Determination of the relative atomic mass of magnesium by back titration

In this following experiment, we are aiming to determine the relative atomic mass of magnesium, using a process known as 'back titration'. The basic outlines of this experiment include a strip of magnesium being allowed to reach with excess hydrochloric acid. The excess acid is then determined by titration with a standard alkali (Sodium Hydroxide) and hence the moles of acid that reacted with the magnesium is found by difference. This can be used to determine the number of mole of magnesium used (explained more thoroughly in analysis section).

The relevant equation is:

2H+ (aq) + Mg (s) --> Mg2+ (aq) + H2 (g)

Safety Precautions

Chemical / Apparatus

Hazard

Precaution

Hydrochloric acid

Irritant

Avoid contact with skin, handle with care, wash off any spileges, wear eye protection.

Sodium Hydroxide

Irritant

See above

Methyl orange indicator

Stains

Handle with care, avoid contact with skin

Apparatus

* 250cm3 volumetric flask

* 25cm3 pipette

* Pipette filler

* Burette

* Burette clamp with stand

* 250cm3 conical flash

* White tile

Reagents

* Clean magnesium ribbon

* Standard 1.0M hydrochloric acid solution

* Indicator (screened methyl orange)

Variables and Fair Testing

A fair test must be ensured at all times, in any experiment, to keep the results as accurate as possible so that appropriate conclusions can be drawn. The main way that I hope to achieve this is by repeating each of my results a further two times so that an average can be taken and any anomalous results can be spotted before they are taken as genuine ones. As well as this I must consider how accurate I want my results to be.

Fixed Variables

Same volume of hydrochloric acid (now reacted with the magnesium): As we're attempting to determine the relative atomic mass of magnesium, then the acid it is dissolved in must be kept constant throughout the experiment, otherwise more than one variable would occur. This would therefore mean our results show nothing of value, and as such determining the excess acid by titration would be of no use.

Same drops of Indicator: The number of drops of indicator being used (methyl orange) must be kept constant throughout. The number of drops should actually be one or two, otherwise distinguishing the colour change during the titration would prove difficult, whereby the possibility of errors would increase. A white tile is also used to help us in the aid of distinguishing the colour change.

Same Apparatus: It is a necessity that all apparatus is kept the same throughout the experiment. Different apparatus differ in their accuracy. For example, 50cm3 beaker would be more accurate than that of a 250cm3. Also, all apparatus should be washed thoroughly before and after each experiment, to avoid any possible contamination.

****BRIGGS***

Same concentration of acid and alkali: In a solution whereby the concentration is high, then the particles that consist are closely packed. This results in more frequent, successful collisions. ****BRIGGS*** Never had time to finish this part off.

Varied

The volume of standard alkali (sodium hydroxide): The aim of the investigation is to determine the relative atomic mass of magnesium by back titration, and therefore, this is the only varied factor in this experiment. For our experiments to be accurate, we expect all results to be concordant, i.e. very similar, within a +- 0.1cm3 range. Results beyond or below this would be considered anomalous, and therefore will be ignored.

Sources of Error

The magnesium may not be clean. It may in parts, potentially be magnesium oxide, and also, magnesium does have a tendency of tarnishing. To prevent this, we could use emery cloth, and wipe the magnesium thoroughly.

The required mass of magnesium needed is 0.3 grams. Weighing this out by trial and error would take us too much time, not to mention how inaccurate it is. So, to overcome this problem, simple algebra is used. We would take 1m of magnesium, weigh this out (assume this value is K), and substitute this into a simple equation. This being.

(0.3 / k) x (? / 100 )

Correct use and handling of apparatus is also another key factor. For good experimental techniques, i.e. handling, setting up and using apparatus, we should always:

* Keep the burette vertically

* Ensure there is no air gap beneath the tap

* Rinse apparatus with distilled water, and then with the solution being used

* Remove the funnel when titrating

* Ensuring the value of the pipette is read at eye level

* Movement during transferring should be kept at a minimum

* Allowing it to drain under gravity

* When empty touching the end on the surface of the solution to leave the correct amount of liquid in the pipette

* Constantly swirl the conical flask whilst running in the solution from the burette.

* Towards the end, add in drops to avoid adding an excess solution

Repeating the experiment is also required. Two concordant titres are required, if not, then repeats are taken.

Method

It is always important to remember that the method integrates ALL sources off error (as stated above), to increase accuracy. The method used is just a procedure, hence is kept simple, and one must recognize that we are constantly keeping ever fair.

1. Weigh out accurately about 0.3g of the magnesium ribbon.

2. Pipette 50.0cm3 of the 1.0M hydrochloric acid into a 250cm3 volumetric flask.

3. Break the magnesium ribbon into small (about 1 cm) pieces and carefully drop it into the flask.

4. Allow the magnesium to react completely with the acid, and then make up to 250cm3 of solution with distilled water.

5. Pipette out 25.0cm3 of this solution into a conical flask. Place this on a white tile.

6. Titrate the excess acid in the solution with standard 0.100M sodium hydroxide, using a few drops of screened methyl orange indicator.

7. Observe and record all results

Results

Weighing

Mass of magnesium ribbon (rough) = 0.32g

Mass of magnesium ribbon (accurate) = 3.154g

Titration

Burette - standard (0.1<) sodium hydroxide solution

Pipette - unknown hydrochloric acid solution

Burette

Rough

1

2

3

Final vol. (cm3)

22.1

46.3

24.3

48.5

Initial vol. (cm3)

0

22.1

0

24.3

Titre (cm3)

22.1

24.2

24.3

24.2

Average tire (cm3)

24.2

Hence 24.2cm3 of standard 0.1M sodium hydroxide solution reacted completely with 25.cm3 of unknown hydrochloric acid solution.

Analysis

The reactions that were undergone during the reaction were between magnesium with the hydrochloric acid, and the sodium hydroxide with the magnesium chloride solution. The equations (both word and symbol) are shown below:

Magnesium + Hydrochloric acid --> Magnesium Chloride + Hydrogen

Mg + 2HCl --> MgCl2 + H2

From our results, we are now able to calculate the number of moles of excess acid in 250cm3 of the solution. To do this, we need to be able to calculate the number of moles in sodium hydroxide, using the following equation:

Number of moles = volume (in dm3) x molarity

Substituting the values gives us:

Number of moles = (25 / 1000) x 0.1

= 2.5 x 10-3 moles

Therefore, for calculating the number of moles in excess hydrochloric acid, we multiply our answer (number of moles in NaOH) by 10. This gives us 0.025 moles.