each of these properties. In Group 2 of the periodic table, chart of the elements arranged according to the periodic law discovered by Dmitri I. Mendeleev and revised by Henry G. J. Moseley. In the periodic table the elements are arranged in columns and rows according to increasing atomic number. It reacts very slowly with cold water. It is not affected by dry air but tarnishes in moist air, forming a thin protective coating of basic magnesium carbonate, MgCO3·Mg(OH)2. When heated, magnesium powder or ribbon ignites and burns with an intense white light and releases large amounts of heat, forming the oxide, , common name for the chemical compound magnesium oxide, MgO. It occurs as colorless, cubic crystals. It is refractory, melting at about 2,800°C. It is very slightly soluble in pure water but is soluble in acids and solutions of ammonium salts. Magnesium reacts with the halogens and with almost all acids. It is a powerful reducing agent and is used to free other metals from their anhydrous halides.
The aim of this experiment is to determine the value of x in the following equation:
Mg + XHCL → MgCLx + X/2 H2
A known amount of magnesium is reacted with a large excess of HCL, and the volume of H2 evolved is measured. As HCL is in excess, all the magnesium will be consumed, and the yield of both MgCLx and H2 depend only on the amount of magnesium. A comparison of the amount of hydrogen produced with the amount of magnesium consumed will enable X to be determined.
APPARATUS AND MATERIALS:
Magnesium ribbon, HCL (0.5M), burette (50cm³), pipette (25cm³), retort stand, electrical balance, watch glass, beaker (1000cm³), gauze, funnel, rubber band, glass rod, thermometer, barometer.
EXPERIMENTAL PROCEDURES:
- Burette was used upside down to collect the hydrogen, but there is an unmarked space between the 50cm³ mark and the tap of unknown volume. The volume of this unmarked space is determined in a clean, dry 50 cm³ burette by pipetting 25.00 cm³ of water into the vertically clamped burette. The burette reading was noted, the burette was drained and repeated. The water in the burette was left for 10 min and occurrence of leaks was checked.
- The piece of magnesium ribbon was cleaned with steel wool. A piece was cut off with scissors within the lengths shown. The ribbon was curled up. A watch glass was teared on the four decimal balance and the magnesium ribbon on the watch glass was accurately weighed between 0.0300 and 0.0360 g, which was then placed inside a 600 cm³ beaker.
- A small filter funnel with a short stem (1.0 – 1.5 cm long) was covered with gauze. It was inverted and placed on the watch glass over the magnesium.
- The beaker was filled with tap water until approximate level of 0.5 – 1.0 cm above the end of the funnel stem. The burette was filled completely with 0.5 M HCL, it was inverted and placed in the water in the beaker, the cork was removed and end of the burette was placed over the stem of the funnel, no entrance of air was ensured.
- The excess off water was removed with a pipette until the level of water was just above the stem of the funnel.
- About 100 cm³ of 0.5 M HCL was added to the beaker, it was stirred using glass rod to ensure complete mixing such that the HCL reached magnesium.
- The solution was stirred to initiate the reaction and it was not stirred further so that the reaction proceeds unaided. After 30 minutes, the watch glass was tapped gently to dislodge any gas bubbles.
DATA/RESULTS:
Mass of magnesium ribbon: 0.0293 g
Volume of small tap space: 28.8 – 25.0 = 3.8 ml
Initial volume of HCl in burette = 40.9 ml
Final volume of HCl in burette = 19.4 ml
Volume oh H2 evolved: 40.9 – 19.4 – 3.8 = 17.7 ml
Initial temperature of the solution: 25°C
Final temperature of the solution: 23°C
DISCUSSION:
The reaction is between a solid magnesium ribbon and a liquid HCL. The reaction happens at the interface. The reaction rate is therefore dependent on surface area of the magnesium ribbon which is steadily growing smaller. The reaction generates bubbles on the surface that grow in size and effectively shield a significant area of the magnesium from the solution before detaching. Finally, unless the reaction is very strongly stirred, you are relying on the fairly weak currents generated by the upwelling bubbles to deliver fresh HCl solution to the surface of the Mg ribbon. Therefore, we must stir it well. After adding the HCL, the temperature is not taken for at least 20 minutes. This is to allow the HCL to reach the magnesium. So that when we stir the solution becomes even. If the temperature is taken immediately after the HCL is poured, we cannot get accurate reading of temperature because HCL would not have reached the magnesium. Any reaction will not happen yet. Therefore, temperature is taken after a while because to wait for the reaction between HCL and magnesium to occur. That chemical reaction might evolve heat. So, to get accurate and stable reading of temperature we have to stir the mixture and wait for even mixing.
The moles of hydrogen present:
Theoretical value of the moles of hydrogen gas, H2 is 2.
