We're being asked to determine which among the statements are incorrect.

Recall that the **rate law** only focuses on the reactant concentrations and has a general form of:

$\overline{){\mathbf{rate}}{\mathbf{}}{\mathbf{law}}{\mathbf{=}}{\mathbf{k}}{\left[\mathbf{A}\right]}^{{\mathbf{x}}}{\left[\mathbf{B}\right]}^{{\mathbf{y}}}}$

k = rate constant

A & B = reactants

x & y = reactant orders

**Given reaction: aA + bB + C → dD + eE**

**Rate Law: Rate = k[A]**^{q}**[B] ^{r}[C]**

**Analyzing each given statement:**

Consider the reaction

*a*A + *b*B + C → *d*D + *e*E

The rate law is known to be

Rate = *k*[A]* ^{q}*[B]

Which of the following statements is __incorrect__?

A) The exponents *q* and *r* are always equal to the coefficients *a* and *b*, respectively.

B) The overall reaction order is *q* + *r* + *s*.

C) The exponent s must be determined experimentally.

D) The symbol *k* represents the rate constant.

E) The exponents *q*, *r*, and s are sometimes non-integers.

Frequently Asked Questions

What scientific concept do you need to know in order to solve this problem?

Our tutors have indicated that to solve this problem you will need to apply the Rate Law concept. You can view video lessons to learn Rate Law. Or if you need more Rate Law practice, you can also practice Rate Law practice problems.

What professor is this problem relevant for?

Based on our data, we think this problem is relevant for Professor Czernuszewicz's class at UH.