7:26 PM Unit 1: Stoichiometry Chemistry 2202 1. 7:26 PM Stoichiometry Stoichiometry deals with quantities used in OR produced by a chemical reaction 2.
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02:37 PM
Unit 1: Stoichiometry
Chemistry 2202
1
02:37 PM
Stoichiometry
Stoichiometry deals with quantities used in OR produced by a chemical reaction
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3 Parts
Mole Calculations (Chp. 2 & 3) Stoichiometry and Chemical
Equations (Chp. 4) Solution Stoichiometry (Chp. 6)
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PART 1 - Mole Calculations
Isotopes and Atomic Mass (pp. 43 - 46)
Avogadro’s number (pp. 47 – 49) Mole Conversions (pp. 50 - 74)
M, MV, NA, n, m, v, N
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Questions
p. 45 #’s 1 – 4p. 46 #’s 1 – 6p. 75 #’s 9 – 12p. 51-53 #’s 5 – 15p. 57 #’s 16 – 19p. 59,60 #’s 20 –
27p. 63,64 #’s 28 - 37
p. 54 #’s 5 – 8p. 65 #’s 2, 4, 5p. 75 #’s 13, 14p. 76 #’s 15, 17–
19,
21-23p. 73 #’s 38 – 43p. 74 #’s 1 – 4p. 76 #’s 26, 27
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PART 1 - Mole Calculations
Percent composition: - given mass (p. 79 - 82) - given the chemical formula (p. 83 - 86)
Empirical Formulas (pp. 87 - 94) Molecular Formulas (pp. 95 - 98) Lab: Formula of a Hydrate
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Questions
p. 82 #’s 1 – 4p. 85 #’s 5 – 8p. 89 #’s 9 – 12p. 91 #’s 13 – 16p. 97 #’s 17 - 20
p. 103 #’s 23 – 24p. 86 #’s 1, 3 – 6p. 94 #’s 1 - 7p. 106 #’s 1 - 3,
6, 7p. 107 – 109
#’s 5 – 23, 25
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Isotopes and Atomic Mass
atomic number - the number of protons in an atom or ion
mass number - the sum of the protons and neutrons in an atom
isotope - atoms which have the same number of protons and electrons but different numbers of neutrons
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Isotopes and Atomic Mass eg.
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Isotopes and Atomic Mass
1735
1737C l C l
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Isotopes and Atomic Mass
1224
1225
1226M g M g M g
not all isotopes are created equal
79 % 10 % 11 %
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Isotopes and Atomic Mass
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Isotopes and Atomic Mass
atomic mass unit (AMU - p.43)- a unit used to describe the mass of
individual atoms- the symbol for the AMU is u- 1 u is 1/12 of the mass of a carbon-12
atom
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Isotopes and Atomic Mass
average atomic mass (AAM)- the AAM is the weighted average of
all the isotopes of an element (p. 45)
p. 14 # 5p. 45 #’s 1 – 4p. 46 #’s 1 – 6 p. 75 #’s 9 - 12 14
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Finding % Abundance
eg. Br has two naturally occurring isotopes. Br-79 has a mass of 78.92 u and Br-81 has a mass of 80.92 u. If the AAM of Br is 79.90 u, determine the percentage abundance of each isotope.
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Let x = fraction of Br-79Let y = fraction of Br-81
x + y = 178.92x + 80.92y = 79.90
Finding % Abundance
y = 1 - x
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x + y = 178.92x + 80.92y = 79.90
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Avogadro’s Number (p. 47)
p. 48
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Avogadro’s Number The MOLE is a number used by chemists to
count atoms The MOLE is the number of atoms contained in
exactly 12 g of carbon-12. In honor of Amedeo Avogadro, the number of
particles in 1 mol has been called Avogadro’s number.
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How big is Avogadro's number?
An Avogadro's number of soft drink cans would cover the surface of the earth to a depth of over 200 miles.
Avogadro's number of unpopped popcorn kernels spread across the USA, would cover the entire country to a depth of over 9 miles.
