4
PASCO Experiment Sensors and Chemical Reactions
I.
Learning Objectives
To learn the use of common sensors (transducers) that will be used in the
laboratory.
To investigate some types of reactions which occur in aqueous solutions.
II.
Background Information
Much of the research carried out in the modern laboratory involves electronic
instrumentation. This is as true in chemistry as it is in high-energy physics. Many of these
instruments are expensive, sophisticated, extremely accurate, and require a Ph.D. and a
team of dedicated technicians to operate. The advent of microelectronics, lasers, and
computers has revolutionized the applications possible for analytical instrumentation in the
laboratory. Unfortunately, many of these developments have made the link between the
scientific principles of the experiment and the output from the instrument less clear for the
undergraduate science student. For the well-experienced scientist or engineer, this is
generally not a problem. For a student who is just getting started in science, however,
instruments can become intimidating "black boxes," spewing forth a mass of meaningless
output signals. In an attempt to circumvent some of these problems, the laboratories in this
general chemistry course adopt a different approach in their use of instrumentation. The
general idea is to work with small, simple versions of the more expensive, commercially
available instrumentation. It is important at this point to note that simple and small does
not mean trivial. With careful preparation, results are possible that are quite precise and
accurate. Many students in the general chemistry curriculum will encounter the more
sophisticated versions of some of the instruments in these experiments later in their
academic/scientific careers.
The basic principle of most scientific instruments is embodied in a sensor capable of
probing/monitoring a physical condition of the system under investigation connected to a
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measuring device of some type. A simple, classical example is the instrument that has
been used for many centuries to measure the thermal energy of a system i.e., the
thermometer (measures hotness or coldness). The sensor in this instrument is a liquid
(typically mercury or alcohol) that expands or contracts as it is heated or cooled by the
environment. The key to the function of this sensor is the standardization of the liquid
expansion to a useable measuring scale. Unfortunately, throughout the more than three
hundred years of development, many different scales have been attached to many different
designs of the thermometer. Even today, there are at least four or five scales commonly
used throughout the world. It is proving an extraordinary difficult task to choose one as a
standardized system (hence, the temperature conversion questions common on freshman
chemistry tests!). It is worth emphasizing that if one particular scale were chosen, the
choice would be arbitrary. Of course, every manufactured thermometer must somehow be
calibrated against a known temperature source. This series of experiments will investigate
several different types of chemical reactions using various sensors, the Science Workshop
500 (SW 500) interface, and DataStudio.
All matter has electronic characteristics. Atoms are composed of negatively charged
electrons and positively charged protons.
Electrons are responsible for the chemical
properties of solids that form the basis for modern electronics. Solids can be divided into
three types, conductors, semiconductors, and insulators, depending on how well the
material conducts electricity. Metals are excellent conductors of electricity because there
are many free electrons that can be easily pushed through the macroscopic structure. Most
nonmetallic elements and compounds contain strongly bound electrons that are fixed in
bonding positions close to the original atoms and are not free to move. These substances
do not conduct electricity and are called insulators. Semiconductor substances have an
intermediate electrical conductivity that can be controlled to any desired level during the
manufacture of the substance.
Combinations of all of these types of materials in
extraordinarily thin microlayers and regions can be arranged in carefully designed circuits
to control the flow of charge in microelectronic devices.
The flow of charges through a substance is known as an electric current. Current
flowing in only one direction is referred to direct current (DC). When the direction of
current flow is oscillating back and forth (in both directions), it is termed alternating
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current (AC). The fundamental law that quantitatively describes the flow of direct current
is called Ohm's law,
V = IR
where V is the voltage, I is the current, and R is the resistance. Voltage (V), sometimes
called potential, is the electrical pressure or force that pushes charge through the
conductor. A good analogy to describe current is to think of it as water flowing through a
pipe. The voltage is analogous to the water pressure that is pushing water through the
pipe. The current (I) is the quantity of electrons passing a point in one second and is
measure in ampere units. One ampere is 6.25 x 1018 electrons passing a point in one
second. All conductors tend to resist the flow of charge to some extent. A conductor has a
resistance (R) of 1 ohm ( ) if a voltage of 1 volt (V) will force a current of 1 ampere
(Amp) through it. Another useful term is power (P), which is the work performed by all
electrical current measured in watt (W) units. One Watt is equivalent to 1 Joule of energy
per second. The power of a direct current is voltage multiplied by current:
P = VI
An electric circuit is any arrangement of electrical materials that will allow a current
to flow. The most basic circuit for direct current flow has a power source (e.g., a battery),
a resistor, and conducting connectors (usually wires). This setup is represented in the
circuit diagram of Figure 4.1. Resistors play a very important role in the design of circuits
for scientific instruments because they limit the flow of current, and they exhibit a
temperature dependence.
This latter property can be extremely useful in designing
sensitive temperature sensors.
