Given the initial rate data for the reaction a + b –––> c, determine the rate expression for the reaction.
[a], m [b], m δ [c]/ δ t (initial) m/s 0.215 0.150 5.81 x 10–4 0.215 0.300 1.16 x 10–3 0.430 0.150 2.32 x 10–3
a) (δ[c]/δt) = 1.80 x 10–2 m –1 s –1 [a][b]
b) (δ[c]/δt) = 3.60 x 10–2 m –1 s –1 [a][b]
c) (δ[c]/δt) = 1.20 x 10–1 m –2 s –1 [a][b]2
d) (δ[c]/δt) = 5.57 x 10–2 m –3 s –1 [a]2 [b]2
e) (δ[c]/δt) = 8.37 x 10–2 m –2 s –1 [a]2 [b]

B) Δ[C]/Δt = 3,60x10⁻² M⁻¹s⁻¹ [A] [B]

Explanation:

For the reaction A + B → C

The formula for rate of reaction is:

Δ[C]/Δt = k [A] [B]

As you have [A], [B] and Δ[C]/Δt information you can multiply [A] times [B] and take this value as X and Δ[C]/Δt as Y. The slope of this lineal regression will be k.

Thus, you must obtain:

y = 3,60x10⁻² X

Thus, rate of reaction is:

B) Δ[C]/Δt = 3,60x10⁻² M⁻¹s⁻¹ [A] [B]

I hope it helps!

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