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If we were able to count atoms at the rate of 10 million per second, it would take about 2 billion years to count the atoms in one mole.
How big is Avogadro's number?
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Avogadro’s Number
1 mole = 6.02214199 x 1023 particles
1 mol = 6.022 x 1023 particles
NA = 6.022 x 1023 particles/mol
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Avogadro’s Number
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Avogadro’s Number
Number of moles Number of atoms
5 mol
0.01 mol
4.65 x 1024 atoms
8.01 x 1021 atoms
7.72 mol
0.0133 mol
6.022 x 1021 atoms
3.011 x 1024 atoms
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Avogadro’s Number
Formulas:
AN
Nn n = # of moles
N = # of particles (atoms, ions, molecules, or formula units)
NA = Avogadro’s #
N = n x NA
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How many moles are contained in the following?
a) 2.56 x 1028 Pb atoms
b) 7.19 x 1021 CO2 molecules
Avogadro’s Number
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Avogadro’s Number
eg. Calculate the number of moles in 4.98 x 1025 atoms of Al.
eg. How many formula units of Na2SO4 are in 5.69 mol of Na2SO4?
# of Na ions? # of Oxygen atoms?
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Avogadro’s Number
1. How many molecules of glucose are in 0.435 mol of C6H12O6?
How many carbon atoms?2. Calculate the number of moles in
a sample of glucose that has 3.56 x 1022 hydrogen atoms.
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Avogadro’s Number
pp. 51 – 53: #’s 5 – 15p. 54: #’s 4 - 8
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Molar Mass
The mass of one mole of a substance is called the molar mass of the substance
eg. 1 mole of Pb has a mass of 207.19 g1 mole of Ag has a mass of 107.87 g
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Molar Mass
The symbol for molar mass is M and the unit is g/mol
eg. MPb = 207.19 g/mol
MAg = 107.87 g/mol
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Molar Mass
The molar mass of a compound is the sum of the molar masses of the elements in the compound
eg. Calculate the molar mass of:a) H2O b) C6H12O6 c) Ca(OH)2
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Molar Mass
H2O has 2 hydrogens and 1 oxygen
2 x 1.01 = 2.021 X 16.00 = 16.00
18.02 g/mol
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Molar Mass
C6H12O6
6 x 12.01 = 72.0612 x 1.01 = 12.126 x 16.00 = 96.00
180.18 g/mol
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Molar Mass
Ca(OH)2
1 x 40.08 = 40.082 x 16.00 = 32.002 x 1.01 = 2.02
74.10 g/mol
Your calculator may not show the zeroes.There should be 2 digits after the decimal when adding molar masses
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Molar Mass
p. 57: #’s 16 – 19 & Molar Masses Handout
1. 151.92 g/mol 7. 58.44 g/mol 2. 120.38 g/mol 8. 100.09 g/mol 3. 105.99 g/mol 9. 44.02 g/mol 4. 100.40 g/mol 10. 248.22 g/mol 5. 74.44 g/mol 11. 115.04 g/mol 6. 78.01 g/mol
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Molar Mass Calculations
AN
Nn
M
mn
N = n x NA
mass
molar
mass
m = n x M
Avogadro’s #
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Molar Mass Calculations
AN
Nn
M
mn
N = n x NA
m = n x M
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Molar Mass Calculations
eg. How many moles are in 25.3 g of NO2?
m = 25.3 gMNO2 = 46.01 g/mol
M
mn
mol/g.
g.