Bt according to our experiment;
pV = nRT
(101.3 kPa) × (17.7ml) = n × (8.3145 J/mol K) × (273.15 K + 23°C)
(101.3 kPa) × (17.7ml) = n × (8.3145 J/mol K) × (296.15 K)
n = 2462.339 ÷ 1793.01
n = 1.37 mole
Therefore, experimentally; The moles of hydrogen gas, H2 present are 1.37 moles.
The Ideal gas law is the of a hypothetical . It is a good approximation to the behaviour of many under many conditions, although it has several limitations. It was first stated by in 1834 as a combination of and . It can also be derived from , as was achieved (apparently independently) by in 1856 and in 1857. The of an amount of is determined by its pressure, volume, and temperature. The modern form of the Ideal gas law equation is:
pV = nRT
where;
p is the absolute of the gas (101.3 kPa );
V is the of the gas;
n is the of the gas, usually measured in ;
R is the (which is 8.314472 −1−1 in );
T is the . Standard temperature: 0°C = 273.15 K
Since it neglects both molecular size and intermolecular attractions, the ideal gas law is most accurate for gases at high temperatures and low pressures. The neglect of molecular size becomes less important for larger volumes, i.e., for lower pressures. The relative importance of intermolecular attractions diminishes with increasing i.e., with increasing temperatures. More sophisticated , such as the , allow deviations from ideality caused by molecular size and intermolecular forces to be taken into account.
A mole (abbreviated mol) of a pure substance is a mass of the material in grams that is numerically equal to the molecular mass in atomic mass units (amu). A mole of any material will contain Avogadro's number of molecules. For example, carbon has an atomic mass of exactly 12.0 atomic mass units -- a mole of carbon is therefore 12 grams. For an of a pure
element, the is approximately equal to the mass in amu. The accurate masses of pure elements with their normal isotopic concentrations can be obtained from the .
One mole of an ideal gas will occupy a volume of 22.4 liters at STP (Standard Temperature and Pressure, 0°C and one pressure).
Avogadro's number;
STP is used widely as a standard reference point for expression of the properties and processes of ideal gases. The standard temperature is the freezing point of water and the standard pressure is one standard atmosphere. These can be quantified as follows:
Standard temperature: 0°C = 273.15 K
Standard pressure = 1 atmosphere = 760 mmHg = 101.3 kPa
R = universal gas constant = 8.3145 J/mol K
Where;
pV = nRT
(101.3 kPa) × V = (1 mole) × (8.3145 J/mol K) × (273.15 K)
V = 22.4 L
Therefore, it is shown that 1 mole of gas at S.T.P. occupies 22.4 L.
If hydrogen gas, H2, leaks out through the stopcock of the inverted burette, we will surely obtain an inaccurate result. The temperature will be drastically affected because the rate of reaction between magnesium ribbon and dilute HCL will change. Volume of hydrogen gas collected will change, and final volume of HCL collected will be affected because of this.
CONCLUSION:
Stoichiometry (sometimes called reaction stoichiometry to distinguish it from composition stoichiometry) is the of (measurable) relationships of the and in a balanced (). It can be used to calculate quantities such as the amount of products that can be produced with the given reactants and percent yield. In general, chemical reactions combine in definite ratios of chemicals. Since chemical reactions can neither create nor destroy matter, nor one element into another, the amount of each element must be the same throughout the overall reaction. For example, the amount of elements x on the reactant side must equal the amount of element X on the product side. Reaction stoichiometry allows us to determine the amount of substance that is consumed or produced by a reaction.
A known amount of magnesium is reacted with a large excess of HCL, and the volume of H2 evolved is measured. As HCL is in excess, all the magnesium will be consumed, and the yield of both MgCLx and H2 depend only on the amount of magnesium. A comparison of the amount of hydrogen produced with the amount of magnesium consumed will enable X to be determined. The reaction is between a solid and a liquid. The reaction happens at the interface. The reaction rate is therefore dependent on surface area of the magnesium ribbon which is steadily growing smaller. The reaction generates bubbles on the surface that grow in size and effectively shield a significant area of the magnesium from the solution before detaching. Finally, unless the reaction is very strongly stirred, you are relying on the fairly weak currents generated by the upwelling bubbles to deliver fresh HCl solution to the surface of the Mg ribbon. The reaction is fastest at the start. As the magnesium is used up, the rate falls. This can be seen on the graph, as the slope becomes less steep and then levels out when the reaction has stopped (when no more gas is produced).The reaction is exothermic, but the dilute acid is in excess and the rise in temperature is only of the order of 3.5˚C. There is some acceleration of the reaction rate due to the rise in temperature.
REFERENCES:
- lectureonline.cl.msu.edu/~mmp/applist/pvt/pvt.htm
- en.wikipedia.org/wiki/Ideal_gas_equation
- hurri.kean.edu/~yoh/calculations/idealgas/idealgas.html
