Chemistry, A molecular approach

Determine the Masses of pieces of Mg.Data on the pictures.  Name:______________________________
Montgomery College – Takoma Park – Silver Spring Campus
Chemical and Biological Division
Principles of Chemistry 1 – CHEM131 –
Determination of the Masses of Pieces of Magnesium
In this experiment the masses of pieces of magnesium ribbon will be determined indirectly. Magnesium
reacts with hydrochloric acid to produce aqueous magnesium chloride and hydrogen gas. From the volume of
hydrogen gas collected, the mass of magnesium can be calculated using Dalton’s Law of Partial Pressures
and the Ideal Gas Law.
Mg (s) + 2 HCl (aq)  MgCl2 (aq) + H2 (g)
Obtain a gas collecting tube and a container with the materials needed to make the “bag” for anchoring
the magnesium ribbon.
1. Pour 5 mL of 6M HCl into the empty gas collecting tube. Carefully fill the tube to the top with deionized
water, so the HCl remains near the bottom. Clamp the tube to a ring stand.
2. Pour tap water into an 800 mL or 1000 mL beaker to within an inch of the top.
3. Take one of the pieces of magnesium ribbon and enclose it in a fiberglass bag, as follows:
a. Coil the Mg strip around the end of a pencil.
b. Take the coil off the pencil and adjust the coil so that no part of the col touches any other part.
c. Place the coil in the center of the fiberglass square and bring the 4 corners of the square together
so that the corners from an envelope enclosing the magnesium coil.
d. Stick one end of the copper through all four corners of the fiberglass and twist it to hold the
magnesium coil in the closed “bag.”
4. Insert the bag with the wire into the top of the gas collecting tube containing the water and acid and
hold it in place with a one-holed rubber stopper. The bag should be about ½” below the bottom of the
stopper and the end of the copper wire should extend from the bag between the stopper and the tube.
When pressing the stopper into the tube, do not cover the hole of the stopper completely with your
finger so the excess air and water can escape from the tube.
5. Place a finger over the end of the tube and invert it into the beaker of water. Be sure to hold the open
end of the tube below eye level while inverting. Clamp the inverted tube to a ring stand, so that the
bottom of the tube is 1/8” above the bottom of the beaker. The more dense hydrochloric acid moves
down the tube, mixing with the water and then coming into contact with the magnesium.
6. Allow the magnesium to react completely with the hydrochloric acid. After all the magnesium has
reacted, record the volume of the hydrogen gas (to the nearest 0.01 mL), temperature of the water (to
the nearest 0.1oC), and the height of the water column above the water leel in the beaker (to the
nearest 1 mm).
Page 1 of 4
7. Record the barometric pressure.
8. Repeat this experiment with the second piece of magnesium ribbon.
9. Untwist the wire from the piece of fiberglass. Rinse the wire, fiberglass square, and stopper; place them
in the container, and return the container to the instructor’s desk.
The table of temperature/vapor pressure of water is in your textbook and in the back of your laboratory
manual. The density of mercury is 13.6 g/mL.
Data Sheet
Unknown Number
Sample 1
Sample 2
Volume of H2 gas (mL)
Temperature of water (oC)
Height of water column (mm)
Vapor pressure of water at above temperature (mmHg)
Barometric pressure (mmHg)
Calculations (Show calculations)
Pressure of the column of water (mmHg)
Pressure of dry hydrogen (mmHg)
Moles of Hydrogen
Mass of Magnesium ribbon (g)
Average mass of magnesium ribbon (g)
Temperature (oC)
Pressure (mm Hg)
Temperature (oC)
Pressure (mm Hg)
Page 2 of 4
Post-Lab Questions
1. If your partner did not take into account the vapor pressure of water when calculating the moles of
hydrogen produced, would the mass of Mg that was calculated at the end of the experiment be too high
or too low?
2. If the strips of Mg ribbon are coated with white MgO, would this cause the mass of Mg that you
calculate to be too high or too low? Why?
3. Instead of a strip of pure Mg ribbon, you are accidentally given a strip of an aluminum-magnesium alloy.
You calculate the mass of the strip as above. Would the mass of the strip be high or lower than
expected? Write the appropriate chemical equations (one equation for the reaction of Mg with HCl and
another equation for the reaction of Al with HCl) and explain your answer.
Page 3 of 4
Pre Lab Assignment – Due at the start of the lab.
1. What volume of hydrogen, measured at 25oC and 1.00 atm, can be produced by dissolving 5.0 g of
magnesium in hydrochloric acid? (Write and balance the equation for the reaction first.)