0146
325
mol.5500
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Molar Mass Calculations
eg. What is the mass of 4.69 mol of water?
n = 4.69 molMwater = 18.02 g/mol
m = n x M = (4.69 mol)(18.02 g/mol) = 84.5 g
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Molar Mass Calculations
Practice: p. 59 #’s 20 - 23p. 60 #’s 24 - 27
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The Mole #4answers
1.a) 0.038 mol 2.a) 17.4 g 3.a) 5631 g b) 3.75 mol b) 1560. g b) 3.73 g c) 0.276 mol c) 528 g c) 10.02 g d) 23 mol d) 97513 g d) 1662.9 g e) 0.0575 mol e) 9328 g e) 981.2 g
4.a) 1.82 mol c) 0.02133 mol e) 0.0000573 mol b) 13 mol d) 4.3423 mol
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eg. How many molecules are in 26.9 g of water?m = 26.9 gMwater= 18.02 g/mol
NA = 6.022 x 1023 molecules/mol
Find N
Particle–Mole-Mass Conversions
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M
mn
02.18
9.26
= 1.493 mol H2O
N = n x NA
= 1.493 X 6.022 x 1023
= 8.99 x 1023 molecules
Particle–Mole-Mass Conversions
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eg. A sample of Sn contains 4.69 x 1028 atoms. Calculate its mass.
N = 4.69 x 1028 NA = 6.022 x 1023 molecules/mol
MSn = 118.69 g/mol
Find m
Particle–Mole-Mass Conversions
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AN
Nn
23
28
10x022.6
10x69.4
= 77,881 mol
m = n x M
= 77881 mol x 118.69 g/mol
= 9.24 x 10 6 g
Particle–Mole-Mass Conversions
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1. Calculate the mass of 4.80 x 1024 water molecules.
2. How many grams are in 2.53 x 1026 CH4 molecules?
3. How many atoms are in 68.0 g of Sn?4. Calculate the number of molecules
in 105 g of C6H12O6.
Particle–Mole-Mass Conversions
47
Particles (N) Moles (n) Mass (m)
4.80 x 1024 H2O molecules
2.53 x 1026 CH4 molecules
68.0 g of Sn
105 g of C6H12O6
Particle–Mole-Mass Conversions
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particles(N)
Moles(n)
Mass(m)
5.98 x 1026 Cu atoms
4.50 g H2O
6.15 mol O3
Particle–Mole-Mass Conversions
02:37 PM 49
particles(N)
Moles(n)
Mass(m)
4.18 x 1022 Al atoms
2.45 g C8H18
0.145 mol S2F4
Particle–Mole-Mass Conversions
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Practice: p. 63 #’s 28 - 33p. 64 #’s 34 – 37p. 76 # 15
moles (n)
mass (m)
particles(N)
x M
÷ M÷ NA
x NA
Particle–Mole-Mass Conversions
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eg. How many molecules are in 4.78 g of glucose?m = 4.78 gMwater= 180.18 g/mol
NA = 6.022 x 1023 molecules/mol
Find N
Particle–Mole-Mass Conversions
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M
mn
180.18
4.78
= 0.02653 mol glucose
N = n x NA
= 0.02653 X 6.022 x 1023
= 8.99 x 1023 molecules
Particle–Mole-Mass Conversions
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Practice: p. 63 #’s 28 - 33p. 64 #’s 34 – 37p. 76 # 15
Particle–Mole-Mass Conversions
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Molar Mass Calculations
Practice:p. 54 #’s 5 - 8p. 65 #’s 2, 4, 5p. 75 #’s 13, 14, p. 76 #’s 15, 17 – 19, 21 -23
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Molar Volume
•The volume of a gas increases when temperature increases but decreases when pressure increases .
•The volume of gases is measured under conditions of Standard Temperature and Pressure (STP)
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Molar Volume
Standard Pressure – 101.3 kPa Standard Temperature – 0 °C Avogadro hypothesized that equal
volumes of gases at the same temperature and pressure contain equal numbers of molecules.
