regents chemistry stoichiometry
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Atomic Mass Unit (amu)
• amu- unit used to for mass of an atom• amu of Oxygen? 16amu
• Why not in grams? – Mass of oxygen is really 2.7 × 10-23g– Atoms are too small, number is too bulky
02
Find the mass of the following atoms
1.Mg = ______
2.Li = ______
3.Cl = ______
4. Al = ______
5. Ca = ______
6. H = ______
24 amu
7 amu
35 amu
23 amu
20 amu
1 amu
Types of Elements
Monoatomic Element
•One atom of an element that is stable enough to stand on its own (VERY RARE)
– not bonded to anything
Diatomic Elements
•Elements whose atoms always travel in pairs (Br,I,N,Cl,H,O,F)
– Bonded to another atom of the same element.
Formula Mass
• mass of an atom, molecule, compound
• O2
• This means that mass of O2 = 2 × _______ amu = ________ amu
• Notice how we use rounded numbers
Subscript = tells you the total number of atoms in the compound/molecule.
16 32
FormulasFormulas
• Molecular – actual ratios of atoms in a molecule or compound.– ex: C4H10 or C6H12O6
• Empirical – simplest ratio of atoms in a compound– ex: C2H5 or CH2O
• Structural – shows the arrangement of the atoms in the actual compound
ExampleExample
• The empirical formula is CH and the molecular mass is 26. What is the molecular formula?
1.C2H2
2.C3H3
3.C4H4
Step 1: find mass of empirical
CH = 13amu
Step 2: Divide molecular mass by empirical mass
26 (given)/13 = 2
Step 3: multiply empirical formula by answer in step 2
CH × 2 = C2H2
ReviewReview- The molecular formula is C3H6. What is the empirical formula?
______________
- Which is an empirical formula?
1.C2H2
2.H2O
3.H2O2
4.C6H12O6
- What is the empirical formula of the compound whose molecular formula is P4O10?
1.PO
2.PO2
3.P2O
4.P8O20
Chemical EquationChemical Equation• Shows which bonds are broken and which
bonds are built. – Numbers of atoms on the left side must equal
number of atoms on the right side of the arrow– After the elements are correctly written, only
the coefficient can be changed. – No coefficient means there is only one
molecule
H2 + O2 H202 atoms H 2 atoms O 2 hydrogen atoms
1 oxygen atom
24 =
2 =
24 =
Fix this equation!Fix this equation!• Formation of salt from sodium and
chlorine gas.
Na + Cl2 → NaCl
Na Cl ClNa ClPRODUCES
Review equationsReview equationsMg + Cl2 MgCl2Ca + HCl CaCl2 + H2
Ca + H20 Ca(OH)2 + H2
Given the incomplete equation: 2N2O5(g)
Complete the balanced equation. 1.2N2(g) + 3H2(g)
2.2N2(g) + 2O2(g)
3.4NO2(g) + O2(g)
4.4NO(g) + 5O2(g)
Reaction TypesReaction Types
Four Basic Types
1.Synthesis
2.Decomposition
3.Single replacement (substitution)
4.Double replacement
SynthesisSynthesis
• Formation of only ONE product from two reactants, but not always.
Examples:
N2(g) + 3H2(g) 2NH3(g)
CH4(g) + 2O2(g) CO2(g) + 2H2O(l) because O2 combines with both metal and nonmetal to form two oxides.
DecompositionDecomposition
• One reactant breaks apart to form several products.
• AKA combustion if products are CO2 and H2O
Examples:
2H2O2(aq) 2H2O(l) + O2(g) hydrogen peroxide decomposes over time to leave behind water.
Single ReplacementSingle Replacement
• A more active metal replaces a less active metal in a compound, or vice versa.
Example
2Fe(s) + 6HCl(aq) 2FeCl3(aq) + 3H2(s)
what happens when a metal becomes corroded by an acid where an element is reacting with a compound.
Double replacementDouble replacement
• Reaction between aqueous compounds.
• Cations and anions switch position.
• If an insoluble precipitate forms, the reaction is an end reaction, otherwise an aqueous mixture of ions
ExampleAgNO3 (aq) + NaCl (aq) NaNO3(aq) + AgCl(s)
ReviewReview
1. Cl2(g) + 2NaBr(aq) 2NaCl(aq) + Br2(l)
2. FeCl3(aq) + 3NaOH(aq) Fe(OH)3(s) + 3NaCl(aq)
3. 2Mg(s) + O2(g) 2MgO(s)
4. H2CO3(aq) H2O(l) + CO3(g)
SRSR
DRDR
SS
D/CombustionD/Combustion
Gram Formula Mass (GFM)Gram Formula Mass (GFM)
AKA
•Molar Mass
•Atomic mass
•Molecular mass (only in covalent)
•Mass of formula in grams
Formula MassFormula Mass
• Sum of atomic masses in the molecule
• What is the formula mass (or molecular mass) of K2CO3?
atom Total of each
mass total
K 2 x 39 = 78
C 1 x 12 = 12
O 3 x 16 = 48
Total = 138
gram formula mass (GFM)gram formula mass (GFM)
• GFM – describes the mass of one mole of a compound– To find GFM, add individual GAM for each
element in the compound• H20 = 1.00 + 1.00 + 16.00 = 18.00 g
• (NH4)2SO4 = N = 2 × 14 = 28
H = 8 × 1 = 8
S = 1 × 32 = 32
O = 4 × 16 = 64
132 g
Gram formula Mass (molar Gram formula Mass (molar mass)= mass in gramsmass)= mass in grams
• Mass of 6.02 x 1023 particles (1 mole of particles).
