Download THE MOLE (pp. 159

THE MOLE (pp. 159 - 164)
Moles to Particles and Particles to Moles
1. Requires a return to dimensional analysis.
2. Use this as a conversion factor:
***** Determine the number of atoms in 2.50 mol Zn.
***** Determine the number of molecules in 11.5 mol H2O.
***** Zinc is used as a corrosion-resistant coating on iron and steel. It is also an essential trace element in
your diet. Calculate the number of moles that contain 4.50 x 10 24 atoms of Zn.
***** Determine the number of moles that contain 3.75 x 1024 molecules of CO2.
***** Determine the number of molecules in 3.25 mol of AgNO 3.
***** Determine the number of moles that contain 2.50 x 10 20 atoms Fe.
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Mass and the Mole
1. The mass in grams of one mole of any pure substance is called its ____________________.
***** Chromium is a transition element used as a coating on metals and in steel alloys to control corrosion;
calculate the mass in grams of 0.0450 mol Cr.
***** Determine the mass in grams of each of the following:
a. 3.57 mol Al
b. 42.6 mol Si
c. 3.45 mol Co
***** Calcium, the fifth most abundant element on earth, is always found combined with other elements
because of its high reactivity. Calculate the number of moles of Ca in 525g Ca.
***** Determine the number of moles in each of the following:
a. 25.5g Ag
b. 125g Zn
c. 1.00g Fe
Mass to Atoms and Atoms to Mass
1. By using molar mass and Avogadro’s number one can convert mass to number of atoms (particles) and
number of atoms (particles) to mass.
******* Gold is one of a group of metals known as the coinage metals (copper, silver, and gold).
Determine the number of atoms of gold in a pure gold nugget having a mass of 25.0g.
***** Determine the number of atoms in each of the following samples:
a. 55.2g Li
b. 0.230g Pb
c. 0.120g Ti
***** A party balloon contains 5.50 x 10 22 atoms of helium gas. Determine the mass of the gas.
***** Determine the mass in grams of each of the following:
a. 1.00 x 1024 atoms Mn
OVER
b. 1.50 x 1015 atoms N
c. 1.50 x 1015 atoms U
Molar Volume
1. Since the mole is the most convenient unit for counting numbers of atoms or molecules, the
______________________________ of a gas is defined as the volume that one mole occupies at 0.00 0C
and 1.00 atm pressure.
2. These conditions are known as standard temperature and pressure: ____________.
3. At STP, _________________________
***** Determine the volume that 0.881 mol of gas at STP will occupy.
***** Determine the number of moles of carbon monoxide inside a container having a volume of 22.8 L.
***** Determine the volume, in liters, occupied by 2.65g of krypton.
***** Determine the following:
a. The mass of xenon gas present in a 2.50L balloon at STP.
b. The volume, in liters, of H occupied by 0.0459 mol of H.
THE MOLE TRAFFIC CIRCLE:
MOLAR MASS OF A COMPOUND (pp. 165 - 170)
1. Molar mass: the sum of the molar masses of all the atoms represented by one mole of formula units.
2. To calculate the molar mass of an ionic compound, you must add the individual molar masses of all the
atoms in the compound.
***** Calculate the molar mass of KI.
***** Calculate the molar mass of Fe2O3.
OVER
***** Calculate the molar mass of Al2(SO4)3
***** Calculate the molar mass of the following compounds:
a. C2H5OH
b. KC2H3O2
c. (C3H5)2S
MASS RELATIONSHIPS or PERCENT COMPOSITION (pp. 171 - 174)
1. One way to determine the amount of an element in a compound is to calculate its percent composition.
2. To determine the percent composition, remember that the molar mass is 100% of the mass represented
by the formula.
3. The percentage of each element in the compound is simply the molar mass of the element divided by the
molar mass of the whole compound and multiplied by 100.
4. The equation looks like this:
***** Determine the percent composition of MgCl2
***** Determine the percent composition of NH4NO3
***** Determine the percent composition of each of the following:
a. Na2SO4
b. H3PO4
OVER
c. NaHCO3
WATER IN A HYDRATE
1. Some salts can bind water molecules within their crystal lattice structures.
2. When this happens, these compounds are known as ____________________.
3. Because they can absorb water, many salts are used as drying agents.
4. Chemical formulas for hydrates look like this:
***** Determine the percent calcium chloride, CaCl2  6H2O, is CaCl2
***** Determine the percent iron (III) chloride, FeCl 3  6H2O, is FeCl3
DETERMINING FORMULAS FROM PERCENT COMPOSITION DATA
(pp. 174 - 185)
Empirical Formula
1. When a new chemical substance is discovered, its discoverers often do not know its formula or its
composition.
