Lab Report and Notebook Format
Analysis of Silver in an Alloy In this experiment an alloy
of silver will be analyzed to determine its silver content. The silver-copper alloy will
be dissolved in nitric acid, the silver will be precipitated as silver chloride, and the
silver chloride will be filtered, washed, dried and its mass determined. From the mass of
the silver chloride formed and the mass of the original sample, you will be able to
calculate the percent of silver in the alloy. Because the results are based on the mass of
a product, this procedure is classified as a gravimetric analysis. Once the silver and copper
are in solution, they can be separated from each other by precipitating the silver as
silver chloride. Silver chloride (AgCl) is very insoluble in water, while copper (II)
chloride (CuCl2) is soluble. Therefore, addition of chloride ions to the
solution will precipitate essentially all of the silver and none of the copper. The silver
chloride precipitate is filtered from the solution.
In this
experiment students study some metals and some nonmetals to find their relative
reactivity. A ranking according to reactivity is called an activity series. Reactions such as these are all
oxidation-reduction reactions. Substances,
which lose electrons during chemical reactions, are said to be oxidized. Those, which gain
electrons, are reduced. Oxidation and reduction reactions must occur simultaneously, and
there must be an equal number of electrons lost and gained. The most reactive metal is the
one most easily oxidized. When a substance readily
loses electrons (and is oxidized), it acts as a good reducing agent. When a substance has
a strong tendency to gain electrons (and be reduced), it also acts as a good oxidizing
agent. An activity series of the
nonmetallic halogens places the most reactive halogen at the top. Like the metals, the
more reactive nonmetals will displace ions of the less reactive halides from solution. In
an activity series of nonmetals, the most reactive halogen is the one most easily reduced.
Finding the Ratio of Moles of Reactants in a Chemical Reaction A balanced chemical
equation gives the mole ratios of reactants and products for chemical reactions. If the
formulas of all reactants and products are known, it is relatively easy to balance an
equation to find out what these mole ratios are. When the formulas of the products are not
known, experimental measurements must be made to determine the ratios. This experiment uses the
method of continuous variations to determine the mole ratio of two reactants. Several
steps are involved. First, solutions of the reactants are prepared in which the
concentrations are known. Second, the solutions are mixed a number of times using
different ratios of reactants. Third, some property of the reaction that depends on the
amount of product formed or on the amount of reactant that remains is measured. This
property may be the color intensity of a reactant or product, the mass of a precipitate
that forms, or the volume of a gas evolved. In this experiment the change of temperature
is the property to be measured. The reactions are all exothermic, so the heat produced
will be directly proportional to the amount of reaction that occurs. Since the experiment
is designed so that the volume of solution is a constant for all measurements, the
temperature change will also be proportional to the quantity of reactants consumed. In the method of continuous
variations, the total number of moles of reactants is kept constant for the series of
measurements. Each measurement is made with a different mole ratio of reactants. The
optimum ratio, which is the stoichiometric ratio in the equation, should consume the
greatest amount of reactants, form the greatest amount of product, and generate the most
heat and maximum temperature change.
Thermochemistry and Hess's Law The
molecular equations for the reactions are as follows: (1)
NaOH(aq) + HCl(aq) (2)
NH4CI(aq) + NaOH(aq) (3)
NH3 (aq) + HCl(aq) There is no single
instrument that can directly measure heat in the way a balance measures mass or a
thermometer measures temperature. However, it is possible to determine the heat change
when a chemical reaction occurs. The change in heat is calculated from the mass,
temperature change, and specific heat of the substance which gains or loses heat. The equation that is used
to calculate heat gain or loss is: q = (grams of substance) x (specific heat) x D T where q = the heat energy
gained or lost and DT is the change in temperature. Since DT = (final temperature minus initial temperature), an increase in
temperature will result in a positive value for both DT and q,
and a loss of heat will give a negative value. A positive value for q means a heat gain,
while a negative value means a heat loss. Acid-base neutralization is
an exothermic process. Combining solutions containing an acid and a base results in a rise
of solution temperature. The heat given off by the reaction (which will cause the solution
temperature to rise) can be calculated from the specific heat of the solution, the mass of
solution and the temperature change. This heat quantity can then be converted to the