Figure 4.1 A Basic DC Circuit
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A transducer is an electronic device for converting energy from one form to another
for the purpose of measurement of a physical quantity.
Common examples include
microphones, loudspeakers, thermometers, position and pressure sensors. Although not
generally thought of as transducers, photocells, LEDs (light-emitting diodes), and even
common light bulbs are transducers.
Efficiency is an important consideration in any transducer. Transducer efficiency is
defined as the ratio of the power output in the desired form to the total power input.
Mathematically, if P represents the total power input and Q represents the power output in
the desired form, then the efficiency E, as a ratio between 0 and 1, is given by:
E = Q/P
If %E represents the efficiency as a percentage, then:
%E = (Q/P) x 100
It is important to point out that no transducer is 100% efficient; some power is always lost
in the conversion process. Usually this loss is manifested in the form of heat. Among the
worst transducers, in terms of efficiency, are incandescent lamps. A 100-watt bulb radiates
only a few watts in the form of visible light. Most of the power is dissipated as heat; a
small amount is radiated in the UV (ultraviolet) spectrum.
PASCO Sensors
Temperature Sensor
The temperature sensor is a transducer that consists of a precision integratedcircuit (IC) chip, whose output voltage is linearly proportional to the Celsius temperature.
This device exhibits a positive temperature coefficient, meaning that the resistance
increases as the temperature increase in the range from 5oC to +105oC. This IC chip
produces a 10-mV signal for every degree Celsius (oC). The voltage reading simply
multiplied by 100 gives the temperature reading in oC.
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Pressure Sensor
All pressure sensors consist of two ports that are separated by a diaphragm. In a
differential gas pressure sensor, both ports are open to the surroundings allowing for the
application of pressure to either side of the diaphram.
The difference between the
pressures applied to these ports are measured. For example, if both ports are left open to
the atmosphere, the effect of atmospheric pressure will equal each other and the difference
between the ports will be zero. If only one port is used and the other port is left open to the
atmosphere, the sensor will measure the pressures relative to atmospheric pressure.
Experiments that would traditionally use a manometer can utilize a differential gas
pressure sensor.
In the general chemistry laboratory, an absolute gas pressure sensor will be used. In
this type of sensor, the pressure is measured against a built-in internal vacuum reference.
If the port is left open to the atmosphere, the sensor will display the value for atmospheric
pressure. The sensor produces an output voltage that is directly proportional to absolute
pressure. The output voltage from the sensor is +1.00 V when the pressure is 100 kPa.
Therefore, the output voltage should be +7.00 V at the top of the range (700 kPa). The
resolution of the sensor is 0.5 kPa.
pH Sensor
The pH sensor consists of a pH electrode and an amplifier. The most widely employed
ion-selective electrode is the glass electrode for measuring pH. A pH electrode responds
selectively to H+, building up a potential difference of 0.05916 V for every factor-of-10
difference in [H+] across the electrode membrane. A factor-of-10 difference in [H+] is one
pH unit, so a change of 4.00 pH units leads to a change in electrode potential of 4.00 x
0.05916 = 0.237 V.
A typical combination electrode, incorporating both glass and reference electrodes in
one body, is shown in Figure 4.2. The line diagram of this cell is
Glass membrane
Ag(s) AgCl(s)
Cl-(aq)
Outer reference electrode
H+(aq, outside)
H+(aq, inside),
H+ outside
glass electrode
(analyte solution)
H+ inside
glass electrode
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Cl-(aq)
AgCl(s) Ag(s)
Inner reference electrode
The pH-sensitive part of the electrode is the thin glass membrane (bulb) at the bottom end
of the electrode in Figure 4.2.
The glass membrane at the bottom of the pH electrode consists of an irregular network
of SiO4 tetrahedra through which Na+ ions move sluggishly but H+ ions are incapable of
penetrating. The glass surface contains exposed O- groups that can bind H+ from the
solutions on either side of the membrane (Figure 4.2). H+ equilibrates with the glass
surface, thereby giving the side of the membrane exposed to the higher concentration of H+
the more positive charge.
Na+ ions that are already in the glass migrate across the
membrane from the positive side to the negative side, so the potential changes by 0.05916
V for a unit change in pH. The amplifier translates the differences in electrical potential
into voltages which are sent to the interface.
Figure 4.2 a) Glass combination electrode with the silver-silver chloride reference
electrode. The two Ag electrodes measure the voltage across the glass membrane. b) Ionexchange equilibria on the inner and outer surfaces of the glass membrane. The pH of the
internal solution is fixed. As the pH of the external solution (the sample) changes, the
electric potential difference across the glass membrane changes.
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Conductivity Sensor
Like the pH sensor, a conductivity sensor has two components, an electrode and an
amplifier.