2. The new compound is put through a series of reactions to determine the elements in it and the mass of
each element present. The chemical formula of the compound can then be calculated.
3. The key is to convert the percent composition data to _______________.
4. Then compare the mole amounts to find the simplest whole number ratio. The ratio tells you how many
atoms of each element are in the chemical formula.
***** Determine the chemical formula of a compound formed by mercury and chlorine given the
following data:
% Hg = 73.90%
%Cl = 26.10%
***** Potassium dichlorate, an orange solid, contains the three elements potassium, chromium and oxygen.
Analysis of the sample gives the following data:
%K = 26.60%
%O = 38.00%
%Cr = 35.40%
OVER
***** Determine the empirical formula of a compound containing carbon, hydrogen and nitrogen given the
following data:
%C = 38.67%
%H = 16.23%
%N = 45.10%
Molecular Formulas
1. Empirical formulas provide information as to what is in the compound by indicating the simplest whole
number ratio of atoms.
2. But, they do not always tell exactly how many atoms of each element are present in a molecule of the
compound. For that one needs the _______________________________.
3. Molecular formula – a formula that gives the type and actual number of atoms in a chemical compound.
4. The molecular formula can provide one with a better idea of the molecule’s properties based on exactly
how many of each atom there will be in a molecule.
5. In some cases the empirical formula is the same as the molecular formula (H 2O).
6. In most findings this is not the case:
Empirical Formula
Molecular Formula
Name
Determining Molecular Formulas
1. In the previous example, the molecular formulas were simply a multiple of the empirical formula.
2. A similar relationship exists between the molar masses of the empirical formula and that of the
compound.
***** Determine the molecular formula of a compound having an empirical formula of CH and a molar
mass of 78.11 g/mol.
***** A compound has the following composition:
76.54% C
12.13% H
11.33% O
If its molar mass is 282.45 g/mol, what is its molecular formula?
STOICHIOMETRY
Stoichiometry and Balanced Equations
1. A balanced chemical equation uses chemical formulas to show the identities and relative amounts of the
substances involved in a chemical reaction.
2. The reaction of powdered iron filings and oxygen is:
One can say four atoms of iron react with three molecules of oxygen to produce two molecules of iron (III)
oxide.
One can also say four moles of iron react with three moles of oxygen to produce two moles of iron (III)
oxide.
3. But one _______________ read this reaction in terms of mass amounts.
4. If we convert from moles to mass, we can then obtain information regarding the masses of the reactants
and products involved:
The total mass of the reactants and total mass of the products are equal.
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Mole Ratios
1. The coefficients in chemical equations indicate the relationships among moles of reactants and products.
2. The relationship between these coefficients can be used to write conversion factors called
_________________________.
3. A mole ratio is a ratio between the number of moles of any two substances in a balanced chemical
equation.
***** Determine all the mole ratios that can exist in the following chemical equation:
2Al (s) + 3Br2 (l) → 2AlBr3 (s)
4. Mole ratios are the key to calculations based upon a chemical equation.
STOICHIOMETRIC CALCULATIONS
Mole – to – Mole Conversions
1. One can determine the moles of a reactant needed or moles of a product formed if one knows the
balanced chemical equation and the number of moles of either one reactant or one product.
***** Determine the number of moles of hydrogen gas produced when 0.0400 moles of potassium are used
in the following reaction:
2K (s) + 2H2O (l) → 2KOH (aq) + H2 (g)
***** Sulfuric acid can be produced as follows:
2SO2 (g) + 2H2O (l) + O2 (g) → 2H2SO4 (aq)
If 12.5 moles of SO2 are used, determine the moles of H2SO4 produced.
If 12.5 moles of SO2 are used, determine the moles of O2 needed.
***** Methane and sulfur react to produce carbon disulfide and hydrogen sulfide:
2CH4 (g) + S8 (s) → 2CS2 (l) + 4H2S (g)
Determine the moles of CS2 produced when 1.50 moles of S8 are used.
Determine the moles of H2S produced when 1.50 moles of S8 are used.
