Which of the following are true statements for the above chemical reaction ?

Question

2H2(g) + O2(g) --> 2H2O(g) ΔH = −486.6kJ

a) A phase change is involved


b) Two moles of hydrogen gas react with one mole of oxygen gas to
produce two moles of liquid water


c) Energy is absorbed by the system, the reaction species


d) If the reaction takes place at constant pressure, q = −483.6 kJ


e) ΔH for reacting four moles of hydrogen gas with two moles of
of oxygen gas is −483.6 kJ

Answer 1

"B"

The mole ratio in the balanced equation is 2 H2 to 1 O2 to 2 H2O

A - no phase change is involved since they are all in the gas (g) phase

C - energy is not absorbed, it is given off since the delta H is negative.

D - nothing is mentioned about pressre in the problem

E - if twice as much of each were to react, you would get twice as much heat given off

Answer 2

E is a wrong answer, you would need to multiple H by 2 to get the correct value. D is the correct answer, because A is a false statement (everything is still a gas), B is wrong because it says liquid water, C is wrong because energy is given off, not absorbed

Answer 3

a) fake: Your reactants and products are the two gases b) fake: in user-friendly terms one mole of Oxygen gas is bypass with on the grounds that oxygen gas is diatomic c)real d)fake: must be - 486.6 kj e) fake: must be -243.3 kj

Answer 4

try e (pretty obvious if you think about it)