Simple conductivity electrodes are constructed of an insulating material
embedded with platinum, graphite, stainless steel or other metallic pieces. These metal
contacts serve as sensing elements and are placed at a fixed distance apart to make contact
with a solution whose conductivity is to be determined. The length between the sensing
elements, as well as the surface area of the metallic piece determine the electrode cell
constant, defined as length/area. The cell constant is a critical parameter affecting the
conductance value produced by the cell and handled by the electronic circuitry.
A cell constant of 1.0 will produce a conductance reading approximately equal to the
solution conductivity. For solutions of low conductivity, the sensing electrodes can be
placed closer together, reducing the length between them and producing cell constants of
0.1 or 0.01. This will raise the conductance reading by a factor of 10 to 100 to offset the
low solution conductivity and give a better signal to the conductivity meter. On the other
hand, the sensing electrodes can be placed farther apart to create cell constant of 10 or 100
for use in highly conductive solutions. The conductivity electrode provided with the
PASCO interface has a cell constant of 1.0 and is designed to achieve optimum
performance over a range of 0 to 20,000 S (Figure 3.4). This performance is achieved by
using a cylindrical cell geometry and platinized platinum conductors embedded on a glass
rod.
The conductivity amplifier has two distinct functions. First, it provides a voltage used
to drive the conductivity electrode. When a potential is applied to the conductivity cell, the
ions in solution are influenced by the charge on the cell's electrodes and begin to migrate
toward the electrodes (Figure 4.3). Then, the second function of the amplifier is to sense
the electrical current the electrode passes when placed in the solution to be tested.
If the voltage and current are known, then the resistance of the cell can be determined
using Ohm's law. If the resistance of the cell is known, taking the inverse of the resistance
and multiplying by the conductivity cell constant will determine the conductivity.
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Figure 4.3 a) Schematic view of the conductivity electrode. b) View of the conductivity
cell in operation.
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Colorimeter
The colorimeter measures the amount of light that is transmitted through a liquid.
Monochromatic light from a LED light source passes through a cuvette containing a
solution sample, as shown in Figure 4.4. Some of the incoming light is absorbed by the
solution. As a result, light of a lower intensity strikes a photodiode. The intensity of the
light passing through the liquid can often be used to determine properties of the liquid such
as the concentration of chemical in the liquid. Substances absorb different amounts of
certain colors of light and transmit other colors. Some substances absorb red but not blue
or green but not orange. The colorimeter can shine the following colors of light through a
liquid: orange (630 nm), green (565 nm), blue (460 nm) and red (697 nm).
to interface
Figure 4.4 Schematic diagram of colorimeter. Monochromatic light, produced by the
LED, passes through the solution contained in the cuvette. Light that is not absorbed by
the solution passes through the cuvette and is measured by the photodiode
The colorimeter is designed to use polystyrene cuvettes that hold approximately 4
mL. The cuvette slot of the colorimeter is designed to give a snug fit to the cuvette and
ensure that it is always in precisely the same position between the LED light source and
photodiode. Two opposite sides of the cuvette are ribbed and are not intended to transmit
the light from the LED. The two smooth surfaces are intended to transmit light. It is
important to position the cuvette correctly in the colorimeter so that the light traveling
from left to right from the LED will pass through the cuvette to the photodiode. This
should be done as shown in Figure 4.5, with the ribbed edges facing away from and toward
you, and the smooth edges facing left and right.
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Figure 4.5 Correct placement of cuvette in colorimeter. Always position the cuvette so the
light beam will pass through the clear sides. The colorimeter door will not close properly
if the cuvette is in the wrong position.
The amount of light that penetrates a solution is known as transmittance.
Transmittance can be expressed as the ratio of the intensity of the transmitted light
measure by the photodiode, It, and the initial intensity of the light beam, Io, as expressed by
the formula:
T = It / Io
The colorimeter produces an output voltage which varies in a linear way with
transmittance, allowing a computer to monitor transmittance data for a solution. The
transmittance of the sample varies logarithmically (base ten) with the product of three
factors: a, the molar absorptivity (extinction coefficient), b, the cell or cuvette width, and
c, the molar concentration.
log(l/T) = abc
In addition, many experiments designed to use a colorimeter require a related
measurement, absorbance. At first glance, the relationship between transmittance and
absorbance would appear to be a simple inverse relationship. As the amount of light
transmitted by a solution increases, the amount of light absorbed might be expected to
decrease proportionally. The true relationship between these two variables is inverse and
logarithmic (base 10). It can be expressed as:
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A = log(1/T)
Combining the two previous equations, the following expression is obtained:
A = abc
This formula more commonly referred to as Beer's Law states that the light
absorbed by a solution depends on the absorbing ability of the solute, the distance traveled
by the light through the solution, and the concentration of the solution. For a given
solution contained in a cuvette with a constant cell width, one can assume a and b to be
constant. Therefore, absorbance is directly proportional the the concentration.
