Transcript 120CH05 - Louisiana Tech University

CHEM 120: Introduction to
Inorganic Chemistry
Instructor: Upali Siriwardane (Ph.D., Ohio State
University)
CTH 311, Tele: 257-4941, e-mail:
[email protected]
Office hours: 10:00 to 12:00 Tu & Th ; 8:00-9:00
and 11:00-12:00 M,W,& F
Chapters Covered and Test dates
• Tests will be given in regular class periods from 9:30-10:45 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
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
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.
• 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
Avogadro’s 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…...
Useful relationship
• # moles X = #g X/molar mass X
Some problems
• How many atoms are there in 5.10 moles of
sulfur? What’s 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?
• 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?
• What’s 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.
Compounds
• The chemical formula:
•
MgO (ion pair)
•
H2O
•
C12H22O11
•
Ca3(PO4)2
•
CuSO4.5H2O vs CuSO4
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 =
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.
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.
• 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
• 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.
• 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
• 5.40. How many moles are in 50.0 g of
• A. CS2
• B. Al2(CO3)3
• C. Sr(OH)2
• D. LiNO3
• 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?
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
• mercury + oxygen ---> mercury(II)oxide
• 10.03g
?
10.83g
•
•
• Easier to use symbols for chem eqns.
• 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
• 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 doesn’t change.
Reactants (reagents)
aA+bB
Coefficent
Conditions
Products
cC+dD
Arrow indicates reaction.
Conditions can include:
Heat (specify temperature) (D)
Solvent
Other information
Chemical equations
• Indicate phases of reactants and products
Fe(s) + 2 HBr(aq)
H2O
FeBr2(aq) + H2(g)
Conditions etc.
(s) = solid
(l) = liquid
(aq) = aqueous,
dissolved in water
(g) = gas
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.
• 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
To balance
• hydrogen + nitrogen  ammonia
• 1. write the symbols for the species in the
rxn
•
• 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 _ H’s, _N’s on both sides and the
lowest set of whole no coefficients have
been used. The equation is balanced.
• 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
Balance
• C2H6 + O2  CO2 + H2O
• H2O2  H2O + O2
• C2H5OH + O2  CO2 + H2O
• KOH + H3PO4  K3PO4 + H2O
• N2O5  N2O4 + O2
Balance
• NH4NO3  N2O + H2O
• NH4NO2  N2 + H2O
• Be2C + H2O  Be(OH)2 + CH4
• NH3 + CuO  Cu + N2 + H2O
Balance
• S2Cl2(s) + NH3(g)  N4S4(s) + NH4Cl(s) + S8(s)
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.
Chemical Reaction
• 3H2(g) +
•
•
•
•
•
•
N2(g) ------> 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
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.
2H2 + O2  2H2O
2H2 + O2  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?
2H2 + O2  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?
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
General prescription
Problems
• How many grams of Al2O3 can be produced
from 15.0 g Al?
• 4 Al(s) + 3O2(g)  2Al2O3(s)
• 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?
• A 4.00 g sample of Fe3O4 reacts with O2 to
produce Fe2O3.
• 4Fe3O4(s) + O2(g) 6Fe2O3(s)
• Determine the no. of grams of Fe2O3
produced.
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
• 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
• 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?
• If the % yield of Fe2O3 in problem was
90.0% what was the actual yield of Fe2O3?
• A 3.5 g sample of water reacts with PCl3
according to :
3H2O + PCl3  H3PO3 + 3HCl.
• How many grams of H3PO3 are produced?