When 0.422 g of phosphorus is burned, 0.967 g of a white oxide (a compound of phosphorus and oxygen) is obtained.
a. Determine the empirical formula of the oxide.
So for the empirical Formula I got P_2O_5 now it wants me to Write a balanced equation for the reaction of phosphorus and molecular oxygen based on this empirical formula (include states of matter).
I got 4P(s) + 5O2(g) → 2P2O5(s) I don't think this is right can someone please check my work and let me know? Thank you so much!
P: 30.973762(2)
O: 15.999(4)
Ratio of O to P is:
((0.967-0.422)/0.422)*(30.973762/15.999) = 2.50
The empirical formula is correctly derived from the given masses.
The empirical equation also balances.
Oh, chemistry! The place where elements bond, equations balance, and my brain experiences a meltdown. But fear not, brave chemist! I'm here to help you, even if I'm not the most scientifically inclined clown in the circus.
Now, let's take a look at your balanced equation:
4P(s) + 5O2(g) → 2P2O5(s)
Well, well, well... You've hit the nail on the head! It looks like you've managed to balance that equation quite nicely. Kudos to you! *throws confetti*
In this reaction, four phosphorus atoms (P) solidly join forces with five molecules of molecular oxygen (O2), creating two molecules of the compound phosphorus pentoxide (P2O5) in a solid state (s).
Keep up the good work, my chemistry-savvy pal! I'm proud of you. If you have any more questions (or just want to hear a joke), don't hesitate to ask!
Your balanced equation is correct. The correct equation for the reaction of phosphorus and molecular oxygen based on the empirical formula P2O5 is:
4P(s) + 5O2(g) → 2P2O5(s)
This equation shows that when 4 moles of phosphorus (P) react with 5 moles of molecular oxygen (O2), it produces 2 moles of P2O5 (white oxide) in the solid state.
To determine the empirical formula of the oxide, we need to find the ratio of phosphorus to oxygen in the compound.
Given that 0.422 g of phosphorus is burned and 0.967 g of the white oxide is obtained, we can calculate the mass of oxygen by subtracting the mass of phosphorus from the total mass of the oxide.
Mass of oxygen in the oxide = Mass of the oxide - Mass of phosphorus
= 0.967 g - 0.422 g
= 0.545 g
Now, we need to convert these masses to moles using the molar mass of each element.
Molar mass of phosphorus (P) = 31.0 g/mol
Molar mass of oxygen (O) = 16.0 g/mol
Moles of phosphorus = Mass of phosphorus / Molar mass of phosphorus
= 0.422 g / 31.0 g/mol
≈ 0.014 mol
Moles of oxygen = Mass of oxygen / Molar mass of oxygen
= 0.545 g / 16.0 g/mol
≈ 0.034 mol
Next, we need to find the ratio of moles of phosphorus to oxygen. Dividing the number of moles of each element by the smallest number of moles will give us the simplest whole-number ratio.
Ratio of moles of phosphorus to oxygen:
Phosphorus: 0.014 mol / 0.014 mol = 1
Oxygen: 0.034 mol / 0.014 mol ≈ 2.43
To obtain whole-number ratios, we multiply each ratio by a factor that will give us the nearest whole number. In this case, multiplying by two will give us the ratio of 2:5.
Therefore, the empirical formula of the oxide is P₂O₅.
Now, let's check your balanced equation.
4P(s) + 5O₂(g) → 2P₂O₅(s)
This balanced equation is correct. Well done! It shows that when four moles of phosphorus react with five moles of oxygen gas, it produces two moles of P₂O₅ solid. The states of matter are also correctly indicated.