Solutions and Chemical Reactions
A homogeneous mixture of two or more substances is called a solution. Usually, the
substance present in a smaller amount is called the solute, and the substance present in the
larger amount is called the solvent. Solutions may be gaseous (e.g., air), liquid (e.g.,
blood), and solid (e.g., steel). Liquid solutions are particularly important in chemical
systems; the universal solvent for these solutions is water. The solution is the medium that
brings reactants close together long enough to allow new associations and bonds to form.
Products appear as bond formation moves toward completion. The solvent almost always
plays an active role in all of these processes, especially when water is the solvent.
The solvent must first dissolve the substances that are eventually going to react. This
deceptively simple process, called dissolution, also involves the breaking of bonds and the
formation of new associations. The solubility of a substance reflects how easily the
solvent can make these changes occur. The solubility of a substance is, in fact, defined as
the concentration of solute in a saturated solution at a specified temperature. The reverse
of dissolution is the coming out" of solution in which the product(s) of a reaction exceed
the solubility and form a new phase that could be a gas, liquid, or solid. Again, the solvent
often plays an active part in this process.
At the molecular level these dynamic aqueous solution processes can be pictured as
shown in Figure 4.6.
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Figure 4.6. Dynamic Aqueous Solution Processes
where
are solute and
are water molecules.
The extent to which each of these processes happens can be described in terms of three
major interactions:
solute-solute interactions
solute-water interactions
water-water interactions
If the solute-solute interaction is dominant, then solute will come out of solution and
two phases will form. Which phase goes up (perhaps a gas) and which phase goes down
(perhaps a solid) depends on external forces, such as gravity.
If the water-water
interaction is dominant, then the solute will be "squeezed out" of the solution and form
another phase. In between these extremes, water-solute interactions can assure a stable,
homogeneous solution.
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The solubility of solids in water is governed by the three major interactions discussed
above. A simple example is the dissolution of sodium nitrate (NaN03) in water. Solid
sodium nitrate has a crystal structure that consists of sodium ions (Na+) ionically bonded to
nitrate ions (N03-). Water is a good solvent for salts such as sodium nitrate because water
molecules are able to move in between the cation and anion and screen the charges from
each other. When this happens the ionic bond weakens, and the water molecules can then
orient around the dissociated (separated) ions with the negative end of the water dipole
pointed towards the cation and the positive end pointed towards the anion. The water
molecules are bonded quite strongly to the ions by the Coulombic attraction of unlike
charges. Bound water molecules surround all cations and anions dissolved in aqueous
solution, and they are said to be hydrated. Figure 3.7 illustrates the process of dissolution
and hydration.
Figure 4.7. Dissolution and Hydration of an lonic Crystal
NOTE: Most chemistry texts, including the one used in Chemistry 21 assume that the
reader understands that cations and anions in aqueous solution exist as hydrated ions;
therefore the formulas for ions are usually written without the bonded solvent molecules.
The extent to which the dissolution process occurs i.e., how soluble the salt is?
depends on the magnitude of the attraction between ions and water molecules (hydration
energy) compared with the attraction between ions in the solid salt (lattice energy). The
battle of forces within the crystal versus those between ions and water molecules is
discussed in a later section on precipitation reactions.
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The strength of the attraction between ions and water molecule dipoles depends on the
charge and the radius of the ion. Small ions with high charge and electronegativity have
the greatest attraction for water molecules; conversely, large ions with a single charge have
the smallest attraction. If the attraction of the metal ion for the negative end of the water
molecule dipole is strong enough, the water molecule may be ripped apart, releasing a
hydrogen ion (H+) that is then hydrated. A good example is the aluminum ion (radius 67
pm and a +3 which charge), can react with water to produce the hydroxy cation
[Al(H2O)5OH]2+ and the hydronium ion H3O+:
[Al (H2O)6]3+ + H2O [Al (H2O)5OH]2+ + H3O+
Compounds that react with water to produce hydronium ions are called acids. The
aluminum ion is therefore called an acidic cation, whereas the sodium ion discussed
earlier is called a nonacidic cation.
Anions in aqueous solution also interact with water, except in this case, the positive
end of the water dipole is attracted to the anion. If the anion has a small size and high
charge, then water molecules are pulled apart and the hydrogen atom bonds to the anion,
releasing a hydroxide ion (OH-). A simple example is the reaction of a carbonate ion
(CO32-) with water,
CO32- + H2O HCO3- + OHwhich produces a bicarbonate ion (HC03-) and a basic (the opposite of acidic) solution.
The carbonate ion is called a basic anion whereas the nitrate ion (NO3-) discussed earlier is
called a nonbasic anion.
The chemistry of ionic solutes in aqueous solution becomes considerably more
interesting and complicated when two solutions, each containing a soluble ionic
compound, are mixed. Generally, in this situation the resulting solution, at the instant of
mixing, contains two different hydrated cations and two different hydrated anions. If
strong solute-solute interaction occurs between any two of the ions, then a chemical
reaction results, and a new product is formed. Often the two remaining ions do not interact
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and remain in solution in the same hydrated state as before the reaction. These unreacted
ions are often called spectator ions. Chemical reactions in aqueous solution may be
divided into many different types: precipitation, acid-base, complexation, gas-forming,
and endothermic/exothermic reactions. A brief discussion of the main characteristics of
each of these reaction types follows.