2. A sample of iron reacted with hydrochloric acid. The liberated hydrogen occupied 40.1 mL when
collected over water at 27oC at a barometric pressure of 750 torr. What is the mass of the sample of
iron? (Because of the collection of gas over water, the vapor pressure of the water must be subtracted
from the barometric pressure to find the pressure of the hydrogen gas.)
3. If you wanted to prepare 50.0 mL of hydrogen, collected over water at 25oC on a day when the
barometric pressure was 730 torr, what mass of aluminum would you react with a hydrochloric acid?
(Balance the equation first.)
_____ Al (s) + _____HCl (aq)  _____AlCl3 (aq) + _____H2 (g)
Page 4 of 4
Chemistry is the science that studies matter and the changes it undergoes. Changes
are divided into two categories: physical and chemical. During a physical change, some
physical property of the substance changes, although the identity of the substance remains
the same. For example, the melting of ice to make liquid water is a physical change:
H2O (s) → H2O (l)
However, when a piece of sodium metal is placed in water, a chemical change occurs. The
element sodium reacts with water to make the compound sodium hydroxide and the element
2 Na (s) + 2 H2O (l) → 2 NaOH (aq) + H2 (g)
When the process is complete, the original substances no longer exist. New substances,
with new properties have been made.
Every change, whether physical or chemical, is accompanied by a change in the
internal energy (E) of the system. In the case of melting ice, the molecules absorb heat from
the surroundings. Their internal energy increases as they liquefy. In the reaction between
sodium and water, the internal energy of these two substances decreases as sodium
hydroxide and hydrogen are produced. The change in internal energy is symbolized by ∆E.
The first law of thermodynamics defines the change in internal energy as ∆E = q + w,
where q refers to the heat exchanged between reaction system and surroundings, and w
refers to the work done by or on the reaction system. A sign convention here is important.
With the reaction system as the point of reference, positive signs are used for q and w if the
system absorbs heat or is worked upon. Negative signs are used if the reaction system
loses heat or does work.
To measure ∆E values, a sealed bomb calorimeter is required. In such a calorimeter,
no work can be done. Therefore, w = 0 and ∆E = q. This q is often labeled qV since the
reaction is carried out at constant volume. In lab, simple calorimeters are usually used. They
are open to the atmosphere. The reaction system can exchange both heat and work with the
surroundings. However, the work component is difficult to measure. Furthermore, it is
usually small compared to the heat component. For that reason, we ignore the work
component and measure just the heat component. Since the atmospheric pressure is
essentially constant for the duration of the reaction, we label the heat as q p. This value, qp, is
not equal to ∆E (since ∆E = q + w, and w is not 0 in this case). The qp term has been named
∆H and symbolizes the change in enthalpy. Enthalpy is similar to, but not the same as,
internal energy. It is often called the “heat content” of a substance.
Because heat is a form of energy, an amount of heat is expressed in units of energy,
such as Joules (J) or calories (cal). Heat will transfer from one object to another because of a
temperature difference between the two objects. Heat flows from hot to cold. The energy
transfer stops when the objects have reached the same temperature. The amount of heat lost
equals the amount of heat gained.
When heat flows into an object, its temperature increases. The rate at which the
temperature increases with each amount (J or cal) of heat is that object’s heat capacity. When
considering a uniform substance, the total heat capacity is proportional to the total mass of the
substance. The heat capacity of each unit of mass (g) of the substance is called the specific
heat capacity, or simply specific heat. Specific heat is usually considered to be the amount of
heat required to raise the temperature of 1 g of a substance by 1 ºC. It is a property unique to
each substance depending on its molecular composition and structure. It is commonly denoted
as c.
The unit of energy, calorie, is defined as the amount of heat required to raise the
temperature of 1 g of water 1 ºC, which means the specific heat of water is 1 cal/gºC. Since 1
cal = 4.184 J, the specific heat of water is 4.184 J/gºC. Water has a relatively high specific heat
for a pure substance due to its extensive hydrogen bonding. The hydrogen-bonding network
allows numerous modes of vibration among all the water molecules, giving water the ability to
store more energy than most other substances while only slightly increasing its temperature.
The specific heat of metals, on the other hand, will generally be much lower, below 1 J/gºC
since they have much simpler atomic structure. For example, the specific heat of iron is 0.4605
J/gºC and that of gold is 0.1256 J/gºC. Interestingly, the relationship between the specific heat
of a metal and its molar mass is also simple. In 1819, the physicists Pierre Dulong and Alexis
Petit discovered that roughly 25 J are required to raise the temperature of 1 mole of most
metals by 1 ºC. This relationship is known and the Law of Dulong and Petit. It is expressed as:
𝑀𝑀 ≅
where MM = molar mass and c = specific heat in units of J/ºC. This law was discovered only a
few years after John Dalton developed the concept of atomic mass and it was one of the few
rules available to early chemists in the study of molar masses.
