In chemistry, compounds can be distinguished by using the empirical formula. The formula provides the simplest positive integer ratio of elements in a compound. The empirical formula is largely useful in determining the ratio of elements within ionic compounds where the structure is of a non-directional nature of bonding where any ion at any time could be surrounded by 4, 6, or 8 oppositely charged ions. This creates a pattern of endlessly repeating lattice of ions they do not exist as a free unit of atoms but in crystal lattices with repeating ions in specific ratios which is why empirical formula is used as a form of identification for defining an ionic-bonded substance.
Calculating the empirical formula:
To calculate the empirical formula for when two reactants undergoes a chemical reaction, the following 5 steps should be taken: 1. Record the masses of all the elements present in a given compound. 2. Convert the masses into moles (dividing by atomic weight in grams). 3. Then divide through by the smallest number of moles to get a ratio.
4. It the numbers are not whole numbers, multiply by a suitable small factor to get a whole number. 5. Finally, round off the numbers in the previous step if applicable to get the prime numbers which indicates the empirical formula. To demonstrate with a simple example; The molecular formula of butane is C4H10, however as the ratio of carbon atoms to hydrogen atoms is 4:10; it can be reduced to the ratio of 2:5. We can see that it is the simplest ratio while remaining as an integer.
Butane can now be represented in empirical formula as C2H5.
Regarding the Mole and its formula:
The mole is the quantity of a substance which contains as many elementary units (atoms, ions, molecules) as there are atoms in exactly 12 grams of carbon-12 isotope. A mole of an element is the mass in grams that is numerically equal to the atomic weight. Also, a mole of a compound is the mass in grams that is numerically equal to the molecular weight. In simpler words;
A mole of a substance is simply the atomic / molecular weight in grams. eg; A mole of copper (atomic weight 63.6) is 63.6grams.
Therefore in a diagram;
The number of atoms or molecules in a mole of any substance is the Avogadro Constant which is 6.02 x . The molar mass is taken to be the relative atomic mass of an element which is the average mass of atoms present in any naturally occurring element relative to the mass of one atom of carbon-12 isotope taken as exactly 12 which gives formula weight (sum of the atomic weights of the atomic species as given in the stated formula for the compound.) The quantitative stoichiometric relationships governing mass and amount is used in the following experiment regarding the combustion reaction of magnesium metal. Magnesium is reacted with oxygen from air in a contained crucible, and the masses before and after the oxidation is measured. The resulting masses are used to calculate the experimental empirical formula of magnesium oxide, which is then compared to the theoretical empirical formula. A crucible and Bunsen burner will be used to heat magnesium metal for burning.
The purpose of this experiment is to perform an experiment of the combustion of Magnesium and gather precise and accurate data of masses and thus find the number of moles of the substance through the stoichiometry mole equation in order to evaluate the empirical formula of Magnesium Oxide.
1 The Bunsen burner was set up with the tripod. The pipe clay triangle was placed over the tripod, ensuring that it is secure. 2. The crucible containing the magnesium was positioned in the pipe clay triangle securely with the lid on. 3. The gas was turned on and the Bunsen burner ignited to a blue flame. 4. The crucible was fired strongly for 5minutes until the bottom of the crucible glowed red over the blue flame to rid of contaminants. 5. The flame was removed and to cool the crucible with lid. 6. A piece of magnesium about 5 cm long was cut. 7. The surface of the magnesium ribbon was thoroughly cleaned with steel wool and its appearance was recorded 8. The cooled crucible and lid was weighed (1st mass to 2dp) 9. The cleaned magnesium was coiled to fit inside the same crucible and covered with the same lid. 10. The crucible containing the magnesium with the lid on was weighed. (2nd mass to 2dp) 11. The crucible containing the magnesium was positioned without the lid onto the pipe triangle setup, ensuring its security. 12. The gas was turned on again and the Bunsen burner was ignited to a red flame (air hole fully open). 13. As the magnesium began to glow, the crucible was covered with its lid carefully with tongs. 14. Heat strongly for about 10 minutes lifting the lid a little VERY carefully occasionally to admit oxygen. 15. Keep heating and lifting the lid until ALL the magnesium turns into gray-white powder or until no further reaction can be witnessed to ensure complete reaction (for around 5 minutes) 16. Turn off the gas and allow the apparatus to cool. 17. Weigh the completely cooled crucible containing magnesium oxide with the lid carefully. (3rd mass to 2dp)
Wear safety glasses. It is important to have eye protection during the combustion of Magnesium as the burning Magnesium in the crucible produces a very bright light which emits a harmful intensity of UV light which can cause eye discomfort or damage. Do not inhale the smoke produced when Magnesium is burned. Magnesium Oxide smoke can cause irritation in the nose, eyes and lungs and in large amounts, may cause metal fume fever. Use tongs at all times when handling hot objects. Careful handling of hot equipment such as the crucible and its lid during the heating is important as the very high temperature can burn skin due to improper or insecure handling. Do not cool the crucible or lid under cold water immediately after heating. This can cause the equipment to crack and the shards may easily pierce the skin. If the crucible crack during the experiment, it is vital that the person discontinue any progress with the experiment and proceed to clean the broken equipment away immediately and place into the broken glass bin.