PRECIPITATION REACTIONS
A precipitation reaction is defined as a reaction that produces a new compound that is
not soluble in an aqueous solution.
Precipitation can be regarded as the reverse of
dissolution. The factors mentioned earlier i.e., hydration energy and lattice energy
play a major role in determining whether an ionic compound will precipitate from solution.
Ionic compounds (salts) in which the cation and anion have approximately the same size
and the same charge tend to form especially stable crystal structures and precipitate from
aqueous solution. Acidic cations and nonbasic anions give rise to soluble salts because
these ions are quite different in size and have much smaller lattice energy than hydration
energy. Although these generalizations are useful, it is often difficult to make exact
predictions, and it is worthwhile to remember a few simple aqueous solubility rules:
1.
2.
3.
4.
5.
Most compounds containing Na+, K+, or NH4+ ions are soluble.
Most compounds containing NO3-, C2H3O2-, ClO4-, Cl-, Br-, or I- ions are soluble.
Most compounds containing Ag+, Pb2+, or Hg22+ ions are insoluble.
Most compounds containing CO32- or OH- ions are insoluble.
Most compounds containing SO42- ions are soluble; however, BaSO4 and CaSO4
are insoluble.
ACID-BASE REACTIONS
In aqueous solution an acid reacts with a base to give a salt and water. An acid/base
reaction, sometimes called neutralization, is characterized by the formation of covalent,
neutral water molecules from acidic hydronium ions (H30+) and basic hydroxide ions
(OH-).
COMPLEXATION REACTIONS
Complexation reactions are closely related to acid/base reactions.
Earlier, it was
pointed out that some metal cations that have small size and high charge can act as acids
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in fact they are called acidic cations. These acidic cations can react with electron-rich
species called ligands to form complexes. A complex may be defined as a chemical
compound in which there is one or more coordinate covalent bonds. A coordinate
covalent bond is a covalent bond in which the shared pair of electrons is provided by the
ligand. A complexation reaction is therefore defined as a reaction in which one or more
coordinate covalent bonds are produced during the formation of product.
GAS-FORMING REACTIONS
Another type of common reaction is one that yields the production of an insoluble gas
as a product. The reactions will produce one of the following gases: hydrogen gas (H2),
carbon dioxide (CO2), sulfur dioxide (SO2), hydrogen sulfide (H2S), and ammonia (NH3).
The table below summarizes the reactions that produce these gases.
Table 4.1 Gas Forming Reactions
Gas
Produced in this type of reaction
H2
Metal + Acid
Example: Mg (s) + 2 HCl (aq) ---> MgCl2 (aq) + H2 (g)
CO2
Metal carbonate or bicarbonate + Acid
Example: Na2CO3 (aq) + 2 HNO3 (aq) ---> 2 NaNO3(aq) + H2O (l) + CO2 (g)
SO2
Metal sulfite or hydrogen sulfite + Acid
Example: NaHSO3 + H2SO4 (aq) ---> SO2 (g) + NaHSO4 (aq) + H2O (l)
H2 S
Metal sulfide + Acid
Example: FeS (s) + H2SO4 (aq) ---> FeSO4 (aq) + H2S (g)
NH3
Ammonium salt + Strong Base
Example: NH4NO3 (aq) + NaOH (aq) ---> NaNO3 (aq) + NH3 (g) + H2O (l)
EXOTHERMIC AND ENDOTHERMIC REACTIONS
During chemical reactions, bonds are broken in the REACTANTS (requiring energy).
This process is followed by bonds forming in the PRODUCTS (releasing energy). If more
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energy is released in forming bonds than is used in breaking bonds, the difference between
the two energies is the amount of energy released to the surroundings from the system.
When a reaction occurs that releases energy to the surroundings in the form of heat, it is
called an exothermic reaction. In an endothermic reaction the opposite holds true more
energy is required to break the bonds of the reactants than is released by forming bonds.
Therefore, energy is transferred from the surrounding to the system
WRITING AND INTERPRETING CHEMICAL EQUATIONS
Chemical reactions in aqueous solution are described very efficiently by writing a
chemical equation in which the component ions of dissolved ionic compounds are written
as separate ions e.g., K+ and I- rather than KI. These chemical equations are called
ionic equations. Reactions that involve spectator ions can be written in a form in which the
spectator ions are deleted (do not appear). These equations are called net ionic equations.
The following simple rules enable the expression of most net ionic equations for a
reaction:
Soluble ionic salts are written as separate ions.
Insoluble compounds are written as complete formula units with the subscript (s).
Covalent compounds (e.g., C02, H2O) are written as molecules.
