Question 2 -
Liquid vitamin C, (ascorbic acid) C6H8O6 , readily reacts with atmospheric oxygen, O2, to form liquid dehydroascorbic acid, C6H6O6 and water.
a) Explain why this is a chemical reaction.
b) Write a balanced equation for this reaction.
c) How many mole of oxygen gas is required to react with 1000.0 mg of vitamin C?
d) Would 1.00 g of oxygen be enough to completely react with the 1000.0 mg of vitamin C? Explain your answer by determining the limiting reagent.
e) What is the theoretical yield of dehydroascorbic acid for the reaction of 1000.0 mg of vitamin C with 1.00 g of oxygen?
Question 3 (based on Topic 6 - chapter 7 of text)
a) Solarcaine aerosol spray is used for temporarily relieving pain as it works as a topical anesthetic numbing the skin. Solarcaine is in fact benzocaine, C9H11NO2 with molar mass 165.2 g/mol.
If a 100.0 mL spray can contains 1.00 g of benzocaine, calculate the pressure within the can and explain why it is likely to explode when heated in a fire where the temperature may be as high as 1300oC ?
b) What characteristic must a compound have to experience the following intermolecular forces?
i. London Dispersion Forces
ii. Dipole-dipole forces
iii. Hydrogen bonding
c) Identify the predominant intermolecular forces in dimethyl ether (CH3OCH3) and ethanol (C2H5OH) and predict which will have the higher boiling point.
http://www.jiskha.com/display.cgi?id=1430350750
Im not sure how to do
Using the coefficients in the balanced equation, convert mols vit C to mols to mols O2.
that part though
Given the number of similar questions the last few days I feel there is an assignment coming due soon for these students.
Perhaps the answer can be found in your lecture notes and some internet research rather than cheating by copy-pasting the entire question set onto a forum.
ive spent 3 days looking for help in text books and google and past lectures and still cant work it out... next thing is to ask people who know what they're doing
Anyone got a good explanation for this question.
Question 2:
a) This is a chemical reaction because there is a change in the chemical composition of the substances involved. The vitamin C (ascorbic acid) reacts with atmospheric oxygen to form dehydroascorbic acid and water, resulting in the formation of new chemical compounds.
b) The balanced equation for this reaction is:
C6H8O6 + O2 -> C6H6O6 + H2O
c) To determine the moles of oxygen gas required, we need to convert 1000.0 mg of vitamin C to moles.
First, we calculate the molar mass of vitamin C:
6 * (12.01 g/mol) + 8 * (1.008 g/mol) + 6 * (16.00 g/mol) = 176.12 g/mol
Now, we can calculate the moles of vitamin C:
1000.0 mg * (1 g / 1000 mg) * (1 mol / 176.12 g) = 0.00567 mol
Since the balanced equation shows a stoichiometric ratio of 1:1 between vitamin C and oxygen gas, we can conclude that 0.00567 mol of oxygen gas is required.
d) To determine if 1.00 g of oxygen is enough, we need to compare the moles of vitamin C available to the moles of oxygen gas required.
The moles of oxygen gas can be calculated by dividing the mass of oxygen by its molar mass:
1.00 g * (1 mol / 32.00 g) = 0.03125 mol
Comparing the moles of vitamin C (0.00567 mol) to the moles of oxygen gas (0.03125 mol), we can see that vitamin C is the limiting reagent. Therefore, there is not enough oxygen to completely react with the 1000.0 mg of vitamin C.
e) The theoretical yield can be calculated based on the stoichiometry of the balanced equation. Since the balanced equation shows a 1:1 ratio between vitamin C and dehydroascorbic acid, the theoretical yield of dehydroascorbic acid will be equal to the moles of vitamin C reacted.
The moles of vitamin C is calculated as mentioned before: 0.00567 mol.
Therefore, the theoretical yield of dehydroascorbic acid is also 0.00567 mol.
Question 3:
a) To calculate the pressure within the can, we can use the ideal gas law equation:
PV = nRT
We are given the volume of the can (100.0 mL), the mass of benzocaine (1.00 g), and the molar mass of benzocaine (165.2 g/mol). We need to convert mL to L and g to moles before plugging into the equation.
First, convert the volume to liters:
100.0 mL * (1 L / 1000 mL) = 0.100 L
Next, calculate the moles of benzocaine:
1.00 g * (1 mol / 165.2 g) = 0.00606 mol
Now, we can calculate the pressure using the ideal gas law equation. Assuming room temperature and atmospheric pressure, we can use the values:
R = 0.0821 L·atm/(mol·K)
T = 298 K
P * 0.100 L = 0.00606 mol * 0.0821 L·atm/(mol·K) * 298 K
P = 1.47 atm
The pressure within the can is approximately 1.47 atm.
The can is likely to explode when heated in a fire because the high temperature of 1300°C will cause the gas inside the can to expand rapidly. Since the can is a closed container, the increase in pressure due to the expanding gas will exceed the can's structural integrity, resulting in an explosion.
b) The characteristics required for the following intermolecular forces are:
i. London Dispersion Forces: These forces occur between all molecules, regardless of polarity. They are caused by temporary fluctuations in the electron distribution around atoms, resulting in temporary dipoles. The strength of London dispersion forces depends on the size and shape of the molecules. Larger and more symmetrical molecules tend to have stronger London dispersion forces.
ii. Dipole-dipole forces: These forces occur between polar molecules that have permanent dipoles. The positive end of one molecule attracts the negative end of another molecule, resulting in an electrostatic interaction. The strength of dipole-dipole forces increases with increasing polarity.
iii. Hydrogen bonding: This is a specific type of dipole-dipole force that occurs when hydrogen is bonded to a highly electronegative atom (such as oxygen, nitrogen, or fluorine) in a molecule, and the hydrogen atom is attracted to another electronegative atom in another molecule. Hydrogen bonding is stronger than regular dipole-dipole forces.
c) Dimethyl ether (CH3OCH3) and ethanol (C2H5OH) both exhibit London dispersion forces and dipole-dipole forces due to their polar nature.
However, ethanol has an additional intermolecular force called hydrogen bonding. The oxygen atom in ethanol is bonded to a hydrogen atom, and this hydrogen atom forms hydrogen bonds with other ethanol molecules. Dimethyl ether does not have this polar hydrogen-oxygen bond and therefore lacks hydrogen bonding.
Since hydrogen bonding is stronger than London dispersion and dipole-dipole forces, ethanol will have a higher boiling point compared to dimethyl ether due to the additional intermolecular force.