• If you weigh 6.02 x 1023 particles (1 mole of particles) of K2CO3, the weight on the scale will be 138 grams
Compound Formula mass, amu
Mass of 6.02 x 1023 particles
(on scale)
K2CO3 138 138g
H2O 18 18g
O2 32 32g
Molecular MassMolecular MassH20 = 1.01 + 1.01 + 16.00 = 18.02 g/mol
(NH4)2SO4 = N = 2 × 14.00 = 28
H = 8 × 1.01 = 8.08
S = 1 × 32.00 = 32
O = 4 × 16.00 = 64
132.08 g/mol
Determine the molecular Determine the molecular Formula of the followingFormula of the following
HNO3
H N O
1×1 1×14 3×16
1 + 14 + 48
63amu
(NH4)2CO3
N H C O
2×14 8×1 1×12 3×16
28 + 8 + 12 + 48
96amu
% Composition% Composition(NH4)2CO3 96amu96amu
N H C O
28/96 8/96 12/96 48/96
×100 ×100 ×100 ×10029.17% 8.33% 12.50% 50.00%
Needs to total 100%
HNO3 63amu63amu
H N O
1/63 14/63 48/63
×100 ×100 ×100
1.59% 22.22% 76.19%
Needs to total 100%
% Given instead% Given instead
• Cmpd is 86% C and 14% H. What is the empirical formula?
C H
86/12 14/1
7.17 C 14 H
7.17/7.17 14/7.17
1 2
CH2CH2
% of water in Na2CO3 • 10H2O (formula mass = 286)?
H2O = 2 + 16 = 18amu
10 × 18am = 180
180/286 × 100 = 62.94%62.94%
RuleRule
1 mole (of molecules)1 mole (of molecules)
EqualsEquals
1 gram molecular mass1 gram molecular mass
EqualsEquals
6.02 x 106.02 x 1023 23 molecules molecules
Equals Equals
22.4 liters (for gases at STP)22.4 liters (for gases at STP)
ExampleExample• What is the formula mass of NO2?
46 gram
• What is the mass of 2 moles of NO2? 2 X 46 gram = 92 grams
• What is the mass of 12 x 1023 molecules of NO2? 12 x 1023 / 6.02 x 1023 = 2
2 x 46 = 92 grams
• What is the mass of 44.8 liters of NO2? 44.8 / 22.4 = 2 2 x 46 = 92 grams
DensityDensity
Density = mass / volume
Usually expressed for gases in grams/liter
Gram formula mass = density at STP (g/L) x 22.4 liters
A piece of aluminum has a mass of grams and a volume of ml. What is its density?
1.35 g/l x 22.4 l = 30.24 grams
Another exampleAnother example
Which gas has a density of 1.70g/l at STP?
1.F2
2.He
3.N2
4.SO2
GFM = density (g/l) x 22.L
= 1.7 g/l x 22.4 liters
GFM = 38.08 grams
Total the mass of each choice to find the answer.
Mole-Mole ProblemsMole-Mole Problems• Answers how many moles
of one element or compound react with a given number of moles of another element or compound.
• How many moles of Ca are needed to react completely with 6 moles of H2O in the following reaction:
Step 1: Balance equation
Step 2: Cross out molecules not needed.
Step 3: Write mole number on top of given formula and an x on the unknown
Step 4: Write mole number on bottom of formula from balanced equation
Step 5: set up proportions
Ca + HCa + H22O O Ca(OH) Ca(OH)22 + H + H2226x
1 mole 2 mole
Review/PracticeReview/Practice
• Given the reaction: CH4 + O2 CO2 + H2O– How many moles of oxygen
are needed for the complete combustion of 3.0 moles of CH4?
1.6.0 moles
2.2.0 moles
3.3.0 moles
4.4.0 moles
• What amount of oxygen is needed to completely react with 1 mole of CH4?
1. 2 moles
2. 2 atoms
3. 2 grams
4. 2 molecules
Avogadro’s numberAvogadro’s number
• measured to be approximately 6.022 x 1023 (to 4 s.f)
• Chemists use the mole in the same way that grocers use the dozen for groups of 12 and stationers use the ream for groups of 500.
• we can use the mole without being overly concerned about exactly how many objects it represents
MoleMole• smallest measurable mass of matter contains
trillions of atoms, so chemists use a unit of amount called the mole (abbreviated mol).– one mole is the number of atoms in 12 g of carbon-
12• one atom of tin-120 has a mass of 120 u, it follows that one
mole of tin-120 atoms will have ten times the mass of one mole of carbon-12 atoms, i.e. 120 g.
– the mass in grams of one mole of atoms of any element will be numerically equivalent to its atomic mass in g/mol.
Review Review
• Which gas sample contains a total of 3.0 x 1023 molecules?
1. 71g of Cl22. 2.0 g of H2
3. 14g of N2
4. 38g of F2
A sample of an unknown gas at STP has a density of 0.630 g/l. What is the gram molecular mass of this gas?
1. 2.81g
2. 14.1g
3. 22.4g
4. 63.0g
• Which quantity represents 0.500 mole at STP?
1. 22.4L of Ar
2. 11.2L of N2
3. 32.0L of H2
4. 44.8L of He
Using Avogadro’s numberUsing Avogadro’s number
• how many atoms are in a sample of silicon that has a mass of 5.23 g.
• Moles can be number of atoms or particles in a molecule
5.23 g silicon
x 1 mol silicon =0.186 mol
silicon
28.09 g silicon
0.186 mol silicon
x 6.022 x 1023 atoms silicon =1.12 x 1023 atoms
silicon
1 mol silicon
MoleMole
– mass of 1 atom = mass of a mole of atoms / 6.022 x 1023
– mass of 1 C atom = 12.01 g / 6.022 x 1023 C atoms
• mass of 1 C atom = 1.994 x 10-23 g
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