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# Stoichiometry

Chapter 11
by

## Sarah Gleason

on 15 April 2013

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#### Transcript of Stoichiometry

The study of quantitative relationships between the amounts of reactants and the amount of products formed in a chemical reaction
It is based on the Law of Conservation of Mass
The amount of that each reactant presents will determine how much the product presents Stoichiometry In terms of molecules, moles and mass Interpreting Chemical Reactions Before you can interpret information about a chemical reaction, they must be balanced!
In a balanced equation the coefficients represent the number of molecules
Also, in a balanced chemical equation the coefficients represent the number of moles
Ex: C3H8 + 5O2 ---> 3CO2 + 4H2O
1 molecules of C3H8 + 5 molecules of oxygen ---> 3 molecules of CO2 + 4 molecules of H2O
1 mole of C3H8 + 5 moles of oxygen ---> 3 moles of CO2 + 4 moles of H20 And the equation in terms of mass should be equal on the reactant and the product side - Law of Conservation of Mass
Ex: C3H8 + 5O2 ---> 3CO2 + 4H2O
Mass of Reactant: 204.1 g
Mass of Product: 204.1 g
204.1 grams = 204.1 grams Interpret the following chemical equations in terms of particles, moles, and mass
1. N2 + 3H2 ---> 2NH2
2. HCl + KOH ---> KCl + H2O This is a ratio between the numbers of moles of any two substances in a BALANCED chemical equation Mole Ratios Example: 2K + Br2 ---> 2KBr
Ratio: 2 mol K / 1 mol Br2
Ratio: 2 mol of K/ 2 mol KBr

Ratio 2: 1 mol Br2/ 2 mol K
Ratio 2: 1 mol Br2 / 2 mol KBr

Ratio 3: 2 mol KBr / 2 mol K
Ratio 3: 2 mol KBr / 1 mol Br2 Now you Try! Determine the possible mole ratios in the following:
1. 4Al + 3O2 ---> 2Al2O3
2. ZnO + HCl ---> ZnCl2 + H2O (Balance first!) This formula will help you convert your known number of moles to the unknown number:

# of mole known x moles of the unknown = moles of unknown
moles of known Stoichiometric Mole to Mole Conversion Example: A disadvantage of burning propane (C3H8) is that carbon dioxide is released as a product (CO2)
How many moles of CO2 are produced when 10.0 mol of C3H8 are burned in excess oxygen in a gas grill?
First - Write out the balanced chemical equation
C3H8 + O2 ---> CO2 + H2O
Balanced: C3H8 + 5O2 ---> 3CO2 + 4H20
Then - write the correct mole ratio of the known -C3H8- to the moles of the unknown - CO2
10.0 mol of C3H8 / ? mol of CO2
Now - set up your conversion factor (on board)