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Molar Volume
Experimental evidence shows the volume of one mole of ANY GAS at STP is 22.4 L/mol
OR VSTP = 22.4 L/mol
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Molar Mass Calculations
AN
Nn
M
mn
N = n x NA
m = n x M
STPV
vn
v = n x VSTP
given volume in Litres
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Molar Mass Calculations
moles (n)
mass (m)
particles(N)
x M
÷ M÷ NA
x NA
volume (v)
x VSTP ÷ VSTP
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Molar Volume
p. 73 #’s 38 – 43
p. 74 #’s 1 – 4
p. 76 #’s 26, 27
WorkSheet: The Mole #6
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Percent Composition (p. 79)
The mass percent of a compound is the mass of an element in a compound expressed as a percent of the total mass of the compound.
mass percen tm ass o f elem en t
to ta l m ass o f com poundX 100%
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Percent Composition
eg. 8.50 g of a compound was analyzed and found to contain 6.00 g of hydrogen and 2.50 g of carbon. Calculate the mass percent for each element.
p. 82 #’s 1 - 466
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Percent Composition
mass percent may be found using the formula & the molar mass of a compound.
eg. Find the percentage composition for CH4
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Percent Composition
M = 12.01 g/mol + 4(1.01 g/mol) = 12.01 g/mol + 4.04 g/mol
= 16.05 g/mol
%25.2
100x16.05
4.04H%
%74.8
100x16.05
12.01C%
p. 85 #’s 5 - 8
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p. 85 #’s 5 - 8p. 86 #’s 1, 3 – 6p. 107 #’s 6 – 10
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Empirical Formulas
An empirical formula gives the simplest ratio of elements in a compound.A molecular formula shows the actual number of atoms in a molecule of a compound.Ionic compounds are always written as empirical formulas
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Empirical Formulas
Compound Molecular Formula
Empirical Formula
butane C4H10
glucose C6H12O6
water H2O
benzene C6H6
C2H5
CH2O
H2O
CH72
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Empirical Formulas (p.87)
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Empirical Formulas
The empirical formula of a compound may be determined by using the % composition of a given compound.
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Empirical Formulas
Method: assume you have 100.0 g of the
compound (ie. change % to g) calculate the moles (n) for each
element divide each n by the smallest n to get
the ratio for the empirical formula75
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Empirical Formulas
eg. A compound was analyzed and found to contain 87.4% N and 12.6 % H by mass. Determine the empirical formula of the compound.
p. 89 #’s 9 - 12
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Empirical Formulas
When finding the EF, the mole ratio may not be a whole number ratio.
eg. A compound contains 84.73% N and 15.27 % H by mass. Determine the empirical formula of the compound.
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p. 90
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Empirical Formulas
eg. A compound contains 89.91% C and 10.08 % H by mass. Determine the empirical formula of the compound.
p. 91 #’s 13 – 16p. 94 #’s 2-4, 6, 7
Answers on p. 109
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MgO Lab
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Molecular Formulas The molecular formula of a compound is
a multiple of the empirical formula.
See p. 95
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Molecular Formulas
To find the molecular formula we need the empirical formula and the molar mass of the compound
eg. The empirical formula of hydrazine is NH2. The molar mass of hydrazine is 32.06 g/mol. What is the molecular formula for hydrazine?
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Molecular Formulas
p. 97 #’s 17 – 20
p. 107, 108 #’s 11 - 14
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CHC analyzer (p. 99 – 101)
1. Describe the operation of a carbon hydrogen combustion analyzer.
2. 22.0 g of carbon dioxide and 10.8 g of water is collected in a CHC analysis. Determine the empirical formula of the hydrocarbon.
p. 101 #’s 21, 2284
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Formula of a hydrate
To determine the formula of a hydrate:
- calculate the moles of water
- calculate the moles of anhydrous compound
- determine the simplest ratio
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Formula of a hydrate
eg. Use the data below to determine the value of x in LiCl• xH2O.