Heat can be directly measured, but it is more convenient to measure it indirectly by
first measuring a temperature difference and then calculating the amount of heat
transferred. To calculate heat, q, we use the equation:
q = mc∆T
where m = mass, c = specific heat, and ∆T = change in temperature, defined as Tfinal – Tinitial.
If Tfinal is higher than Tinitial, then ∆T is positive and q is positive. An object with a positive q
has gained heat during the energy transfer. If Tfinal is lower than Tinitial for an object, then its
∆T is negative, causing its q to be negative. A negative q value for an object signifies that it
has lost heat during the energy transfer.
Heat exchanges between system and surroundings are measured in special
containers called calorimeters. They may be as sophisticated as a bomb calorimeter, or
simple as a Styrofoam coffee cup or thermos. The system is whatever substances undergo
the change. The surroundings are all the other contents of the container as well as the
container itself. If the system loses heat, all of it goes to the surroundings.
Once ∆H data has been determined for a reaction, it is often written at the end of
the balanced equation. Its value corresponds to the balanced equation when interpreted
in terms of moles. For example:
H2O (s) → H2O (l)
∆H = +6.0 kJ
2 Na (s) + 2 H2O (l) → 2 NaOH (aq) + H2 (g) ∆H = –367 kJ
In the first case, when one mole of solid water (ice) undergoes melting, 6.0 kJ of heat
are absorbed by the water molecules. In the second case, when two moles of solid sodium
react with two moles of liquid water, two moles of aqueous sodium hydroxide and one mole
of gaseous hydrogen are produced, and 367 kJ of heat are given off to the surroundings.
Note that there is no such thing as positive or negative heat. The positive and negative signs
simply signify the direction of heat flow, into or out of the system.
Sometimes a chemical product is made in one step; in other cases, several steps
may be required to make the desired product. For example, substance “E” can be made as
(Step 1)
A+B→ C
(Step 2)
The net reaction can be found by adding the two steps:
(Net) A + B + D → E
The C’s in steps 1 and 2 cancel. The value of ∆H for the net reaction can be found by adding
the ∆H values for each step. If the values are:
(Step 1)
∆H = +100. kJ
(Step 2)
∆H = –300. kJ
then the value of ∆H for the net reaction is –200. kJ.
In the first part of this lab, we will determine the specific heat of a metal. In the
second part, we will determine the value of ∆H for two processes. One is a chemical
change, the other is a physical change. Then, we will use this lab data as well as additional
data for two other reactions to determine ∆H for a third reaction by means of Hess’s law.
A. Determination of the specific heat of a metal
1. Fill a 400-mL beaker about ¾ with deionized water and begin heating it to a boil.
2. While the water is heating, weigh your sample of unknown metal in a large stoppered
test tube.
3. Pour the metal into a dry container and weigh the empty test tube and stopper.
4. Place the metal back into the test tube and loosely place the stopper into the top.
5. Place the test tube into the hot water in the beaker. The water level in the beaker should
be higher than the top of the metal sample.
6. Wait for the water to begin boiling then keep the test tube in boiling water for at least 10
minutes to ensure that the metal reaches the temperature of the boiling water. Add small
amounts of water as necessary to maintain the level of the water above the metal.
7. While the water is boiling, weigh your calorimeter. Place about 40 mL of deionized water
in the calorimeter and weigh it again.
8. Cover the calorimeter and insert a digital thermometer through a hole in the lid. Make
sure the tip of the thermometer is completely immersed in the water and then measure
the temperature of the water in the calorimeter to the nearest 0.1 ºC.
9. Take the test tube containing your metal sample out of the boiling water, remove the
stopper, and pour the metal into the water in the calorimeter. Try to do this as smoothly
and quickly as possible to avoid heat loss by the metal during the transfer. Also, be
careful that no water that has adhered to the side of the test tube drops into the
calorimeter during the transfer of the metal.
10. Replace the calorimeter cover and agitate the contents.
11. Record the maximum temperature observed.
B. Determination of ∆H for the reaction HCl (aq) + NH3 (aq) → NH4Cl (aq)
1. Measure out 50.0 mL of 2.50 M NH3. Pour it into your calorimeter. Measure and
record the temperature of the solution.
2. Measure out 50.0 mL of 2.00 M HCl in a clean, dry graduated cylinder. Measure and
record its temperature. NOTE: The temperatures of the two solutions are probably
almost the same. If they are not, use the average of the two as the initial temperature
of the chemical system.
3. Add the HCl solution to the NH3 solution in the calorimeter. Stir, and record the
highest temperature reached.