enthalpy change for the reaction in terms of kJ/mole by using the concentrations of the
reactants. According to Hess, if a
reaction can be carried out in a series of steps, the sum of the enthalpies for each step
should equal the enthalpy change for the total reaction. Another way of stating
"Hess's Law" is: If two chemical equations can algebraically be combined to give
a third equation, the values of DH for the two equations can
be combined in the same manner to give DH for the third
equation. An examination of the acid-base equations above shows that if equation (2) is
subtracted from equation (1), equation (3) will result. Therefore, if the value of DH for equation (2) is subtracted from that of equation (1), the
enthalpy change for equation (3) should result. Molecular Mass of a Volatile Liquid It is often useful to know
the molecular mass of a substance. This is one of the properties that helps characterize
the substance. If the substance is a volatile liquid, one common way of determining its
molecular mass involves using the ideal gas law, PV = nRT. Since the liquid is volatile,
it can easily be converted to a gas. While it is in the gas phase, its volume, temperature
and pressure are measured. The ideal gas law will then allow the calculation of the number
of moles of the substance present: n
= PV/RT n
= grams gas/M
M = grams gas x RT/PV
Study of the Kinetics of a Reaction In
this experiment students study the kinetics of a chemical reaction. The reaction
is called a "clock" reaction because of the means of observing the
reaction rate. The reaction involves the oxidation of iodide ion by bromate ion
in the presence of an acid: 6
I-(aq) + BrO3-(aq) + 6 H+(aq)
The
reaction is somewhat slow at room temperature. Its rate depends on the
concentration of the reactants and on the temperature. The rate law is a
mathematical expression that relates the reaction rate to the concentrations of
reactants. If we express the rate of reaction as the rate of decrease in
concentration of bromate ion, the rate law has the form: where the square brackets
refer to the molar concentration of the indicated species. The rate is equal to
the change in concentration of the bromate ion, -Δ[BrO3-
], divided by the change in time for the reaction to occur, At. The term "k"
is the rate constant for the equation, and changes as temperature
changes. The exponents x, y, and z are called the "orders" of the
reaction with respect to the indicated substance, and show how the concentration
of each substance affects the rate of reaction. One
purpose of the experiment is to determine the total rate law for the process. To
do this we must measure the rate, evaluate the rate constant, k,
and determine the order of the reaction for each reactant, the values
of x, y, and z. A second goal is to determine the activation energy for the
reaction. Lastly we see the effect a catalyst has on the reaction rate. To
find the rate of the reaction we need some way of measuring the rate at which
one of the reactants is used up, or the rate at which one of the products is
formed. The method that we use is based on the rate at which iodine forms. If
thiosulfate ions are added to the solution they react with iodine as it forms in
this way: I2(aq)
+ 2 S2O32-(aq)
Reaction
(1) is somewhat slow. Reaction (2) proceeds extremely rapidly, so that as
quickly as iodine is produced in reaction (1), it is consumed in reaction (2).
Reaction (2) continues until all of the thiosulfate is used up. After that,
iodine begins to increase in concentration in solution. If some starch is
present, iodine will react with the starch to form a deep blue‑colored
complex that is readily apparent. Carrying
out reaction (1) in the presence of thiosulfate ion and starch produces a
chemical "clock." When the thiosulfate is consumed, the solution turns
blue. In
all of our reactions we use the same quantity of thiosulfate ion. The blue color
appears when all the thiosulfate is used up. An examination of equations (1) and
(2) shows that 6 moles of S2O32- are needed to
react with the 12 formed from 1 mole of BrO3-.
Knowing the amount of thiosulfate used allows the calculation of the amount of
12 that is formed, and also the amount of BrO3- that
has reacted at the time of the color change. The reaction rate is expressed as
the decrease in concentration of BrO3- ion divided
by the time it takes for the blue color to appear. The
experiment is designed so that the amounts of the reactants that are consumed
are small in comparison with the total quantities present. This means that the
concentration of reactants is almost unchanged during the reaction, and
therefore the reaction rate is almost a constant during this time.
Analysis of a Commercial Bleach In this experiment the amount of hypochlorite ion present in a solution of bleach is determined by an oxidation-reduction titration, the iodine-thiosulfate titration procedure. In acid solution, hypochlorite ions oxidize iodide ions to form iodine, I2. The iodine that forms is then titrated with a standard solution of sodium thiosulfate. The analysis takes place in a series of steps: (1) Acidified iodide ion is added to hypochlorite ion solution, and the iodide is oxidized to iodine.