mass of crucible = 26.35 g
crucible + hydrate = 42.15 g
crucible + anhydrous compound= 34.94 g
86
mwater = 42.15 – 34.94 = 7.21 g H2O
mLiCl = 34.94 – 26.35 = 8.59 g LiCl
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eg. Na2CO3. xH2O
crucible = 15.96 g
crucible + hydrate = 22.19 g
crucible + anhydrous compound = 19.67 g
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eg. CoCl2.xH2O
crucible = 151.96 gcrucible + hydrate = 164.35 g crucible + anhydrous compound = 158.23 g
p. 103; # 24 Lab: pp. 104-105
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Formula of a hydrate
mass of empty beaker
mass of beaker & hydrate
mass of beaker & anhydrous compound
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Review – Chp. 3
p. 86 #’s 1, 3 – 6p. 94 #’s 1 – 7p. 103 # 23pp. 107–109 #’s 5 – 19, 21, 22
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Test
p. 45 #’s 1 – 4p. 46 #’s 1 – 6p. 75 #’s 9 – 12p. 51-53 #’s 5 – 15p. 57 #’s 16 – 19p. 59,60 #’s 20 –
27p. 63,64 #’s 28 - 37
p. 54 #’s 5 – 8p. 65 #’s 2, 4, 5p. 75 #’s 13, 14p. 76 #’s 15, 17–
19,
21-23p. 73 #’s 38 – 43p. 74 #’s 1 – 4p. 76 #’s 26, 27
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Test
p. 82 #’s 1 – 4p. 85 #’s 5 – 8p. 89 #’s 9 – 12p. 91 #’s 13 – 16p. 97 #’s 17 - 20
p. 103 #’s 23 – 24p. 86 #’s 1, 3 – 6p. 94 #’s 1 - 7p. 106 #’s 1 - 3,
6, 7p. 107 – 109
#’s 5 – 23, 25
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Test Friday – Dec. 4
1. Formula of a hydrate- lab activity- p. 103 # 24
2. % composition- given mass
p. 82 #’s 1 – 4 - given formula
p. 85 #’s 5 - 8p. 86 #’s 3 & 4
p. 103 # 23p. 107 # 5
3. EF and MF- What is an empirical formula- Finding EF from % composition pp. 89, 91 #’s 9 - 16 (watch for .5 or .333)- Finding MF from EF and molar mass p. 97 #’s 17 – 20 p. 108 #13
4. Mole calculations – Chp. 2
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Stoichiometry (Chp.4)
Stoichiometry is the determination of quantities needed for, or produced by, chemical reactions.
Ratios from balanced chemical equations are used to predict quantities.
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Stoichiometry – p. 111
Clubhouse sandwich recipe
96
Clubhouse sandwich recipe
Slices of Toast
Slices of Turkey
Strips of Bacon
# of Sandwiche
s
12
27
66
100
Fill in the missing quantities:
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Mole Ratios
A mole ratio is a mathematical expression that shows the relative amounts of two species involved in a chemical change.
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Mole Ratios
A mole ratio Comes from a balanced chemical
equation Shows the relative amounts of the
reactants/products in moles Looks like
99
coeffic ien t
coeffic ien trequ ired
g iven
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N2(g) + 3 H2 (g) → 2 NH3 (g)
20
66
140
81
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C3H8 + 5 O2 → 3 CO2 + 4 H2O
n n xcoeffic ien t
coeffic ien trequ ired g ivenrequ ired
g iven
How many moles of CO2 are produced when 31.5 mol of O2 react?
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C3H8 + 5 O2 → 3 CO2 + 4 H2O
How many moles of H2O are produced when 1.35 mol of O2 react?
02:37 PM 102
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C3H8 + 5 O2 → 3 CO2 + 4 H2O
How many moles of C3H8 are needed to react with 0.369 mol of O2?
02:37 PM 103
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C5H12 + O2 → CO2 + H2O
How many moles of CO2 are produced when 6.35 mol of O2 react?
