extremely   large  numbers  of  atoms,  ions,  or  molecules!  We  use  the  mole  (the  chemist’s  dozen)  to   express  quantities  of  substances.  The  mole  relates  macroscopic  quantities  of   substances,  such  as  their  masses  expressed  in  grams,  to  the  #  of  particles  they  contain.   o Mole:  An  amount  of  a  substance  that  contains  Avogadro’s  number  of  particles   (atoms,  ions,  molecules,  or  formula  units)   o Avogadro’s  number  (N ):  6.022A  ×  10 .  The  number  of  carbon  atoms  in   exactly  12  grams  of  the  carbon-­‐12  isotope   o Molar  Mass:  the  mass  of  1  mole  of  a  substance.   o Be  able  to  convert  between  the  mass  of  a  substance  and  the  amount  (#  of   moles)  and  vice  versa.     Chapter  3:  Atomic  Structure,  Explaining  the  Properties  of  Elements   • What  is  radiant  energy?  What  is  the  electromagnetic  spectrum?   o Electromagnetic  Radiation:  any  form  of  radiant  energy  in  the  electromagnetic   spectrum   • Know  the  order  of  the  different  regions  in  the  electromagnetic  spectrum  (Figure   3.1  in  textbook).  In  order  of  increasing  wavelength:   Gamma  rays  <  X-­‐rays  <  UV  <  Visible  <  IR  <  Microwaves  <  Radio  waves     6  of  10   Fall  2015_CH1010_Dr.  Kreider-­‐Mueller     • Know  the  different  parts  of  a  wave:  amplitude,  peak,  trough,  wavelength,  frequency   • How  are  frequency  and  wavelength  related?   c  =  νλ where  c  =  speed  of  light  (2.998  ×  10  m/s)   ν =  frequency  (Hz  or  s )   λ  =  wavelength  (m) • Understand  the  development  of  the  current  model  of  the  electronic  structure  of   atoms   o Understand  what  the  Balmer-­‐Rydberg  equation  tells  us  (YOU  DO  NOT  NEED   TO  USE/MEMORIZE  THIS  EQUATION  FOR  THE  EXAM!)   o What  is  a  continuous  spectrum?  What  does  it  look  like?   o What  is  an  emission  (line)  spectrum?  What  does  it  look  like?   § Atomic  emission  spectra:  characteristic  patterns  of  bright  lines   produced  when  atoms  are  vaporized  in  high-­‐temperature  flames  or   electrical  discharges.   o What  is  an  absorption  spectrum?  What  does  it  look  like?   § Atomic  Absorption  spectra:  characteristic  patterns  of  dark  lines   produced  when  an  external  source  of  radiation  passes  through  free,   gaseous  atoms.   o Use  the  concepts  of  energy  levels  and  orbitals  to  explain  the  occurrence  of   emission  and  absorption  spectra   o Understand  the  development  of  the  dual  wave/particle  nature  of  the   electrons.   § The  photoelectric  effect:  the  release  of  electrons  from  a  material  as  a   result  of  electromagnetic  radiation  striking  it.   • Understand  the  experimental  setup   • Be  able  to  solve  word  problems  involving  the  photoelectric  effect     7  of  10   Fall  2015_CH1010_Dr.  Kreider-­‐Mueller   o Work  function  (Φ):  the  amount  of  energy  needed  to   dislodge  an  electron  from  the  surface  of  a  material   o Threshold  Frequency  (ν ):  the  mi0imum  frequency  of   light  required  to  produce  the  photoelectric  effect   • Does  the  intensity  of  a  beam  of  light  affect  the  number  of   electrons  ejected  from  a  metal?  If  so,  how?   § What  does  it  mean  when  we  say  energy  is  quantized?   • Quantized:  having  values  restricted  to  whole-­‐number  multiples   of  a  specific  base  value   • Quantum:  the  smallest  discrete  quantity  of  a  particular  form  of   energy   • Photon:  a  quantum  of  electromagnetic  radiation   § Know  De  Broglie’s  Hypothesis  and  how  to  use  it  in  a  word  problem   λ  =  h/(mv)       −34 where  h  =  Planck  constant  (6.626  ×  10  J•s)   m  =  mass       v  =  velocity   § Know  what  Heisenberg’s  Principle  tells  us  (YOU  DO  NOT  NEED  TO   USE/MEMORIZE  THIS  EQUATION  FOR  THE  EXAM!)   • The  principle  that  one  cannot  simultaneously  know  the  exact   position  and  the  exact  momentum  of  an  electron   • Know  how  to  use  Quantum  Numbers:   o Quantum  Number:  one  of  four  related  numbers  that  specify  the  energy,  shape,   and  orientation  of  orbitals  in  an  atom  and  the  spin  orientation  of  electrons  in   the  orbitals.   o What  is  an  orbital?   § Regions  in  an  atom  where  the  probability  of  finding  an  electron  is  high   2 § Defined  by  the  square  of  the  wave  function  (ψ )   o How  do  you  determine  possible  values  of  n,  l,  m, l m ?s  (section  3.9  of  textbook)   o What  do  each  of  the  different  quantum  numbers  tell  us  about  an  orbital?   o What  letters  correspond  to  each  value  of  l?   o Know  what  a  shell  is,  and  which  quantum  number  it  refers  to   § Shell:  orbitals  with  the  same  value  of  n  are  in  the  same  shell   o Know  what  a  subshell  is,  and  which  quantum  number  it  refers  to   § Subshell:  orbitals  with  the  same  values  of  n  and  l  are  in  the  same   subshell   • Be  able  to  identify  an  orbital  (s,  p,  d)  based  on  it’s  shape  and  orientation  on  a  3D  set  of   axes   • What  is  a  node?  How  many  nodes  are  in  a  p  orbital?  How  many  are  in  a  d  orbital?   o Node:  a  location  in  a  standing  wave  that  experiences  no  displacement         8  of  10   Fall  2015_CH1010_Dr.  Kreider-­‐Mueller   • Electron  configurations:   o Know  how  to  write  and  identify  electron  configurations  using  the  standard   (expanded)  form  and  shorthand  versions.     o Be  able  to  use/generate  an  orbital-­‐filling  diagram   o Know  the  following  Principles  &  Rules:   § Pauli  Exclusion  Principle:  No  2  electrons  in  an  atom  can  have  the  same   values  of  their  4  quantum  numbers   § Hund’s  Rule:  If  2  or  more  orbitals  with  the  same  energy  are  available,   one  electron  goes  in  each  until  all  are  half  full.  The  electrons  in  the  half-­‐ filled  orbitals  all  have  the  same  value  of  their  spin  quantum  number.   § Aufbau  Principle   •  Lower  energy  orbitals  fill  before  higher  energy  orbitals.     • An  orbital  can  hold  only  2  electrons,  which  must  have  opposite   spins  (Pauli  Exclusion  principle).     • If  2  or  more  degenerate  orbitals  are  available,  1  electron  goes   into  each  until  all  are  half-­‐full  (Hund’s  Rule)   o  Know  that  with  increasing  energy,  the  gap  between  one  shell  and  the   next  shell  decreases.  (Ex.  there  is  a  large  energy  gap  between  1s  and  2s   orbitals,  but  a  smaller  energy  gap  between  2s  and  3s  orbitals)       o Be  able  to  explain  anomalous  electron  configurations  for  Cr  and  Cu   § Cr:  [Ar]3d 4s   § Cu:  [Ar]3d 4s   1 o Recognize  the  relationship  between  an  element’s  electron  configuration,  the   number  of  valence  electrons,  and  its  Group  Number.  