4. Pour the contents of the calorimeter down the drain. Dry the calorimeter with a paper
towel. When finished, dry the calorimeter before beginning part B.
C. Determination of ∆H for the process NH4Cl (s) → NH4Cl (aq)
1. On the balance, weigh out a 4.500 – 5.000 g sample of solid ammonium chloride,
NH4Cl. USE YOUR SPATULA. Record the actual mass you use.
2. Place 50.0 mL of room temperature water in the calorimeter. Measure and record its
3. Add the NH4Cl to the water. Stir and record the lowest temperature reached.
4. Pour the calorimeter contents down the drain.
Name ________________________________________________ Date ______________
Be sure to use the correct number of significant figures for the volumes and
temperatures you record below!
A. Determination of the specific heat of a metal
Unknown number
Mass of stoppered test tube and metal
___________ g
Mass of test tube and stopper
___________ g
Mass of calorimeter
___________ g
Mass of calorimeter and water
___________ g
Mass of water
___________ g
Mass of metal
___________ g
Initial temperature of water in calorimeter
__________ ºC
Initial temperature of metal (assume 100 ºC
unless directed to do otherwise)
__________ ºC
Equilibrium temperature of metal and water
in calorimeter
__________ ºC
B. Determination of ∆H for the reaction HCl (aq) + NH3 (aq) → NH4Cl (aq)
volume of 2.50 M NH3
________ mL
temperature of 2.50 M NH3
volume of 2.00 M HCl
________ mL
temperature of 2.00 M HCl
average temperature
temperature of mixture (Tfinal for solution)
C. Determination of ∆H for the process NH4Cl (s) → NH4Cl (aq)
Volume of H2O
________ mL
mass of NH4Cl (s)
________ g
temperature of water (Tinitial for solution)
temperature of mixture (Tfinal for solution)
Name ________________________________________________ Date ______________
Show your work clearly, in the space provided, for each of the following calculations. Enter
the answer on the line provided. The specific heat of water is 4.184 J/gC o. Assume this is
also the specific heat of the dilute aqueous solutions we use here.
A. Determination of the specific heat of a metal
ΔTwater (Tfinal – Tinitial)
__________ ºC
__________ ºC
___________ J
Specific heat of the metal
________ J/gºC
Approximate molar mass of the metal
________ g/mol
B. Determination of ∆H for the reaction HCl (aq) + NH3 (aq) → NH4Cl (aq)
mass of solution
__________ g
(its density is 1.03 g/mL)
q for solution
(q = mcΔT)
__________ J
q for reaction
__________ J
The amount of heat gained by the solution equals the amount of heat
lost by the reaction system. Watch your signs!
moles of NH3 used
__________ mol
Remember, you know the volume and the molarity.
moles of HCl used
__________ mol
Is there a limiting reactant?
Name ________________________________________________ Date ______________
q for one mole of limiting reactant
__________ J/mol
_________ kJ/mol
(Hint: this is just a unit conversion.)
C. Determination of ∆H for the reaction process NH4Cl (s) → NH4Cl (aq)
mass of solution
__________ g
(NH4Cl and water. Assume the density of water is 1.00 g/mL.)
q for solution
(q = mcΔT)
__________ J
q for dissolution reaction
__________ J
Watch your signs!
moles of NH4Cl used
__________ mol
q for one mole of NH4Cl
__________ J/mol
_________ kJ/mol
(Hint: this is just a unit conversion.)
Name ________________________________________________ Date ______________
Application of Hess’s Law to the Lab
You have now determined the value of ∆H for a chemical change and a physical change
involving ammonium chloride. Use your lab data as well as the data given below to determine
the value of ∆H for the decomposition reaction. Show all your work below.
NH4Cl (s) → NH3 (g) + HCl (g)
Consider the following as you make your determination:
NH4Cl (s) → NH4Cl (aq)
∆H = ____________ kJ
HCl (aq) + NH3 (aq) → NH4Cl (aq)
∆H = ____________ kJ
NH3 (g) → NH3 (aq)
∆H = –34.6 kJ
HCl (g) → HCl (aq)
∆H = –75.1 kJ
Pre-Lab Assignment
Name ________________________________________________ Date ______________
How much heat is lost when 75.0 g copper (c = 0.385 J/goC) at 57.2oC are cooled to
Using the data from the table in Appendix II B in your textbook, or some other
reference source, determine the value for the enthalpy change (ΔH) for reaction in
Part B and reaction in Part C of this lab. (Refer to section 6.9 in your textbook if
necessary). Make sure that you use the value for the correct form of the substance
[For instance, ΔHf° HCl (aq) is different than ΔHf° HCl (g)].

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