02:37 PM 104
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Al(s) + Br2(l) → AlBr3(s)
How many moles of Br2 are needed to produce 0.315 mol of AlBr3?
p. 115 #’s 4 – 7p. 117 #’s 8 - 10
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Mole to Mole Stoichiometry
1. How many moles of copper would be produced if 20.5 mol of copper (II) oxide decomposes?
02:37 PM 106
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Mass to Mole Stoichiometry
1. How many moles of water are produced when 20.6 g of CH4 burns?
02:37 PM 107
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Mass to Mole Stoichiometry
2. How many moles of nitrogen gas are needed to produce 6.75 g of NH3 in a reaction with hydrogen gas?
108
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Mass to Mole Stoichiometry
3. How many moles of silver would be produced if 10.0 g of silver nitrate reacts with copper metal?
109
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Mass to Mole Stoichiometry
4. How many moles of water are produced when 20.6 g of C6H12 burns?
110
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Mass to Mole Stoichiometry
5. How many moles of water are produced when 20.6 g of CH4 burns?
02:37 PM 111
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Mass to Mole Stoichiometry
02:37 PM 112
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Mole to Mass Stoichiometry
1. What mass of CaCl2 is produced when
4.38 mol of Ca(NO3)2 reacts with NaCl?
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Mole to Mass Stoichiometry
2. What mass of copper would be produced if 20.5 mol of CuO decomposes?
02:37 PM 114
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Four step stoichiometry
1. Write a balanced chemical equation
2. Calculate moles given
3. Mole ratio – find moles required
4. Calculate required quantity
m = n x M v = n x Vstp N = n x NA
115
M
mn
STPV
vn
AN
Nn
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Mass to Mass Stoichiometry
1. How many grams of water are produced when 5.45 g of C3H8 burns?
02:37 PM 116
02:37 PM 117
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Mass to Mass Stoichiometry
eg. Calculate the mass of HCl needed to react with 3.56 g of Fe to produce FeCl2.
02:37 PM 118
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Stoichiometry (Chp.4)
eg. What mass of CO2 gas is produced when 45.9 g of CH4 burns ?
Step #1
CH4 + 2 O2 → CO2 + 2 H2O
45.9 g ? g
119
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eg. How many moles of HCl needed to react with 3.56 g of Fe to produce FeCl2.
121
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Mole Calculations (p. 121 #13)
3.56 g ? g
= 0.06374 mol Fe
Fe + 2 HCl → FeCl2 + H2
55.85g/mol
g3.56nFe Step #2
Step #3
Fe mol 1
HCl mol 2 x Fe mol 0.06374nHCl
= 0.12748 mol HCl122
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Mole Calculations
p. 122 #15Given 32.0 g of sulfur (M = 256.56 g/mol) Find mass of ZnS
#2 n = 0.1247 mol S8
#3 n = 0.9976 mol ZnS(M = 97.45 g/mol)
#4 m = 97.2 g ZnS123
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Mole Calculations
p. 123 #18Given 33.5 g of H3PO4 (M = 98.00 g/mol) Find mass of MgO
#2 n = 0.3418 mol H3PO4
#3 n = 0.5128 mol MgO(M = 40.31 g/mol)
#4 m = 20.7 g MgO124
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Mole Calculations
p. 123 #17Given 25.0 g of Al4C3 (M = 143.95 g/mol) Find volume of CH4
#2 n = 0.174 mol Al4C3
#3 n = 0.522 mol CH4 (MV = 22.4 L/mol)
#4 m = 11.7 L CH402:37 PM 125
How many moles of aluminum chloride can be produced from the reaction of chlorine and 10.8 mol of aluminum ?
Cl2(g) + Al(s) → AlCl3(s)
02:37 PM 126
How many moles of magnesium are needed to react with 27 g of iodine to form magnesium iodide?
02:37 PM 127
How many grams of nitrogen are needed to react with 14.0 mol of oxygen to produce nitrogen dioxide ?
N2(g) + O2(g) → NO2(g)
02:37 PM 128
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Mole Calculations (p. 121 #14)
2.34 g ? L
#2 n = 0.05086 mol NO2
#3 n = 0.01272 mol O2
(MV = 22.4 L/mol)
#4 n = (0.01272)(22.4) = 0.285 L O2129
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Limiting Reactant (p. 128)
10.0 g of Li reacts with 15.0 g of Br2. Calculate the mass of LiBr produced.