The example that follows develops the net ionic equation for a precipitation reaction
between silver nitrate solution and hydrochloric acid solution.
Since the two solutes AgNO3 and HCI are obviously soluble and therefore dissolved,
the reactants are shown in equation form as
Ag+(aq) + NO3-(aq) + H+(aq) + Cl-(aq)
In this reaction an off-white precipitate comes out of solution. The precipitate must be
either AgCl(s) or HNO3(s). In this instance it is easy to make the decision because we know
that all common inorganic acids such as HNO3 are soluble in water. The precipitate must
be AgCl(s), and the spectator ions are H+ and NO3-. The full ionic equation is
Ag+(aq) + NO3-(aq) + H+(aq) + Cl-(aq) AgCl(s) + H+(aq) + NO3-(aq)
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The net ionic equation is
Ag+(aq) + Cl-(aq) AgCl(s)
The long form of this equation would be
AgNO3 (aq) + HCI(aq) AgCI(s) + HNO3 (aq)
Not only is the net ionic equation more concise, but it also suggests that any soluble Ag+
salt and any soluble Cl- salt would give the same reaction.
Hence, many chemical
reactions are summarized in this single net ionic equation.
III. Preparation for Experiment 5
Proper Use of a Desiccator
A desiccator is a container that can be sealed from the atmosphere and holds a drying
agent (desiccant) such as CaCl2. If the chemical samples placed in a desiccator contain
small amounts of water, the water is slowly removed by the drying agent. The drying
agent also removes any atmospheric water trapped in the desiccator.
The first step in setting up a desiccator is to check that the desiccant is not wet or
clumped.
If the desiccant is dry, any grease left on the ground glass seals of the
desiccator and lid should be removed by wiping thoroughly with a paper towel. Once the
desiccator seal is cleaned, a thin film of grease is spread over the ground glass surface on
the desiccator and lid. The desiccator is sealed by sliding the lid into place. The lid is then
gently moved back and forth to create a complete seal.
If the desiccant is wet or clumped, it should be discarded in a waste container
designated by the TA. The desiccator should be cleaned and dried. New solid, dry CaCl2
is available in each lab. Approximately 200 mL of new desiccant (CaCl2) should be added
and distributed evenly on the bottom of the desiccator. The ceramic desiccator plate is
placed on top of the desiccant with the shiny side facing up. Any excess grease should be
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wiped from the edges of the ground glass seals with a paper towel. Once the desiccator
seal is cleaned, a thin film of desiccator grease is spread over the ground glass surface on
the desiccator and lid. The desiccator is sealed by sliding the lid into place. The lid is then
gently moved back and forth to create a complete seal.
A hot sample must NEVER be placed in a sealed desiccator. Heated samples are
always allowed to cool in the desiccator with the lid slightly open (cracked). Remember,
hot air expands. If a hot sample is placed in the desiccator, and the lid is immediately
sealed, the lid will pop off and usually break. Desiccator lids are expensive!
Dehydration of Calcium Hydroxide Sample
1.
Wash a weigh bottle with micro solution and then rinse thoroughly with distilled
water. Using a pencil, write your initials on the cloudy top edge of the bottle.
2.
Heat the weigh bottle (without the cover) in the drying oven (110oC) for 15 minutes.
3.
Remove the weighing bottle from the oven with tongs. Let cool. Do not handle the
dried bottle (with bare hands) even once cool.
4.
Place content of 1 bag of calcium hydroxide, Ca(OH)2, into the weigh bottle.
5.
Place the weigh bottle without the cover in the drying oven for 2 hours.
6.
Carefully remove the weigh bottle from the oven and allow the sample to cool for 10
minutes on the bench top.
7.
Place the weigh bottle in a desiccator until the next laboratory.
IV. Stepwise Procedure
Heat of Solution
A temperature sensor measures the change in temperature of a solution in a calorimeter
during a chemical reaction in this experiment.
A.
Hardware Setup
1. Start the computer.
2. Check that the Science Workshop 500 interface is connected to the power source. When
the interface is properly connected, the green power light is illuminated.
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3. Connect the computer to the interface using the cable provided. Attach the cable to the
USB port of the computer.
4. Check that the temperature sensor is connected to analog port A and the pH sensor is
connected to analog port B.
B.
Software Setup
1. Open Data Studio.
2. Select CREATE EXPERIMENT.
3. The program looks for the interface to initialize. If the interface is not connected, properly
it prompts the user to SCAN or PICK the interface for initialization. Select SCAN. If this
does not work, see a TA.
4. If connected properly, three windows automatically open. On the left hand side is a DATA
window and a DISPLAYS window, both of which are parts of the SUMMARY feature. In
the center is the experimental setup window which displays the 500 interface and Sensor
options.
5. In the experimental setup window, click on Port A of the interface box, scroll down to the
Temperature Sensor and double click on it. This selects the sensor and automatically
connects it to port A (the pressure sensor icon should now be shown to be connected to that
port).