2 H+(aq) + ClO-(aq)
+ 2 I-(aq) (2) Iodine is only slightly soluble in water. It dissolves very well in an aqueous solution of iodide ion, in which it forms a complex ion called the triiodide ion. The triiodide ion is yellow in dilute solution, and dark red-brown when concentrated.
I2(aq) + I-(aq)
(3) The triiodide is titrated with a standard solution of thiosulfate ions, which reduces the iodine back to iodide ions:
I3-(aq)
+ 2 S2O32-(aq)
During this last reaction the red-brown color of the triiodide ion fades to yellow and then to the clear color of the iodide ion. It is possible to use the disappearance of the color of the I3- ion as the method of determining the end point, but this is not a very sensitive procedure. Addition of starch to a solution that contains iodine or triiodide ion forms a reversible blue complex. The disappearance of this blue colored complex is a much more sensitive method of determining the end point. The quantity of thiosulfate used in step (3) is directly related to the amount of hypochlorite initially present. Determination
of the Equilibrium Constant This experiment is designed to
measure the equilibrium constant for the formation of the thiocyanatoiron(III)
complex ion: Fe3+(aq)
+ HSCN(aq) The complex ion is a wine-red
color in solution. A series of solutions of FeSCN2+ with known
concentrations are prepared. The absorbance of each is measured using a
spectrophotometer, and a Beer's law calibration graph is constructed relating
absorbance to concentration. Next a series of solutions with varying amounts of
Fe3+ and HSCN are prepared and the absorbance of each is measured.
Concentrations of the complex ion are determined from the calibration graph.
Stoichiometry is used to determine concentrations of other species, and the
value of the equilibrium constant for the formation of the complex ion is
computed for each solution. The values for the equilibrium constant are then
averaged. Laboratory skills involve measuring different volumes of solutions with precision and diluting to volume in a volumetric flask, use of a spectrophotometer to measure absorbance of solutions, construction and use of a calibration curve for the spectrophotometric measurements of the thiocyanatoiron(III) complex ion, stoichiometric calculations to determine concentrations of species in solution, and calculation of the equilibrium constant from the data gathered. Determination of
the Solubility Product of an Ionic
Compound The
solubility product constant, Ksp, is a particular type of equilibrium
constant. The equilibrium is formed
when an ionic solid dissolves in water to form a saturated solution. The
equilibrium exists between the aqueous ions and the undissolved solid. A
saturated solution contains the maximum concentration of ions of the substance
that can dissolve at the solution's temperature. The equilibrium equation
showing the ionic solid lead chloride dissolving in water is: PbCl2(s)
The
solubility product expression is: Ksp
= [Pb2+] [Cl-]2 where
square brackets refer to molar concentrations of the ions. A knowledge of the Ksp
of a salt is useful, since it allows us to determine the concentration of ions
of the compound in a saturated solution. This allows us to control a solution so
that precipitation of a compound will not occur, or to find the concentration
needed to cause a precipitate to form. The
solubility product which will be determined by this experiment is that of the
strong base, calcium hydroxide, Ca(OH)2 This
experiment uses a microscale technique to determine the solubility product of
calcium hydroxide, Ca(OH)2. A solution of calcium nitrate is diluted and the
diluted solutions are combined in a well plate with sodium hydroxide. Calcium
hydroxide precipitates in the wells that contain concentrations of calcium and
hydroxide ions that exceed the solubility product. The first well where no
precipitate is present is assumed to be a saturated solution, and the ion
concentrations are used to calculate the solubility product. The process is
repeated using a dilution of the sodium hydroxide solution. Students must use good
technique to measure small solution volumes and make dilutions using Beral
capillary pipets. Calculations involve determination of the concentrations of
the diluted solutions and calculation of the solubility product. Acid-base titrations can be used to measure the concentration of an acid or base in solution, to calculate the formula (molar) mass of an unknown acid or base, and to determine the equilibrium constant of a weak acid or weak base. Titration is a method of volumetric analysis—the use of volume measurements to analyze an unknown. The purpose of this experiment is to standardize a sodium hydroxide solution and use the standard solution to titrate an unknown solid acid. The equivalent mass of the solid acid is determined from the volume of sodium hydroxide added at the equivalence point. The equilibrium constant, Ka, of the solid acid is calculated from the titration curve obtained by plotting the pH of the solution versus