CHEM 120: Introduction to Inorganic Chemistry

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CHEM 120 Introduction to Inorganic Chemistry

• Instructor Upali Siriwardane (Ph.D., Ohio State
  University)
• CTH 311, Tele 257-4941, e-mail
  upali_at_chem.latech.edu
• Office hours 1000 to 1200 Tu Th 800-900
  and 1100-1200 M,W, F

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Chapters Covered and Test dates

• Tests will be given in regular class periods 
  from  930-1045 a.m. on the following days
• September 22,     2004 (Test 1) Chapters 1 2
• October 8,         2004(Test 2)  Chapters  3,
  4
• October 20,         2004 (Test 3) Chapter  5 6
  
• November 3,        2004 (Test 4) Chapter  7 8
• November 15,      2004 (Test 5) Chapter  9 10
• November 17,      2004 MAKE-UP Comprehensive
  test (Covers all chapters
• Grading
• ( Test 1 Test 2 Test3 Test4 Test5)
  x.70 Homework quiz average x 0.30 Final
  Average
•                               5

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Chapter 5. Calculations and the Chemical Equation
5.1 The Mole Concept and Atoms The Mole and
Avogadro's Number Calculating Atoms, Moles, and
Mass 5.2 Compounds The Chemical Formula 5.3
The Mole Concept Applied to Compounds 5.4 The
Chemical Equation and The Information It Conveys
A Recipe for Chemical Change Features of a
Chemical Equation The Experimental Basis of a
Chemical Equation 5.5 Balancing Chemical
Equations 5.6 Calculations Using the Chemical
Equation General Principles Use of Conversion
Factors Theoretical and Percent Yield
Pharmaceutical Chemistry The Practical
Significance of Percent Yield
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The mole concept and atoms

• In ch 1 we learned that 1 amu 1.661 x
  10-24 g
• So if the average mass of a gold atom is196.97
  amu x 1.661 x 10-24 g 3.27 x 10-22 g
  1 amu
• a very
  small no.

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• Where did I get the 196.97 anu for the mass of
  one Au atom?
• From the ____________________!!!
• If I write amu after these nos. it implies that I
  have the mass of ______ atom of that element (in
  amu).
• But 3.27 x 10-22 g is too small an amt to work
  with in the lab.
• What to do?
• Scale up to quantities that we can handle by

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Avogadros number

• Defining one mole (mol) as amt of substance that
  contains as many elementary entities (atoms,
  molecules, ions , etc) as there are in atoms in
  exactly 12 g of the carbon-12 isotope. This is
  determined experimentally and is...

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Useful relationship

• moles X g X/molar mass X

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Some problems

• How many atoms are there in 5.10 moles of sulfur?
  Whats the mass of 5.10 moles of S?
• How many moles of calcium atoms are in 1.16 x
  1024 atoms of Ca? How many grams?

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• Which of the following has more atoms 1.10g of
  hydrogen atoms or 14.7 g of chromium atoms?
• How many moles are in 0.040 kg Na?

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• Whats the mass, in grams, of one atom of
  potassium?
• One atom of some element has a mass of 1,45 x
  10-22 g. Identify the element.

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Compounds

• The chemical formula
• MgO (ion pair)
• H2O
• C12H22O11
• Ca3(PO4)2
• CuSO4.5H2O vs CuSO4

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The mole concept applied to compounds

• The formula weight of a species is the sum of
  atomic masses (amu) of the atoms in a species.
• Formula weight of NH3
• For an ionic compound
• MgF2

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Covalent Cpds ? Molecular weight ?moles ?
(Molar Mass) Ionic Cpds ? Formula weight ?
Formula units ? (Molar Mass)

• In general we talk about
• moles for covalent compounds
• formula units rather than moles of ionic
  compounds.

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Molar mass

• Mass of one mole of NH3
• Mass of 6.022 x 1023 molecules of NH3 is
• Mass of one molecule of NH3 is.

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• Mass of one mole of MgF2 is
• Mass of one formula unit of MgF2 is
• Mass of 6.022 x 1023 formula units of MgF2 is

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• Calc the molar mass of Ca(NO3)2.
• Calc the molar mass of a compound if 0.372 mol of
  it has a mass of 152g.

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• 5.38. How many grams of each are required to have
  0.100 mol of
• A. NaOH
• B. H2SO4
• C. C2H5OH
• D. Ca3(PO4)2

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• 5.40. How many moles are in 50.0 g of
• A. CS2
• B. Al2(CO3)3
• C. Sr(OH)2
• D. LiNO3

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• Calc the no. of C, H, and O atoms in 1.50 g of
  glucose (C6H12O6).
• What is the average mass of one C3H8 molecule?
• What is the mass of 5.00 x 1024 molecules of NH3?

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Law of conservation of mass

• Mass is neither created nor destroyed in an
  ordinary chemical rxn.
• Or the sum of the masses of the reactants is
  equal to the sum of the masses of the products

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• mercury oxygen ---gt mercury(II)oxide
• 10.03g ? 10.83g
• Easier to use symbols for chem eqns.

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• reactants ? products
• lhs rhs
• may indicate physical state by (s), (g), (l),
  (aq)-aqueous solution
• Remember that H2,N2,O2,F2,Cl2,Br2,I2 occur as
  diatomics in nature and are used as diatomics in
  chemical eqns

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• To balance Have to have same no of each kind of
  atom on both sides of the eqn. The bonding
  arrangement changes, but the no of each kind of
  atom doesnt change.