130
Limiting Reactant (p. 128)
02:37 PM 131
2 Li Br2 2 LiBr
10.0
15.0
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Limiting Reactant (p. 128)
The Limiting Reactant (LR) OR Limiting Reagent (LR) is the substance that is completely used in a chemical reaction.
The Excess Reactant is the reactant that is left over after a reaction is complete.
132
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Limiting Reactant (p. 128)
eg. 2.00 g of NaI reacts with 2.00 g of Pb(NO3)2. Determine the LR and calculate the amount of PbI2 produced.
write a balanced equation find n for each reactant (Step #2) find moles produced by each
reactant (Step #3) 133
Pb(NO3)2 + 2 NaI → 2 NaNO3 + PbI2
2.00 g 2.00 g
nPb(NO)3 = 2.00 g 331.21 g/mol = 0.006038 mol
nNaI = 2.00 g 149.89 g/mol = 0.013343 mol
02:37 PM 134
nPbI2 = 0.006038 mol Pb(NO3)2 x 1 mol PbI2
1 mol Pb(NO3)2
= 0.006038 mol PbI2
nPbI2 = 0.013343 mol NaI x 1 mol PbI2
2 mol NaI
= 0.006672 mol PbI2
mPbI2 = 0.006038 mol x 460.99 g/mol
= 2.78 g PbI2
02:37 PM 135
What mass of calcium carbonate will be produced when 20.0 g of calcium phosphate reacts with 15.0 g of sodium carbonate?
(14.2 g)
3 Na2CO3 + Ca3(PO4)2 → 3 CaCO3 + 2 Na3PO4
02:37 PM 136
What mass of barium hydroxide will be produced when 10.0 g of barium nitrate reacts with 30.0 g of sodium hydroxide? (6.56 g)
Ba(NO3)2 + NaOH → Ba(OH)2 + NaNO3
= 0.03826 mol Ba(NO3)2 = 0.750 mol NaOH
02:37 PM 137
g/mol 40.00
g30.00nNaOH g/mol 261.35
g10.00n
23 )Ba(NO
Using Ba(NO3)2
= 0.03826 mol Ba(OH)2
Using NaOH
= 0.375 mol Ba(OH)2
02:37 PM 138
23
223Ba(OH) )Ba(NO mol 1
Ba(OH) mol 1 x )Ba(NO mol 0.03826n
2
NaOH mol 2
Ba(OH) mol 1 x NaOH mol 0.750n 2
Ba(OH)2
g 6.56
g/mol 171.35 x mol 0.03826m2Ba(OH)
What volume of hydrogen gas at STP will be produced when 10.0 g of zinc metal reacts with 20.0 g of hydrogen chloride?
Zn + 2 HCl → H2 + ZnCl2
02:37 PM 139
02:37 PM 140
02:37 PM 141
Law of Conservation of Mass (p. 118)
In a chemical reaction, the total mass of reactants always equals the total mass of products.
eg. 2 Na3N → 6 Na + N2
When 500.00 g of Na3N decomposes 323.20 g of N2 is produced. How much Na is produced in this decomposition?
142
02:37 PM
Law of Conservation of Mass( p. 118)
eg. To produce 90.1 g of water, what mass of hydrogen gas is needed to react with 80.0 g of oxygen?
eg. If 3.55 g of chlorine reacts with exactly2.29 g of sodium, what mass of NaCl willbe produced?
143
02:37 PM
The theoretical yield is the amount of product that we calculate using stoichiometry
The actual yield is the amount of product obtained from a chemical reaction
Percent yield (p. 137)
144
02:37 PM
Percent yield (p. 137)
percent y ie ld actua l y ie ld
theoretica l y ie ld x 100
145
p. 139 #’s 31, 32, & 33
DEMO: silver nitrate + copper
Equation:
Mass AgNO3 =
Mass Cu =
02:37 PM 146
DEMO: silver nitrate + copper
Mass of filter paper and precipitate =
Mass of empty filter paper =
Mass of precipitate =
02:37 PM 147
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