6. Click on Port B of the interface box and scroll down to the pH sensor. Double click on the
pH sensor to add it to Port B.
7. At the bottom of the Experiment Set-up screen, change the sample rate to 1 reading every 2
seconds.
8. At the top of the Experiment Setup window, select the CALIBRATE SENSORS tab. A
new window will appear. Select the pH sensor from the pull down menu at the top of the
window.
9. Obtain a wash bottle and rinse the tip of the pH electrode several times with distilled water.
Carefully dry the tip of the pH sensor with a paper towel.
10. Place the pH electrode in the high pH buffer standard solution (pH 10) and wait for 30
seconds for the voltage reading to stabilize. Enter a value of 10 in the standard value box
for calibration point 1. Click READ FROM SENSOR to record the voltage for calibration
point 1.
11. Rinse the pH electrode with distilled water, dry with a paper towel, and place the electrode
in the low pH buffer standard solution (pH 4). Wait about 30 seconds for the voltage
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reading to stabilize. Enter a value of 4 in the standard value box for calibration point 2.
Click READ FROM SENSOR to record the voltage for calibration point 2.
12. Click OK and return to the Experimental Setup window
13. Save the activity.
Select a file name with less than 10 characters and devoid of
punctuation.
C. DISPLAY SETUP
1. Create a table and graph of the temperature data for the temperature sensor in oC by
dragging the temperature icon from the DATA window on top of the table and graph icons
in the DISPLAYS window.
2. Save the activity. SELECT A FILE NAME WITH LESS THAN 10 CHARACTERS AND DEVOID
OF PUNCTUATION.
D. CALORIMETER SET UP
Obtain two styrofoam cups and a lid from the TA bench. Place one styrofoam cup inside the
other. Place a stir bar in the cup. Set the apparatus on a stir plate placed next to metal support
rod on the bench top as shown in Figure 4.8. Using a clamp, fix the temperature sensor
suspended in the styrofoam cup without touching the base of the cup.
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E. Data Collection
a. Obtain a bag of ammonium chloride and measure ~1 gram of ammonium chloride in a
weigh boat. Record the mass with three decimal places in DATA TABLE I.
b. Place 30 mL of distilled water in the styrofoam cup and cover the cup with a lid.
b. When everything is ready, turn on the stir plate and click the START button in DataStudio
to begin data recording.
c. After five seconds, lift one edge of the lid and add the ammonium chloride to the
styrofoam cup. Immediately replace the lid on the cup.
d. Continue to collect data until the temperature does not change any further.
e. Click on the STOP button to end data recording. Run #1 should appear in the list in the
DATA window. Rename this run ammonium chloride.
f.
Remove the temperature sensor and rinse the sensor with distilled water.
g. Discard the solution in the cup down the sink and rinse the cup with distilled water.
h. Save the activity.
i.
Repeat the procedure with ~1 gram of sodium carbonate instead of ammonium chloride.
Rename this run sodium carbonate
7. Data Analysis
a. Click on the Statistics button () in the left hand area of the Table display.
b. In the Statistics area that opens at the bottom of the Table, use the minimum and maximum
values for both runs as it corresponds to the initial and final temperature readings and
record the data in DATA TABLE I.
c. Calculate change in temperature by subtracting initial temperature final temperature.
Then calculate moles of compound by dividing mass by the molecular weight of the
compound . Find molar heat of solution (T/mol) for ammonium chloride and sodium
carbonate and record all of the calculated values in DATA TABLE 1.
Heat of Reaction
1. Reacting Sodium Bicarbonate with Citric Acid.
a. Place 30 mL of 0.5 Molar citric acid into the styrofoam cup. Place the lid on the cup and
place the temperature sensor into the citric acid solution.
b. Weigh out ~3.5 g of solid sodium bicarbonate in a weigh boat and record the mass with
three decimal places in DATA TABLE II.
b. When everything is ready, turn on the stir plate and click the START button in DataStudio
to begin data recording.
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c. After 20 seconds have elapsed, add the sodium bicarbonate to the citric acid solution.
d. Record data until a minimum temperature has been reached and the temperature reading
begins to increase.
e. Click on the STOP button to end data recording. Run #3 should appear in the list in the
DATA window. Rename this run sodium bicarbonate + citric acid.
f.
Remove the temperature sensor and rinse the sensor with distilled water.
g. Discard the solution in the styrofoam cup down the sink and rinse the cup with distilled
water.
h. Save the activity.
2. Reacting Hydrochloric Acid and Magnesium.
a. Place 30 mL of 1 Molar HCl into the styrofoam cup. Place the lid on the cup and place the
temperature sensor into the citric acid solution.
b. Obtain a piece of shiny magnesium metal from your TA and determine its mass to three
decimal place. Record this value in DATA TABLE II.
b. When everything is ready, turn on the stir plate and click the START button in DataStudio
to begin data recording.
c. After 20 seconds have elapsed, add the magnesium ribbon to the hydrochloric acid
solution.
d. Record data until a maximum temperature has been reached and the temperature reading
begins to decrease.
e. Click on the STOP button to end data recording. Run #4 should appear in the list in the
DATA window. Rename this run magnesium + hydrochloric acid.
f.
