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SECTION 1 LEARNING TARGETS12.1.1 – I can explain how balanced
equations apply to both chemistry and everyday life.
12.1.2 – I can interpret balanced chemical equations in terms of moles, representative particles, mass, and gas volume at STP.
12.1.3 – I can identify the quantities that are always conserved in chemical reactions.
USING EVERYDAY EQUATIONS A balanced chemical equation provides
the same kind of quantitative information that a recipe does.
USING BALANCED CHEMICAL EQUATIONS Chemists use balanced chemical
equations as a basis to calculate how much reactant is needed or product is formed in a reaction.
EXAMPLE: Tiny tike has decided to make 288
tricycles each day. How many tricycle seats, wheels, and pedals are needed?
INTERPRETING CHEMICAL EQUATIONS A balanced chemical equation can be
interpreted in terms of different quantities, including; numbers of atoms, molecules, or moles, mass, and volume.
NUMBER OF ATOMS A balanced equation indicates that the
number and type of each atom that makes up each reactant also makes up each product.
NUMBER OF MOLECULES Here the coefficients tell you how many
molecules will react and form. Much like how many atoms.
MOLES A balanced equation also tells you the
number of moles of reactants and products.
The coefficients tell you this. You’ll use this most often.
MASS A balanced equation obeys the law of
conservation of mass. Using a mole relationship you can
relate number of moles to mass of either reactants or products.
VOLUME If you are at STP the equation tells you
about volumes of gases. You can use mole relationships for this
also.
EXAMPLE: Balance the following equation.
___C2H4(g) + ___O2(g) → ___CO2(g) + ___H2O(g)
Interpret the balanced equation in terms of relative numbers of moles, volumes of gas at STP, and masses of reactants and products.
SECTION 2 LEARNING TARGETS12.2.1 – I can construct mole ratios from
balanced chemical equations and apply these ratios in stoichiometric calculations.
12.2.2 – I can calculate stoichiometric quantities from balanced chemical equations using units of moles, mass, representative particles, and volumes of gases at STP.
WRITING AND USING MOLE RATIOS Mole ratio – a conversion factor derived
from the coefficients of a balanced chemical reaction interpreted in terms of moles.
In chemical calculations, mole ratios are used to convert between moles of reactant and moles of product, between moles of products, or between moles reactants.
MOLE-MOLE CALCULATIONS The easiest way to see these is to do
an example. W is the unknown, G is the given
quantity. a and b are the coefficients from the
balanced chemical equation.
4Al(s) + 3O2(g) → 2Al2O3(s) How many moles of aluminum are
needed to form 3.7 moles of aluminum oxide?
MASS-MASS CALCULATIONS In a reaction things are not measured
in moles (no scale does this). Instead things are measured in mass
then transferred to moles.
STEPS IN SOLVING A MASS-MASS PROBLEM
1. Change the mass of G to moles of G (mass G → mol G) by using the molar mass of G.
2. Change the moles of G to moles of W (mol G →mol W) by using the mole ratio from the balanced equation.
EXAMPLE: Acetylene gas (C2H2) is produced by
adding water to calcium carbide (CaC2).CaC2(s) + 2H2O(l) → C2H2(g) +
Ca(OH)2(aq) How many grams of acetylene are
produced by adding water to 5.00g of calcium carbide?
OTHER STOICHIOMETRIC CALCULATIONS In a typical stoichiometric problem, the
given quantity is first converted to moles.
Then the mole ratio from the balanced equation is used to calculate the number of moles of the wanted substance.
Finally, the moles are converted to any other unit of measure related to the unit mole, as the problem requires.
EXAMPLE: How many molecules of oxygen are
produced by the decomposition of 6.54g of potassium chlorate (KClO3)?
2KClO3(s) → 2KCl(s) + 3O2(g)
EXAMPLE: The equation for the combustion of
carbon monoxide is: 2CO(g) + O2 → 2CO2(g)
How many liters of oxygen are required to burn 3.86L of carbon monoxide?
SECTION 3 LEARNING TARGETS12.3.1 – I can identify the limiting
reagent in a reaction.13.3.2 – I can calculate theoretical yield,
actual yield, or percent yield given appropriate information.
LIMITING AND EXCESS REAGENTS When cooking you know you need the
right amounts of ingredients for the recipe to turn out.
In a chemical reaction, an insufficient quantity of any of the reactants will limit the amount of product that forms.
Limiting reagent – determines the amount of product that can be formed in the reaction. This is the one used up first in a reaction. The reaction can only “go” until this
reactant is completely used up. Excess reagent – reactant that is not
completely used up in a reaction.
How would the amount of products formed if you started with four molecules of N2 and three molecules H2?
EXAMPLE: The equation for the complete
combustion of ethene (C2H4) is: C2H4(g) + 3O2(g) → 2CO2(g) + 2H2O(l)
If 2.70mol of C2H4 is reacted with 6.30mol of O2, identify the limiting reagent.
The reactant that is present in the smaller amount by mass or volume is not necessarily the limiting reagent.
EXAMPLE: The heat from an acetylene torch is
produced by burning acetylene (C2H2) in oxygen:
2C2H2 + 5O2 → 4CO2 + 2H2O How many grams of water can be
produced by the reaction of 2.4 mol C2H2 with 7.4 mol O2?
PERCENT YIELD Your grades are usually expressed as a percent
Chemists use similar calculations when products are formed based on balanced equations.
In theory all reactions would produce at 100%. In reality they don’t.
%100totalright
Theoretical yield – the maximum amount of product that could be formed from given amounts of reactants.
Actual yield – the amount of product that actually forms when the reaction is carried out.
The percent yield is a measure of the efficiency of a reaction carried out in the laboratory.
Percent yield can be lowered by: Impure reactants. Loss of product in filtration or transferring. If reactants or products have not been
carefully measured.
EXAMPLE: When 84.8g of iron (III) oxide reacts
with an excess of carbon monoxide, iron is produced.
Fe2O3(s) + 3CO → 2Fe(s) + 3CO2(g) What is the theoretical yield of iron?