Mass of Magnesium
Mass of oxygen
Mass of Magnesium Oxide
Mass of crucible + lid
Total mas of Mg oxide in crucible + lid:
Total mass of crucible + lid + magnesium:
Percent composition of Magnesium in compound:
Mass of Mg in 1 mole/ Mass of MgO in 1 mole % composition of Oxygen in compound:
Upon observation, the 5cm Magnesium ribbon had a slightly greasy texture and a brittle and coarse surface. It had a hazy, dark metallic sheen to its appearance. After polishing its surface front and back thoroughly with steel wool, there was a change in its appearance. It had a shiny and glossy silver lustre with a smooth clean surface, no longer feeling waxy. The steel wool’s purpose was to remove the oxide layer of carbonate and sulphate which has coated the Magnesium ribbon’s surface due to its slow oxidation in air with CO2 and SO2 and other possible contaminants (which may have caused the strip to feel greasy). The oxide coating on the Magnesium would have made the ribbon more resilient to ignite immediately and thus hinder the combustion of the metal and prolong the time it takes for the metal to fully combust. The procedure of rubbing Magnesium’s surface with steel wool was beneficial in order to expose fresh Magnesium to facilitate the contact of the ribbon with the flame quicker and thus a faster complete combustion.
The crucible and lid used had minimal surface stains on the outside however it was heavily contaminated with residual substances towards the inside base. By firing the equipment under a blue Bunsen flame thoroughly, it became apparent that any moisture or volatile materials present are burnt off by 5minutes to reveal a clean crucible free of stains or moisture. The purpose of firing the crucible at a high temperature was to quickly eradicate any stubborn chemicals which may have resisted cleaning by water, as well as any moisture the crucible may hold to provide a clean and dry equipment which can ensure accuracy and validity of calculations of masses.
In order to activate the reaction of Magnesium, a source of energy was needed. The flame provided a source of heat which prompted a chemical reaction to proceed. When the magnesium was supplied with energy in the crucible, it reacted with a limited quantity of oxygen by using the lid to prevent high exothermic energy (Magnesium would react vigorously if heated in the presence of unobstructed air flow) and the escape of any magnesium oxide during the combustion. It became oxidized to become an ionic compound Magnesium Oxide.
After the experiment of combusting Magnesium, the residue in the crucible is observed to be in a fine white powder form of Magnesium Oxide, an ionic compound. The exothermic reaction of combusting Magnesium produced a very bright light within the crucible due to the rapid heating of the Magnesium, where it quickly absorbs energy through ionisation. Magnesium ionises to become a cation while Oxygen ionises to an anion, forming an ionic bond due to electrostatic forces.
In this experiment, through the combustion reaction of Magnesium, a word equation forms between Magnesium, Oxygen and the ionic compound Magnesium Oxide. Magnesium + Oxygen Magnesium Oxide
When ignited, magnesium has reacted with oxygen to form the products Magnesium Oxide. By taking the mass of equipments’ used and its mass during and after the reaction, the mass of Magnesium, Oxygen and Magnesium Oxide can be calculated. The mass of the reactants should very closely or mirror the mass of products by Law of Conservation of Mass.
(mass 1) Crucible + lid = 24.31g (weight after ridding of visible contaminants on crucible) (mass 2) Crucible + lid + magnesium = 24.39g (weight after polished magnesium is placed in fired crucible + lid) (mass 3) Crucible + lid + product = 24.44g (weight of reacted substance MgO in the crucible with lid) To calculate the mass of Magnesium metal = Mass 2 – Mass1 = 24.39g – 24.31g Mass of Magnesium = 0.08g
To calculate the Mass of Oxygen incorporated = Mass 3 – Mass 2 (the increase in mass corresponds to the mass of oxygen) = 24.44g – 24.39g Mass of Oxygen = 0.05g
To calculate the mass of oxide product formed = Mass 3 – Mass 1 (to validate through law of conservation of mass) = 24.44g – 24.31 Mass of Magnesium Oxide = 0.13g
Now that the mass of each substance which took part in the reaction is found, the number of moles can be found through the relationship below. Number of Moles = Mass / Molar Mass
The number of Moles can be calculated by knowing the mass of individual substances in the experiment divided by the molar mass (given on the periodic table as atomic mass number) to experimentally determine empirical formula for the ionic oxide.