EQUALS - 30.0 moles of CO2 NOW it's your TURN! 1. Methane and sulfur react to form carbon disulfide (CS2)
___ CH4 + ____S8 ---> ____CS2 + ______H2S
A. Balance the equation
B. Calculate the number of moles of CS2 produced from 1.50 mol of S8
C. How many moles of H2S are produced? If you know the moles, but you want to find the mass of another product or reactant Stoichometric Mole to Mass Conversions Example: Determine the mass of NaCl, commonly called table salt, produced when 1.25 mol of Chlorine gas (Cl2) reacts vigorously with excess sodium.
You know the moles of the reactant (Cl2) and now must determine the mass of the product (NaCl)
FIRST you must write out the BALANCED chemical equation
2Na + Cl2 ---> 2NaCl
SECOND you must convert moles of Cl2 to moles of NaCl using the mole ratio equation.
2 moles of NaCl / 1 mol of Cl2 = mol ratio
THIRD set up using equation of the known moles vs. unknown moles
1.25 mol Cl2 = 2 mol NaCl / 1 mol Cl2
= 2.50 mol NaCl
LAST convert from moles to mass - (use the molar mass of NaCl)
2.50 mol NaCl = 58.44 g NaCl / 1 mol Nacl
= 146 g NaCl YOUR TURN!! Titanium tetrachloride (TiCl4) is extracted from Titanium Oxide (TiO2) using chlorine (Cl2) and coke (carbon)
TiO2 + C + 2Cl2 ---> TiCl4 + CO2
A. What mass of Cl2 gas is needed to react with 1.25 mol of TiO2?
B. What mass of C is needed to react with 1.25 mol of TiO2?
C. What is the mass of ALL the products formed by a reaction with 1.25 mol of TiO2? ANSWERS:
A = 177 g Cl2
B = 15.0 g C
C = 292 g (of all the products) This can be used when preparing for a lab - in knowing the how much of each reactant to use when producing a certain mass of your product Mass to Mass Stoichiometry Example: Ammonium Nitrate (NH4NO3) produces N2O and H2O when it decomposes. Determine the mass of H2O produced from the decomposition of 25.0 g of solid NH4NO3
FIRST Write out the balanced equation
NH4NO3 --> N2O + 2H2O
Then convert from mass to moles
Answer will be .312 mol of NH4NO3
The next conversion is from moles to moles
Answer will be .624 mol of H2O
Last step is to convert from moles to mass
A. 2CH4 + S8 ---> 2CS2 + 4H2S
B. 3.00 mol CS2
C. 6.00 mol H2S Practice Questions One of the reactions that is used to inflate air bags involves sodium azide: NaN3 ---> 2Na + 3N2
Determine the mass of N2 produced from the decomposition of 100.0 grams of NaN3.
Mass to moles
Moles to moles
Moles to mass ANSWER: 65.6 g N2 Because a chemical reactions stops when one of the reactants is all used up Limiting Reactants RARLEY in nature does the reactants present the specific amount specified by the balanced chemical equation
USUALLY at least one of the reactants is in excess and is left over at the end of the chemical reaction - known as the EXCESS REACTANTS
The reactant that is limited, or there is less of that reactant, is known as the LIMITING REACTANT. They determine the amount of product formed Are the Car BODIES or Car TIRES the "limiting reactant"? The car BODIES there are not enough to use all the tires You are used to seeing and working with problems that are in the correct ratio - but that doesn't always happen in nature How to determine the Limiting Reactant Let's look at an example:
In the following chemical reaction - water is formed
On the product side, after the reaction, you have more oxygen left over (the red dots)
So oxygen is our excess reactant (there is some left over) and hydrogen is our limiting reactant (stopped the formation of water) Now see if you can find the excess and limiting reactants?? Limiting reactant = Hydrogen
Excess reactant = Nitrogen Now we will look at how to Calculate the amount of a Product when a REACTANT IS LIMITING! Example: S8 + 4Cl2 ---> 4S2Cl2
If 200.0g of sulfur reacts with 100.0 g of chlorine, what mass of disulfur dichloride is produced?
First determine what the limiting reactant is! - You need to find the number of moles for EACH reactant to do this. - Convert the masses of each reactant to moles (mass to moles)
Cl2 = 1.410 mol
S8 = 0.7797 mol
Then determine if the mole ratios are correct based off of the BALANCED chemical equation
In the chemical equation the ratio of Cl2 to S8 is 4:1
So compare your moles above to the correct ratio - divide the number of Chlorine moles by Sulfur
You find that it is only 1.808 NOT 4 - therefore chlorine is the limiting reactant!
Now calculate the amount of product formed:
Multiply the given number of moles of the limiting reactant x the mole ratio compared to the product (S2Cl2 ) over the limiting reactant (Cl2)
Then convert the moles to mass (grams) by multiplying by the molar mass of the product
Will be = 190.4 g of S2Cl2 What Else Can You Find? How to analyze the excess reactant Same Example: to find out how much of the excess reactant remains
You need to makes a mole to mass conversion to determine the mass of sulfur needed to completely react with 1.410 mol of Chlorine
Multiply the number of moles of Chlorine by the ratio of S8 (unknown) over Cl2 ( the known) = 0.3525 mol S8
The convert from moles to mass, multiply the number of moles by the molar mass of S8 = 90.42 g of S8 needed
Then knowing that 200.00 grams of sulfur (number from original question) is available and you only used 90.42 grams
200.00 - 90.42 = 109.6 grams of S8 in excess The reaction between solid white phosphorus (P4) and oxygen produces solid tetraphosphorus decaoxide (P4O10).
A. Determine the mass of P4O10 formed from 25.0 grams of P4 and 50.0 grams of O combined.
B. How much of the excess reactant remains after the reaction stops? Lets try another example: PART A:
1. First identify the limiting reactant
You determine the number of moles of EACH reactant by going from mass to moles
2. Then find the actual ratio of O2 to P4 based on the chemical equation and compare your mole data - then you find you limited reactant
Excess reactant will produce MORE than needed
3. Then calculate the moles of the product formed
Calculate from moles to mass (mole to mole first, then mole mass)
Answer = 57.3 g P4O10 PART B:
1. Use the moles of the limiting reactant to find the moles of the excess reactant
Mole of P4 times ratio of unknown to known
2. Convert the moles of the excess reactant to mass
Moles x molar mass
3. Subtract the total number of O2 available (given in the problem) by the amount used (what you just calculated)
Answer = 17.7 g O2 in excess Your Products? Chemical reactions do not always follow the perfect balanced equation and produce the perfect amount
Some reactions don't fully complete themselves for many reasons ( like part of the product being destroyed, or sticking to the side of the container, and ect.) Theoretical Yield:
This is the maximum amount of the product that can be formed from a reactant if the reaction goes perfect
Actual Yield:
Since chemical reactions rarely produce the theoretical yield this is the amount of the product ACTUALLY produced when a chemical reaction occurs The Percent Yield:
This is how chemist know if they are producing the desired amount of the product
It is ratio of the actual yield to the theoretical yield expressed as a percent Example #1:
Solid silver chromate (Ag2CrO4) forms when excess potassium chromate (K2CrO4) is added to a solution containing .500g of silver nitrate (AgNO3).
Determine the theoretical yield of Ag2CrO4.
Then Calculate the percent yield if the reaction yields .455 g of Ag2CrO4. To find the theoretical yield of Ag2CrO4 - do a mass to mol conversion (3 steps)
(Mass to moles) First go from the mass given of silver nitrate (.500g) to the moles = .00294 mol AgNO3
(Moles to Moles) Then go from moles of AgNO3 to moles of Ag2CrO4 = .00147 mol Ag2CrO4
(Moles to Mass) Last go from moles of Ag2CrO4 to the mass = .488 g of Ag2CrO4 Then find the percent yield:
Take the actual yield (.455g) and divide by the theoretical yield (.488g) and multiply by 100
Will equal 93.2% of Ag2CrO4 Now it's Your Turn:
1. Zinc reacts with iodine in a synthesis reaction:
Zn + I2 --> ZnI2
A. Determine the theoretical yield if 1.912 mol of zinc are used
B. Determine the percent yield if 515.6 g of the product is recovered
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