Elemental  families  have   similar  electron  configurations   § Core  Electrons:  Electrons  in  the  filled,  inner  shells  in  an  atom  or  ion   that  are  not  involved  in  chemical  reactions   § Valence  Electrons:  electrons  in  the  outermost  occupied  shell  of  an   atom  having  the  most  influence  on  the  atom’s  chemical  behavior   § Group 1A atom: [Noble Gas] ns 1 2 § Group 2A atom: [Noble Gas] ns 2 4 § Group 6A atom: [Noble Gas] ns np § Group 7A atom: [Noble Gas] ns np 2 5 § How  do  electron  configurations  help  us  explain  why  metals  tend  to  form   cations  and  nonmetals  tend  to  form  anions?   o Be  able  to  write/identify  electron  configurations  for  ions   § s  block  elements:  form  monatomic  cations  by  losing  all  the  outer-­‐shell   electrons  leaving  their  ions  with  the  electron  configuration  of  the  noble   gas  immediately  preceding  them  in  the  periodic  table.     • Ex:  Na  =  [Ne]3s   Na  =  [Ne]   § p  block  elements:  form  monatomic  anions  by  gaining  enough  electrons   to  completely  fill  its  valence-­‐shell  p  orbitals.  The  ion  formed  has  the     9  of  10   Fall  2015_CH1010_Dr.  Kreider-­‐Mueller   electron  configuration  of  the  noble  gas  at  the  end  of  its  row  in  the   periodic  table   2 5 − 2 6 • Ex:  Cl  =  [Ne]3s 3p   Cl  =  [Ne]3s 3p  or  [Ar]   § d  block  elements:  the  electrons  in  orbitals  with  the  highest  n  value   ionize  first  .  When  forming  cations,  transition  metals  lose  outer-­‐shell  s   electrons  first,  followed  by  d  electrons   • Ex:  Ni  =  [Ar]3d 4s   Ni  =  [Ar]3d   • Ex: Fe: [Ar]3d 4s 6 2 Fe : [Ar]3d 6 Fe : [Ar]3d 5 • Trends  to  know  (and  be  able  to  explain  the  trend…think  about  Z ):   eff o Effective  Nuclear  Charge  (Z ):  the effattraction  toward  the  nucleus  experienced   by  an  electron  in  an  atom;  the  positive  charge  on  the  nucleus  reduced  by  the   extent  to  which  other  electron  in  the  atom  shield  the  electron  from  the  nucleus   o Trend  for  atomic  radius  of  neutral  atoms  (Figure  3.34  in  textbook)   o Ionic  radius  (Figure  3.35  in  textbook)   § When  you  form  a  cation,  the  radius  shrinks  (ex:  the  ionic  radius  of  Na  is   smaller  than  the  atomic  radius  of  Na).  Why?   − § Whey  you  form  an  anion,  the  radius  expands  (ex:  the  ionic  radius  of  Cl   is  larger  than  the  atomic  radius  of  Cl).  Why?   o Ionization  Energy   + − § M  +  Energy  →  M  +  e              M  +  Energy  →  M  +  e              M  +  Energy  →  M  +  e              …and  so  forth   § Look  at  Figure  Table  3.2  in  your  book  (Think  about  when/why  there  are   large  jumps  between  successive  ionization  energies)   § Be  able  to  make  predictions  about  ionization  energies  for  various   elements   o Electron  Affinity   § Which  group  in  the  periodic  table  has  the  highest  (most  negative)   electron  affinity?  Which  group(s)  in  the  periodic  table  have  the  lowest   (values  close  to  zero)  electron  affinity   § Look  at  Figure  3.37  in  your  book   § Be  able  to  make  predictions  about  electron  affinities  for  various   elements   § Group  1A  elements  have  low  electron  affinities  so  they  tend  to  lose  their   ns  electron   § Group  2A  elements  have  low  electron  affinities  so  they  tend  to  lose  their   2 ns  electrons   § Group  7A  elements  have  large  electron  affinities  so  they  tend  to  gain  one   electron,  adopting  the  electron  configuration  of  the  neighboring  noble   gas   § Group  8A  elements  are  inert,  and  don’t  generally  gain  or  lose  electrons       10  of  10  


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