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Dr. Streit Week 6 Notes

by: Rachel Ferrell

Dr. Streit Week 6 Notes CHEM 1030 - 003

Marketplace > Auburn University > Chemistry > CHEM 1030 - 003 > Dr Streit Week 6 Notes
Rachel Ferrell
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Fundamentals Chemistry I
John D Gorden
Class Notes
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This 5 page Class Notes was uploaded by Rachel Ferrell on Friday February 26, 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 52 views. For similar materials see Fundamentals Chemistry I in Chemistry at Auburn University.


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Date Created: 02/26/16
Rachel  Ferrell   CHEM  1030   2/23/16     Chapter  5:  Cont.     Ionic  Compounds  and  Bonding:   • =electrostatic  attraction  that  hold  oppositely  charged  ions  together  in  an  ionic  compound   +   -­‐ o ex.  Na  +  Cl→Na +  Cl →combine  to  form  NaCl  or  sodium  chloride   • chemical  formula=  denotes  the  elements  of  the  ionic  compound  and  the  ratio  that  they  occur  in   • lattice=  a  3D  array  of  cations  and  anions  in  an  ionic  compound   o think  of  like  a  crystal  structure   o Lattice  energy=  amount  of  energy  required  to  convert  a  mole  of  ionic  solid  to  constitute   ions  in  the  gas  phase   § Basically  is  a  measure  of  how  stable  the  ionic  compound  is   o This  formation  of  ionic  bonds  releases  a  large  amount  of  energy   o Lattice  energy  depends  on  the  magnitude  of  charges  and  distance  between  them   o High  lattice  energy→small  radius/distance→most  stable   o High  lattice  energy→higher  charges→most  stable   § If  charges  are  the  same→look  at  the  distance   § Is  distances  are  the  same→look  at  the  charges   § Ex.  Arrange  MgO,  CaO,  and  SrO  by  increasing  lattice  energy   2+ 2+ 2+ • Mg=Mg ,  Ca=Ca ,  Sr=   • Since  same  charge,  so  order  by  atomic  radius   • Answer=  SrO<CaO<MgO   Naming  Ions  and  Ionic  Compounds:   • Monotomic  cation→  add  –ion  to  the  element   • Monoatomic  anion→  add  –ide  to  the  element   • Some  metals  can  form  cations(not  anions)  of  more  than  one  possible  charge   2+ o Ex.  Fe =  Fe(II)   o              Fe =  Fe(III)   Formulas  for  Ionic  Compounds:   • ionic  compounds→always  between  metals  and  nonmetals   • electrically  neutral→sum  of  charges  must  be  zero   o ex.  Al ,  O →Al O 2  3switch  charges  to  be  the  subscript  of  the  opposite  element)   • To  name  ionic  compounds:   o NaBr→  sodium  bromide   o FeCl →iro2(II)  bromide   o CaO→calcium  oxide   o Mg N → 3  m2gnesium  nitride   o Fe S →2 3n(III)  sulfide   o Exceptions   § Transition  metals  can  form  more  than  one  possible  charge;  therefore  the  roman   numerals  must  be  used   § Always  be  sure  to  check  that  the  overall  charge  is  zero   Covalent  Bonding  and  Molecules:   • Elements  with  similar  properties  share  electrons  to  give  each  atom  more  stability  and  a  noble  gas   configuration   • Lewis  Theory  of  Bonding   o Depicts  the  bond  formation  of  H 2  as  2  Hs  sharing  their  electrons   • Covalent  bond=  pair  of  shared  electrons   • Molecule=  combo  of  two  atoms  held  together  by  chemical  bond   o can  be  an  element  or  a  compound   • Law  of  Definite  Proportions=  different  samples  of  the  same  compound  always  have  the  same   mass  ratio  of  elements   • Law  of  Multiple  Proportions=  if  2  elements  can  combine  with  each  other  to  form  2  or  more   different  compounds,  then  the  ratio  of  masses  of  1  element  that  combines  with  a  fixed  mass  of  the   other  element  can  be  expressed  in  small  whole  numbers   o ex.  CO 2  can  also  be  CO→ratio  of  O  to  C  between  these  molecules  is  about  2:1   • Diatomic  Molecules=  contains  2  atoms;  can  be  the  same  element  or  different  elements   o Heteronuclear=  elements  are  different   o Homonuclear=  elements  are  the  same   • Polyatomic  moleules=  more  than  2  atoms     Molecular  Formulas:   • Chemical  formula  denotes  the  composition  of  the  substanc2  (H O)   • Molecular  formula=  shows  the  exact  number  of  atoms  of  each  element  in  a  molecule     o using  the  dot  structure   • some  elements  have  2  or  more  distinct  forms  known  as  allotropes   o ex.  