Deterioration of buildings, bridges, and other structures through the rusting of iron costs millions of dollars every day. although the actual process also requires water, a simplified equation (with rust shown as fe2o3) is: 4 fe (s) + 3 o2 (g) → 2 fe2o3 (s) δhrxn = - 1.65 * 103 kj (a) what is the δhrxn when 0.250 kg of iron rusts?

Question

Travis Barron

Chemistry

31 August, 03:46

Deterioration of buildings, bridges, and other structures through the rusting of iron costs millions of dollars every day. although the actual process also requires water, a simplified equation (with rust shown as fe2o3) is: 4 fe (s) + 3 o2 (g) → 2 fe2o3 (s) δhrxn = - 1.65 * 103 kj (a) what is the δhrxn when 0.250 kg of iron rusts?

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Rebekah · Answer 1

From the balanced chemical reaction, we can see that for every 4 moles of Fe (iron) formed, the heat released or heat of reaction is 1.65 x 10^3 kJ.

So calculate the moles of iron first:

moles iron = 250 g / (55.85 g/mol) = 4.48 mol

So the heat of reaction is:

δhrxn = (-1.65 x 10^3 kJ / 4 mol) * (4.48 mol) = - 1.85 kJ