Admin stuff

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Admin stuff

• OWLs remaining OWLs due Mon.
• Reading Current Ch. 5.1-5.7
• Next test Nov. 20 (Mon, 6-730 PM)
• Exam will be in Bartlett 65
• Will not be multiple choice like sample exam

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Solution chemistry
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Demo

• Strong electrolytes
• Weak electrolytes
• Non-electrolytes
• If we mix equal amounts of acid and base and test
  the solution conductivity with a light bulb, its
  intensity will be
• (A) Bright (B) Dim (C) Zero

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Acid-base reactions

• strong acid strong base --gt react to produce
  water
• strong acid weak baseweak acid strong base
  acids and bases always react
• weak acid weak base
• Acids and bases always react with each other
• When they react, one or more protons is
  transferred in the reaction
• When they react, a salt is formed
• HCl(aq) NaOH(aq) -gt H2O(aq) NaCl(aq)
• HCl(aq) NH3(aq) -gt NH4Cl(aq)
• H2CO3(aq) 2 KOH(aq) - gt K2CO3(aq) H2O(aq)
• NH3(aq) CH3COOH(aq) -gt NH4CH3COO(aq)

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Oxoacid nomenclature

• -ate anions form -ic acids
• sulfuric acid (H2SO4)
• nitric acid (HNO3)
• -ite anions form -ous acids
• sulfurous acid (H2SO3)
• nitrous acid (HNO2)

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Acid nomenclature

• Can be named two ways
• Molecular hydrogen anion name hydrogen
  chloride (HCl) hydrogen sulfide (H2S)
• Acid hydro ic acid hydrochloric
  acid (HCl) hydrosulfuric acid (H2S)

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Gas-forming reactions

• Practical examples include
• Baking soda (NaHCO3 sodium bicarbonate)
• when mixed with acid, CO2 gas forms
• Limestone (CaCO3) also gives off CO2 in the
  presence of acid

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Gas-forming reactions

• (1) Gas forms as a result of decomposition of a
  product
• H2CO3 --gt H2O CO2
• H2SO3 --gt H2O SO2

• Examples
• K2CO3(aq) 2 HI(aq)
• -gt 2 KI(aq) H2O(aq) CO2(g) (not H2CO3)
• Li2SO3(aq) 2 HBr(aq) -gt (not H2SO3)
• 2 LiBr(aq) H2O(aq) SO2(g)

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Mentos?

• http//eepybird.com/exp214.html

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Gas-forming reactions

• A gaseous molecule is a reaction product
• S2- 2H --gt H2S CO2
• NH4 OH --gt NH3(g) H (aq)

• Examples
• Na2S (aq) 2 HCl(aq) -gt 2 NaCl(aq) H2S(g)
• NH4Cl(aq) NaOH(aq) -gt NaCl(aq) NH3(g)

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Gas-forming reactions

• Conversion of ionic species into a covalently
  bonded molecule
• 2H --gt H2

• Example
• Zn (s) 2HCl(aq) -gt ZnCl2(aq) H2(g) (redox)
• http//video.google.com/videoplay?docid-213426665
  4801392897

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Oxidation numbers

• We can track which atoms are losing or gaining
  electrons by calculating their oxidation numbers
  before and after a chemical reaction
• Oxidation numbers are calculated like formal
  charges, with one very different assumption
• Electrons in a bond are shared unequally, with
  the more electronegative atom getting all of the
  bond electrons

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Oxidation numbers

• Rules for determining oxidation numbers
• The sum of oxidation states within a compound
  must equal its charge.
• Each atom in a pure element has an oxidation
  number of zero (elements are neutral)
• For monoatomic ions, the oxidation state is equal
  to the ion charge

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Oxidation numbers

• Rules for determining oxidation numbers
• The preferred oxidation number of an element in a
  compound is generally the one the gives it a full
  shell of electrons (same as its ionic charge)
• Halogens (F,Cl,Br,I) prefer -1
• Oxygen prefers -2
• Alkali metals prefer 1
• Alkaline earth metals prefer 2
• Lanthanides, Sc, Y, and La prefer 3
• Transition metals have multiple possible
  oxidation states

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Oxidation numbers

• Notable exceptions
• When two elements both want a negative oxidation
  state, only the more electronegative one gets it.
• (Cl, Br, I have oxidation numbers of -1, except
  when combined with O or F)
• The oxidation number of H is 1 with non-metals
  (C,O,N,F,S,P, etc.), but is -1 when combined with
  metals (Na, Mg, Ca, Fe, etc.)
• Need to calculate the oxidation number of
  transition metals last!

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Test cases

• O2 gas
• MnO2
• Mn2O3
• KMnO4
• Ba(ClO3)2
• LaRuO3

O 0 (neutral element) O -2 Mn must be 4 O
-2(x3) Mn must be 6/2 3 O -2(x4) K
1 Mn 7 O -2 Ba 2 Cl must be 5 O
-2 La 3 Ru must be 3
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Redox reactions

• A redox reaction is one in which oxidation
  numbers of elements change during the reaction
• Short for reduction-oxidation reaction
• Reduction when oxidation number gets reduced
• Oxidation when oxidation number gets increased

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Redox reactions

• Examples include
• Fe2O3(s) 3 CO(g) -gt 2 Fe(s) 3 CO2(g)
• 2 Ag(aq) Cu(s) -gt 2 Ag(s) Cu2(aq)
• In each case, electrons are shifted around

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-6 (for molecule)
-4 (for molecule)
3
2
4
0
-2
-2
-2
Fe 3 -gt 0 C 2 -gt 4
Ag 1 -gt 0 Cu 0 -gt 2
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Redox reactions

• Fe2O3(s) 3 CO(g) -gt 2 Fe(s) 3 CO2(g)
• 2 Ag(aq) Cu(s) -gt 2 Ag(s) Cu2(aq)
• When an element is reduced, its oxidation number
  is reduced (Fe 3 -gt 0).
• When an element is oxidized, its oxidation number
  is increased (C 2 -gt 4)
• A reducing agent will reduce the charge on
  another species (while being oxidized itself,
  since it is donating electrons)
• An oxidizing agent will oxidize another species
  (while being reduced itself, since it taking away
  electrons).
• Most common oxidizing agent

oxygen
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Demonstrations

• H2 O2 --gt H2O
• Zn HCl --gt ZnCl H2
• C3H8 O2 --gt CO2 H2O
• (NH4)2Cr2O7 --gt N2 4 H2O Cr2O3
• Fe2O3 Al --gt Fe Al2O3
• Gummy bear (on CD)

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Common oxidizing agents

• oxygen Mg O2 -gt MgO

• halogens Mg Cl2 -gt MgCl2

• nitric acid 2NO3- Cu 4H
• -gt Cu2 2NO2 H2O

• dichromate 3 CH3CH2OH 2 Cr2O72- 16 H
• -gt 3 CH3COOH 4 Cr3 11 H2O

• permanganate MnO4- 5 Fe2 8 H
• -gt Mn2 5 Fe3 4 H2O

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Common reducing agents

• hydrogen Mn2O3 3 H2 -gt 2Mn 3 H2O

• elemental metals 2 Na 2H2O -gt 2NaOH H2

• carbon/CO Fe2O3 3 CO -gt Fe 3 CO2