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.