Remove the temperature sensor and rinse the sensor with distilled water.
g. Discard the solution in the styrofoam cup down the sink and rinse the cup with distilled
water.
h. Save the activity.
3. Data Analysis
a. Click on the Statistics button () in the left hand area of the Table display.
b. In the Statistics area that opens at the bottom of the Table, find the minimum and
maximum values for both run and record the data in DATA TABLE II.
c. Calculate change in temperature by subtracting initial temperature final temperature.
Then calculate moles of compound by dividing mass by the molecular weight of the
compound . Find molar heat of solution (T/mol) for sodium bicarbonate in citric acid and
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magnesium in hydrochloric acid and record all of the calculated values in DATA TABLE
II.
Antacid Titration
A. Display Setup
1. Click and drag the pH icon in the DATA window on top of the Graph icon in the
DISPLAYS window. This will create a graph of pH vs. time.
2. Save the activity.
B. Sampling Options
1. In this experiment, the run will stop automatically after 1000 seconds. To set-up this
sampling option click on the SAMPLING OPTIONS button in the Experiment Setup
Window.
2. The SAMPLING OPTIONS window will open. Click on the AUTOMATIC STOP tab.
3. Select Time and enter 1000 seconds. Click OK to save the settings.
C. Data Collection Whole Antacid Tablet
1. Select ONE of the following antacids and obtain the indicated number of tablets: Equate
(2 tablets); Rolaids (3 tablets); or Tums (3 tablets).
2. Write down the active ingrediates for the antacid in DATA TABLE III.
3. Using a graduated cylinder, measure 150 mL of 0.1 Molar hydrochloric acid (HCl) into a
clean and dry 400-mL beaker.
4. Put a stir bar in the beaker and place the beaker on the magnetic stirrer.
5. Use the 3 prong clamp on the support rod to postion the pH electrode in the HCl solution
so that is will not interfere with the stir bar.
6. Turn on the magnetic stirrer.
7. Click the START button to begin recording data.
8. After 10 seconds, drop the antacid tables into the solution.
Data recording will
automatically stop after 1000 seconds. Run #5 will appear in the Data window. Rename
this run whole antacid tablet.
9. When data Recording stops, turn off the stirrer. Remove the pH sensor from the solution.
10. Rinse the pH sensor in distilled water and gently dry the sensor.
11. Dispose of the solution down the drain.
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12. Save the activity.
D. Data Collection Crushed Antacid Tablet
1. Obtain the indicated number of tablets for the antacid you selected above: Equate (2
tablets); Rolaids (3 tablets); or Tums (3 tablets). Use the bottom of a scintillation vial to
crush the tablets in a weigh boat.
2. Using a graduated cylinder, measure 150 mL of 0.1 Molar hydrochloric acid (HCl) into a
clean and dry 400-mL beaker.
3. Put a stir bar in the beaker and place the beaker on the magnetic stirrer.
4. Use the 3 prong clamp on the support rod to postion the pH electrode in the HCl solution
so that is will not interfere with the stir bar.
5. Turn on the magnetic stirrer.
6. Click the START button to begin recording data.
7. After 10 seconds, drop the antacid tables into the solution.
Data recording will
automatically stop after 1000 seconds. Run #6 will appear in the Data window. Rename
this run crushed antacid tablet.
8. When data Recording stops, turn off the stirrer. Remove the pH sensor from the solution.
9. Rinse the pH sensor in distilled water and gently dry the sensor.
10. Dispose of the solution down the drain.
11. Save the activity.
E. Data Analysis
1. In the Graph display, click on the AUTOSCALE button to resize the graph to fit the data.
2. Click on the SMART CURSOR button. The cursor will change to a cross-hair when you
move it into the data area of the graph display. The X- and Y-coordinates will appear
beside the crosshair.
3. Move the cursor to the place in the graph of the whole tablet where the pH begins to rise.
Record the Y-coordinate as the initial pH value, pH1, in DATA TABLE III. Record the Xcoordinate as the time, T1.
4. Move the cursor to the place in the graph when the pH begins to level off. Record the Ycoordinate as the final pH value, pH2, and the X-coordinate as the final time, T2 in DATA
TABLE III.
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5. Determine the difference in pH by subtracting the initial pH value from the final pH value.
Determine the difference in time by subtracting the initial time from the final time.
Convert the time to minutes and record the values in DATA TABLE III.
6. Calculate the change in pH per minute by dividing the difference in pH by the change in
time.
7. Repeat the data analysis procedure for the crushed tablet and record your final values
(pH/Time for the whole and crushed tablets) on the board.
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