CONVERTING TO MOLES:
Number of Moles of Magnesium: N = 0.08g / 24.31g/mol
= 0.00329082692 moles Number of Moles of Oxygen: N = 0.05g/16x2g/mol (oxygen is doubled as it exists as a diatomic molecule) = 0.0015625 moles
DIVIDE BY SMALLEST MOLE VALUE
The number of moles of Magnesium is larger than the number of moles of Oxygen: 0. 00329(to5dp) moles > 0.00156 (to5dp) moles
By dividing each by the smallest mole value of 0.00156 of Oxygen, Magnesium and Oxygen mole ratio is calculated respectively. Magnesium Mole ratio
Oxygen Mole ratio
Multiply UNTIL WHOLE:
The next step is to multiply any decimal numbers by a small whole number and do the same for the other whole number ratio until the decimal number reaches a near whole number (which can then be rounded up). Magnesium Mole ratio is in a decimal number of 2.11, and as it is very close to 2(with 0.11 extra of being a whole number), the number 2.11 is round down to 2 so the process of multiplying until whole is omitted. We obtain the mole ratio as follows: Magnesium : Oxygen 2 : 1 With 2 moles of Magnesium reacting with 1 mole of Oxygen, this should suggest 2 moles of MgO after balancing the equation. 2Mg(s) + O2(g) → 2MgO(s) To confirm, the law of conservation of mass is applied which states that in a balanced equation, matter is conserved: Total mass of reactants = Total mass of products
Thus, by adding the mass on the left hand side (reactants), it should equal the right hand side (product). Magnesium + Oxygen = 0.08g + 0.05g Magnesium Oxide = 0.13g It is found that there is no difference in mass between the left side of 0.13g of Magnesium + Oxygen to the right side of 0.13g of Magnesium Oxide in the equation considering the possibility of experimental errors, which represents a positive outcome as an theoretical equation is established and proved to be true. The empirical formula for Magnesium oxide is MgO, which is the correct formula and thus the aim of this experiment has been met.
The experiment demonstrated the ability for a substance to exist in the empirical formula composition as the simplest ratio of elements present in the compound. It also demonstrated quantitative stoichiometric relationships between the number of moles, mass and molar mass in a chemical reaction. The theoretical result for the combustion product of Magnesium and Oxide is MgO, which in comparison to the experiment result of MgO was proven to be the same. This experiment had demonstrated the basic chemical reaction and the change in states between elements in order to form a stable ionic compound. Through the ionic bond between a metal and a non-metal, Magnesium Oxide was the product of two reactive elements Magnesium and Oxygen.
Theoretical laws of conservation of mass and constant composition, lead to the formation of a relationship between the reactants (Magnesium and Oxygen) and product (Magnesium Oxide). The total mass of the products of a reaction must equal the total mass of the reactants. (0.8+0.5 = 0.13) The coefficient of a substance indicated the amount of portions each substance existed in, based on the law of conservation of mass. (2 moles of Magnesium, 1 mole of Oxygen, and 2 moles of Magnesium Oxide).
And lastly, the empirical formula of a compound gave the lowest whole-number ratio of the atoms that is the identical with the mass ratios measured by experiment. (MgO) Any portion of a compound will have the same ratio of masses as the elements in the compound. Metal and a non-metal > reaction of 2 substances taking place> Ionic bond of elements>changes in states and formation of a stable compound >the construction of an unbalanced chemical equation -> evaluation of masses of the reactants to products> the law of conservation of mass/and constant composition ->the masses of left side equals right side provide moles to balance the overall equation >establish empirical formula for end product. 3 experimental errors that may have effect on result
Magnesium Oxide forms fumes which may escape the crucible when allowing a passage for oxygen to pass when the lid is lifted. Incomplete combustion of Magnesium (as no stirring rods was used to check/sift through the oxide to prevent calculation errors) Unthorough firing of crucible and lid (due to contact between tongs and crucible, certain areas may be missed) 3 improvements to method to improve results:
Monitor the reaction of Magnesium with oxygen carefully, and keep the lid in place on the crucible containing the magnesium instead of off in step 11. Heat the magnesium for five minutes longer, lightly rotating the crucible at its base to ensure complete combustion. Fire the crucible and the lid twice to ensure that its’ completely dry and clean.
An experiment was performed to calculate the empirical formula of magnesium oxide by comparing the masses of solid magnesium metal to the magnesium oxide solid product in a crucible. The concept of stoichiometry mole equation leads to finding the empirical formula of magnesium oxide. One major finding during the experiment was that burning magnesium caused its mass to increase as it reacted to oxygen. The amount of mass increase is proportional and able to be calculated through weighing the mass of product Magnesium Oxide and subtracting the original Magnesium mass to calculate the mass of Oxygen which partook in the combustion.
The Empirical formula had indicated the proportion of Magnesium to oxygen (1:1) and identifies the compound to be Magnesium Oxide. The balance was very accurate in taking precise measurements of masses and the supervision of the experiment during the burning had been careful to prevent any loss of Magnesium Oxide mass to escape which in turn caused my mass results to apply to the law of conservation of mass without any experimental errors.
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