O2  and 3  O   • structural  formula=  shows  not  only  elemental  composition,  but  also  the  general  arrangement  of   atoms     o shows  polarity,  etc.   Empirical  Formula:   • molecular  substances  can  be  represented  using  empirical  formulas=  whole  number  ratio  of   elements  in  a  compound   o the  simplest  chemical  formula   o ex.  Molecular  formula=  N 2 4   o              Empirical  formula=  NH   2 • empirical  and  molecular  formulas  can  often  be  the 2same  (H O)   • ex.  Problems:   o C H 6 →12 O6   2 o C H 5 →5HN5   o N O→2 O   2 Naming  Molecular  Compounds:   • Binary  molecular  compound→just  2  elements  (two  nonmetals)   • How  to  name   o 1)  Name  the  first  element  in  formula   o 2)Name  the  2  element,  change  ending  to  –ide   • ex.  HCl→hydrogen  chloride   • Greek  prefixes  used   o Mono=1   o Di=2   o Tri=3   o Etc.   • Rule=  don’t  use  mono-­‐  for  the  first  element,  only  for  the  second   o Ex.  Carbon  monoxide  not  monocarbon  monoxide   • Ex.  Problems:   o 1)NF →nit3ogen  trifluoride   o 2)  N O →  dinitrogen  tetroxide   2 4 o 3)  Sulfur  tetrafluoride→  SF   4 o 4)  Tetraphosphorus  decasulfide→  P S 4 10 Compounds  Containing  Hydrogen:   • usually  don’t  conform  to  normal  naming  system  used  for  molecular  compounds   • many  compounds  with  Hydrogen  have  common  names:   o B H =2ib6rane   o SiH =  Silane   6 o NH = 4Ammonia   o PH = 3 Phosphine   o H O= 2  Water   o H S= 2  Hydrogen  sulfide   • Acid=  produces  H  ions  when  dissolved  in  water   o ex.  HCl  is  an  acid  when  dissolved  in  water→   o Naming  Acids:   § 1)  Remove  –gen  from  hydrogen   § 2)change  –ide  ending  on  2  element  to  –ic   o ex.  Hydrochloric  acid   o a  compound  must  contain  at  least  one  ionizable  hydrogen  atom  to  be  an  acid  when   dissolved   § ionizable  hydrogen  is  one  that  separates  from  the  molecule  when  dissolved  in  water   + and  becomes  an  H  ion   Organic  Compounds:   • also  do  not  follow  the  normal  naming  system  for  molecular  compounds   • organic  molecule=  has  carbon  and  hydrogen,  sometimes  other  atoms   • alkanes=simplest  hydrocarbons   o 1  carbon=  methane   o 2  carbon=  ethane   o 3  carbon=  propane   o 4  carbon=  butane   o 5  carbon=  pentane   • many  organic  compounds  have  functional  groups   o often  determine  a  molecules  reactivity   o Alcohol=  -­‐OH   o Aldehyde=  -­‐CHO   o Carboxyl(acid)=  -­‐COOH   o Amine(base)=  -­‐NH  2 Covalent  Bonding  in  Ionic  Species:   • Polyatomic  ions=  combination  of  2  or  more  atoms,  held  together  by  covalent  bond   • Compounds  with  polyatomic  ions→follows  same  rules  as  naming  ionic  compounds     o Molecule  must  be  neutral  overall   o Ex.  Ca  and  P4 →Ca (PO3) 4 2 • Common  polyatomic  ions  below:  MEMORIZE   •   • Ex.  Problems:   o 1)  Fe (SO ) →  Fe ,  SO →  iron(III)  sulfate   2 4 3 4 o 2)  Al  (OH) → 3  Al ,  OH →  aluminum  hydroxide   o 3)  Hg O→ 2  Hg ,  O →  mercury  (I)  oxide   • Oxoanions   o =polyatomic  ions  that  contain  1  or  more  oxygen  atoms  and  one  atom  (central  atom)  of   another  element   o Naming  Oxoanions:   § 1)  Start  with  oxoanions  that  end  in  –ate  and  as  reference  point   § 2)  Ion  with  one  more  O=  per…ate  ion   • ClO =3hlorate,  ClO =  4erchlorate  ion   § 3)  Ion  with  one  less  O=  -­‐ite  ion   • ClO =2  chlorite  ion   § 4)  Ion  with  2  less  O=  hypo…ite  ion   • ClO=  hypochlorite  ion   o Must  memorize  common  –ate  oxoanions   § Chlorate=  ClO 3 -­‐ -­‐ § Nitrate=  NO 3 § Phosphate=  PO   4 3-­‐ § Sulfate=  SO  -­‐ 4 § Bromate=  BrO   3-­‐ • Oxoacids   + o Polyatomic  ions  that  have  1  or  more  oxygen  atom  plus  central  atom,  AND  can  release  H ion   when  dissolved  in  water   o Naming  oxoacids:   § 1)  An  acid  with  -­‐ate→change  to  –ic  acid   • HClO =ch3oric  acid     § 2)  An  acid  with  -­‐ite→change  to  –ous  acid   • HClO = 2 chlorus  acid   § 3)Prefixes  in  oxoanion  names  stay  the  same  for  corresponding  acids   • HClO = 4 perchloric  acid   • HClO=  hypochloric  acid   o Oxoacids  can  be  monoprotic  or  polyprotic   § Monoprotic=  only  has  1  ionizable  H  atom   § Polyprotic=  has  more  than  1  ionizable  H  atom   § Names  of  anions  indicate  the  number  of  remaining  hydrogen   • H 3O = 4 phosphoric  acid   • H 2O = 4 dihydrogen  phosphate  ion   2-­‐ • HPO = 4 hydrogen  phosphate  ion   • PO =4  phosphate  ion        


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