the volume of sodium hydroxide added. The concentration of the NaOH solution must be accurately known. To "standardize" the NaOH, that is, to find its exact molarity, the NaOH is titrated against a solid acid, potassium hydrogen phthalate (abbreviated KHP). In this experiment the equivalent mass of an unknown acid is determined by titration. An acid may contain one or more ionizable hydrogen atoms in the molecule. The equivalent mass of an acid is the mass that provides one mole of ionizable hydrogen ions. It can be calculated from the molar mass divided by the number of ionizable hydrogen atoms in a molecule. The acid, a solid crystalline substance, is weighed out and titrated with a standard solution of sodium hydroxide. From the moles of base used and the mass of the acid, the equivalent mass of the acid is calculated. The equivalent mass is determined by titrating an acid with a standard solution of NaOH. Since one mole of NaOH reacts with one mole of hydrogen ion, at the equivalence point the following relation holds: Vb x Mb = moles base = moles H+ EM = grams acid/moles H+
The acid is then titrated a second time with the standard solution of sodium hydroxide and the course of the titration is followed using a pH meter. A titration curve is constructed by graphing pH on the vertical (y) axis versus the volume of NaOH on the horizontal (x) axis. The value of the equilibrium constant (Ka.) for the dissociation of the weak acid is determined by analyzing the titration curve. A graph of pH versus volume of NaOH added for the titration of a weak acid with NaOH is obtained by carefully following the titration with a pH meter. There is a significant change in pH in the vicinity of the equivalence point. The value of the equilibrium constant for the dissociation of the acid is determined from a titration curve by considering the pH when the acid is "half-neutralized." If the dissociation
of the acid is represented as: HA + H20
↔ H30+ + A-
When the acid is half neutralized, [HA] = [A-], these terms cancel in the above equation, and Ka = [H30+]. Therefore, when the acid is half-neutralized, the pH = pKa. The point where pH is equal to pKa can be determined from the graph.
Oxidation-reduction reactions are studied by constructing various electrochemical cells and measuring the voltage generated. From these measurements, a reduction series is generated, the concentration of copper ions in solution determined, and the Ksp of silver chloride calculated. In Part 1 of this laboratory a "standard" table of electrode potentials in order of ease of reduction is constructed. The series of microscale half-cells is constructed by placing a piece of metal into a 1.0 M solution of its ions for each metal in the series. The metals chosen are copper, iron, lead, magnesium, silver, and zinc. The half-cells are connected by a salt bridge constructed of a strip of filter paper soaked in a solution of potassium nitrate. The zinc half-cell is chosen as the reference standard and is assigned a value of 0.00 volts. All potentials are measured with respect to the zinc electrode. These voltage values should correlate with those found in published tables, but they will differ by the value of E° for the standard zinc electrode. In Part 2, the Nernst equation is applied to the voltage measurement of a cell with nonstandard copper ion concentration. A solution of 0.0010 M Cu2+ is prepared, and the voltage of the cell: Zn(s) | Zn2+(1.0 M) || Cu-(0.0010 M) | Cu(s) is measured. The measured voltage is compared to that calculated from the Nernst equation. The table of standard potentials assumes that all ion concentrations are 1.0 M, gas pressures are 1 atm, and temperature is 25 °C. Calculations of potentials under nonstandard conditions can be made using the Nemst equation: RT E = E° - —— In Q nF where E = the measured cell potential, E° = the standard cell potential, R is the gas constant (8.314 J/ mol-K), T is the temperature (K), n = the number of moles of electrons transferred as shown by the oxidation-reduction equation, and F is the Faraday constant (9.65 x 104 C/mol). Q is the reaction quotient: the actual concentrations of products and reactants substituted into the equilibrium constant expression. Using base 10 or common logarithms the expression can be written: 2.303 R T E = E° - ————— log Q nF Substituting for the constants 2.303, R, and F, and using a temperature of 25 °C (298K) the expression can be simplified to: 0.0592V E = E° - ———— log Q n A measurement of the cell potential, E, under nonstandard conditions, can be used to calculate the value of Q, which can then be used to determine unknown concentrations of ions actually present in a solution. In the final section of the lab, the solubility product constant of silver chloride, AgCl, is determined from the Nernst equation and the voltage of a cell in which the zinc half-cell is connected to a solution containing a trace of Ag+ ions in a 1.0 M solution of sodium chloride, NaCl.
Last Updated on
04/16/2008 |
NaCl(aq) + H2O(l)
FeSCN2+(aq) + H+(aq)