Investigation To Find Out The Empirical Formula


Investigation To Find Out The Empirical Formula Of Magnesium Oxide. Essay, Research Paper

`Aim. To find out

the empirical formulae for magnesium-oxideIntroduction. The

empirical formula of a compound tells us the types of atoms present in a

compound as well as the simplest whole-number ratio of the different types of

atoms. The empirical formula does not tell us the actual number of atoms in the

molecule. The

mass of Mg + the mass of O2=mass of MgxOx. Knowing the mass of magnesium used

and the mass of magnesium oxide produced you can determine the mass of oxygen

used. The ratio between number of moles of magnesium used and number of moles

of oxygen used can be calculated and the empirical formula can be written on

the basis of this ratio. Prediction.?? I predict that magnesium can and

will join with oxygen to become magnesium oxide (MgO) because magnesium is in

group 2 and can has ?2 bands. Both magnesium and oxygen need to join each other

to lose or gain electrons. Magnesium: ? Oxygen: Mg²+O²¯=MgOWhen magnesium and oxygen coincide they

cancel each other out and therefore are completely compatible.Apparatus. 1.

Bunsen burner 2.

Tripod 3.

Gauze 4.

Crucible and Lid 5.

Tongs 6.

Heatproof mat 7.

Safety glasses 8.

MagnesiumDiagram. ? Method. 1. Place a

clean, dry crucible (check for any cracks) with its cover on a clay triangle

supported by a tripod. Adjust the flame of a Bunsen burner so that the crucible

is just above the hot part of the flame. Heat the crucible strongly for five

minutes.? The bottom of the crucible

should glow red during the heating.2. Remove

the crucible, using the crucible tongs, and place it on the ceramic pad to

cool.? After the crucible has cooled to

room temperature (this will take about ten minutes), weigh it with its cover on

an analytical balance to 0.001 g. Record the weight on the data sheet. 3. Weigh

out 0.3g of magnesium and place it into the crucible.? Weigh the magnesium, crucible and cover on an analytical

balance.? Record the weight to 0.001 g. 4.

At frequent intervals, remove the burner and slightly lift the crucible lid

using tongs, quickly replacing the lid so as to lose as little magnesium oxide

’smoke’ as possible. Repeat the process patiently until the magnesium ceases to

flare up, then remove the lid and heat strongly to make sure that combustion is

complete and all the metal has been converted into oxide. You should not be

able to see any shiny metallic surfaces. Allow the crucible to cool then

replace the lid and reweigh the whole. Record the mass. 5.

Repeat this 5 times so that you will be able to gain an average result.? Fair Test. To

make this a fair test we weighed the empty crucible and lid before

adding magnesium. This means that the different weights of the crucibles could

be accounted for. This allowing any minor errors in the experiments result to

be averaged out. To make this a lot more accurate we carried the test out 5

times. Before

each time we carried out the experiment we cleaned the crucible and lid

thoroughly making sure that there was no excess foreign materials of any sort

to tamper with the chemical reaction.Variables. – The amount of magnesium used in the experiment. – How long it had been left burning for. – How much air you let in it when lifting up the lid. If I were to do this experiment again I would carry it out

in a fume cupboard to make sure nothing but pure oxygen will be allowed to come

in contact with the burning magnesium. ? Potential reasons for error Hot

Magnesium will react with Nitrogen in the air if there is insufficient O2

forming Magnesium Nitrate (Mg3N2). When

magnesium metal is burned in pure oxygen, the only product is magnesium

oxide. If a sample of Mg is weighted before and after combustion,

then the increase in mass is equal to the mass of oxygen that is combined with

Mg. From a knowledge of the mass of Mg and the mass of combined O, the

empirical formula of magnesium oxide can be calculated. There

is only one problem with this straight forward approach. It is more usual

to burn things in air than in pure oxygen. Air is about 80%

nitrogen. So burning Mg in air results not only in the formation of

magnesium oxide, but also a small amount of a "byproduct", magnesium

nitride. Not all the Mg was burned. There is not

much to do about this, just make sure that you have followed the method

correctly and look out for all the signs of the magnesium being present. I

recommend doing this experiment in a fume cupboard so that any products of the

combustion do not escape with the gas as you lift up the lid of the crucible to

let oxygen in. at lest it will stay in one concentrated area and not around the

rest of the room where it won?t get used at all.

Safety. ????????? As

we were dealing with chemicals and heat during the course of this test it was

vital that it was carried out with great caution. To achieve this we: – Made sure that the work surface was

clean and tidy with no unnecessary objects lying around on it. – Were cautious about burning magnesium as

it burns a bright white light, which is possible to blind any one looking straight

at it.? – Were careful to wear out protective

glasses on at all times. – Used tongs to move the equipment about

as we were dealing with hot things. – Stood up during the experiment. – Never left the apparatus unattended and

kept a close eye on it at all times noting the changes taking place. – Put the Bunsen burner on the safety

flame when not in use to prevent any mishaps – Made sure we didn?t inhale any ?smoke?

when the magnesium was being burned. Results. Crucible+lid

(g) Crucible+lid+Mg

(g) Mg

(g) Crucible+lid+MgO

(g) O

(g) 1 19.9 20.2 0.3 20.4 0.2 2 25.2 25.6 0.4 25.9 0.3 3 24.8 25.1 0.3 25.3 0.2 4 16.3 16.6 0.3 16.8 0.2 5 11.8 12.1 0.3 12.3 0.2 Total

Weight (grammes) 1.6 1.1 RAM 24 16 Molecules 0.07 0.069 Conclusion. From my results I can tell that 0.067 molecules of magnesium had

combined with 0.069 moles of oxygen, I predict that 1 mole of magnesium will

combine with 0.069/0.067 or 1.031 moles of oxygen. I think that if this experiment had been carried out in a proper

laboratory with no interferences you would find that the ratio for this

combustion would be very close to 1:1. Which would justify my prediction to be



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