Mole – to – Mass Conversion
1. Add one step to the railroad tracks to convert moles of desired material to mass.
***** Determine the mass (g) of sodium chloride produced when 1.25 moles of Cl 2 reacts with Na metal.
2Na (s) + Cl2 (g) → 2NaCl (s)
***** Determine the mass (g) of chlorine gas needed to react 1.25 moles of titanium (IV) oxide in the
following reaction:
TiO2 (s) + C (s) + 2Cl2 (g) → TiCl4 (s) + CO2 (g)
OVER
***** Calcium hydride reacts with water according to the following reaction:
CaH2 (s) + 2H2O (l) → Ca(OH)2 (aq) + 2H2 (g)
Determine the mass (g) of Ca(OH)2 produced when 2.50 moles of CaH 2 react.
Determine the mass (g) of H2O required to react with 2.50 moles of CaH2.
Mass – to – Mass Conversions
1. This is more typical of what a chemist needs to know.
2. These require conversions of mass to moles, then the use of the mole ratio and then a conversion of
moles to mass.
***** Ammonium nitrate decomposes when heated:
NH4NO3 (s) → N2O (g) + 2H2O (g)
Determine the mass (g) of H2O produced when 25.0g of NH4NO3 decomposes.
***** Methyl salicylate, C8H8O3, also known as oil of wintergreen, is most often made in a synthesis
reaction between methanol and salicylic acid, C7H6O3. Determine the mass (g) of salicylic acid required to
produce 325g of methyl salicylate.
C7H6O3 (s) + CH3OH (l) → C8H8O3 (s) + H2O (l)
LIMITING REACTANTS
Why Do Reactions Stop?
1. Reactions stop when one of the reactants is used up. The reactant that is used up first is called the
_________________________. It limits the extent of the reaction and, thereby, determines the amount of
product.
2. A portion of all of the other reactants remains after the reaction stops. These leftover reactants are
called _________________________.
3. Previous calculations were based on having the reactants present in the amounts needed.
4. To determine the amount of product formed when one reactant is limiting, one must first determine
which reactant is the limiting reactant.
Calculating the Product When a Reactant is Limited
***** Disulfur dichloride is used to vulcanize rubber, a process that makes rubber harder, stronger, and
less likely to become soft when hot or brittle when cold. It is produced by reacting sulfur and chlorine as
follows:
S8 (l) + 4Cl2 (g) → 4S2Cl2 (l)
Determine the mass (g) of S2Cl2 produced when 200.0g of S8 and 100.0g of Cl2 react.
OVER
***** Determine the mass (g) of tetraphosphorus decoxide, P 4O10, produced when 25.0g of phosphorus, P 4,
reacts with 50.0g of oxygen, O2.
P4 (s) + 5O2 (g) → P4O10 (s)
***** Silver nitrate and sodium phosphate react to produce silver phosphate and sodium nitrate:
3AgNO3 (aq) + Na3PO4 (aq) → Ag3PO4 (s) + 3NaNO3 (aq)
If 200.0g Na3PO4 and 200.0g AgNO3 are used in this reaction, determine:
a. the mass of solid silver phosphate produced
b. the limiting reactant
c. the reactant in excess
d. the mass of excess reactant that remains after the reaction is complete
PERCENT YIELD
1. Most chemical reactions do not succeed in producing the predicted amount of product.
2. The _________________________ is the maximum amount of product that can be produced from a
given amount of reactant.
3. The _________________________ is the amount of product actually produced when the chemical
reaction is carried out in an experiment.
4. The ratio of the actual yield to the theoretical yield expressed as a percent is the
_________________________ of a product. It is a measure of the reaction’s efficiency.
5. The actual yield comes from actual experimentation while the theoretical yield comes from
stoichiometric calculations.
***** When potassium chromate, K2CrO4, is added to a solution containing 0.500g silver nitrate, AgNO 3,
solid silver chromate, Ag2CrO4, is formed.
2AgNO3 (aq) + K2CrO4 (aq) → Ag2CrO4 (s) + 2 KNO3 (aq)
Determine the theoretical yield of silver chromate.
If 0.455g of silver chromate is actually produced, calculate the percent yield.
***** Zinc reacts with iodine in a synthesis reaction. Determine the theoretical yield of zinc iodide if
125.0g of zinc are used. Determine the percent yield if 515.6g of zinc iodide are actually produced.
Zn (s) + I2 (s) → ZnI2 (s)
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