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(D)

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Chemical equations

• Indicate phases of reactants and products

Fe(s) 2 HBr(aq)
Conditions etc.
(s) solid
(g) gas
(l) liquid
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Balancing chemical eqns

• Use correct formulas for the reactants and
  products (if word eqn at start)
• Balance by putting coefficients (nos) in front of
  the formulas. You may not change the formulas!
  These coefficients are called the stoichiometric
  (measure of mass) coefficients.
• By convention use the lowest set of whole no.
  coefficients to balance.

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• Start by balancing elements that appear only once
  on each side of the equation
• Balance remaining elements
• Check your balanced equation!
• To predict products--do an experiment

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To balance

• hydrogen nitrogen ? ammonia
• 1. write the symbols for the species in the rxn

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• Now figure out how to get the same no of atoms of
  each kind on both sides by using whole no
  coefficients in front of the species.
• As H2 _N2 ? NH3, then
• H2 _N2 ? _ NH3, then
• _ H2 _N2 ? _ NH3
• Now have _ Hs, _Ns on both sides and the lowest
  set of whole no coefficients have been used. The
  equation is balanced.

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• 3H2 N2 ? 2NH3
• 3 mol of H2 reacts with 1 mol of N2 to form 2 mol
  of NH3
• 3 molecules of H2 reacts with 1 molecule of N2 to
  form 2 molecules of NH3
• 6H 2N reacts to give 6H and 2N
• 6g of H2 reacts with 28 g of N2 to form 34g of
  NH3
• Note that

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Balance

• C2H6 O2 ? CO2 H2O
• H2O2 g H2O O2
• C2H5OH O2 g CO2 H2O
• KOH H3PO4 g K3PO4 H2O
• N2O5 g N2O4 O2

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Balance

• NH4NO3 g N2O H2O
• NH4NO2 g N2 H2O
• Be2C H2O g Be(OH)2 CH4
• NH3 CuO g Cu N2 H2O

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Balance

• S2Cl2(s) NH3(g) g N4S4(s) NH4Cl(s) S8(s)

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Calculations using the chemical eqn

• Quantitative study of reactants and products in a
  chemical reaction
• How much product will be formed?
• How much reactant is needed?
• Use coefficients in a balanced equation to
  convert between moles of different substances in
  a chemical reaction.

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Chemical Reaction

• 3H2(g) N2(g) ------gt 2NH3(g)  
• 3 mol H2 (reactant) 1 mol N2 (reactant)
  consumed
• 3 mol H2 (reactant) 2 mol NH3 (products)
  produced
• 1 mol N2 (reactant) 2 mol NH3 (products)
  produced
• 3 x 2 (6) g H2 (reactant) 1x 28 (28)mol N2
  (reactant) consumed
• 3 x 2 H2 (6) (reactant) 2x 17 (34) NH3
  (products) produced
• 1 x 28 (28) g N2 (reactant) 2 x 17 (34) NH3
  (products) produced

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Chemical reactions

• Hydrogen reacts with nitrogen to form NH3.
• theoretically it is should be that when 6 g of
  hydrogen reacts completely with 28 g of
  nitrogen, 34 g of ammonia is formed.
• However in real chemical reactions actual __ g of
  hydrogen reacting with __ g of nitrogen, __ g of
  ammonia is produced need be experimentally
  determined.

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2H2 O2 g 2H2O
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2H2 O2 g 2H2O

• How many moles of H2 is needed to completely
  react with19.8 mol O2?
• How many moles of H2O are formed when 25.4 mol of
  H2 react?

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2H2 O2 g 2H2O

• How many moles of H2 react with 38 g of O2?
• What mass of H2O is formed when 59.0g of H2
  reacts completely with O2? How much O2 reacted
  in this case?

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Mass relationships in chemical equations

• Mole-to-mole conversions
• use mole ratios as conversion factors
• Mass-to-mole and mole-to-mass conversions
• use molecular weights as conversion factors
• Mass-to-mass conversions
• do in multiple steps

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General prescription
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Problems

• How many grams of Al2O3 can be produced from 15.0
  g Al?
• 4 Al(s) 3O2(g) g 2Al2O3(s)

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• C3H8 O2 ? CO2 H2O balance
• How many mol of O2 does it take to completely
  burn 7.0 mol of C3H8?
• How many mol each of CO2 and H2O are produced?
• How many grams of oxygen does it take to
  completely burn 25.0 g of C3H8?
• How many grams each of CO2 and H2O are produced
  when 25.0 g of C3H8 is burned?

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• A 4.00 g sample of Fe3O4 reacts with O2 to
  produce Fe2O3.
• 4Fe3O4(s) O2(g)g 6Fe2O3(s)
• Determine the no. of grams of Fe2O3 produced.

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Theoretical and percent yield

• How good an experimentalist are you?
• What if 100 of reactants are not converted to
  desired products?
• Frequently happens because of side reactions
  (other products), handling, etc.
• 100 amount is theoretical yield
• Amount obtained is actual yield

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• Theoretical yield - amount of product that would
  result if all limiting reagent gave only product
• Actual yield - the amount of product actually
  obtained from a reaction (almost always less than
  the theoretical yield)
• Percent yield - calculated by
• yield actual yield ? 100
• theoretical yield

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• Theoretical yield is what we calculate assuming
  100 conversion of reactants to products.
• In the combustion of 33.5g of C3H6, 16.1 g of H2O
  is isolated. What is the percent yield?

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• If the yield of Fe2O3 in problem was 90.0 what
  was the actual yield of Fe2O3?

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• A 3.5 g sample of water reacts with PCl3
  according to 3H2O PCl3 g H3PO3 3HCl.
• How many grams of H3PO3 are produced?