`ArarrB` , Rate of this reaction is expressed by `-(d[A])/(dt) " or "+(d[B])/(dt)` . What is the significance of (-) sign and (+) sign in this case ?

The rate of the reaction , ArarrB , is expressed by -(d[A])/(dt) . What does the (-) sign imply ?

In the reaction ArarrB , if the change in concentration of A in the time interval Deltat " is " Delta[A] , then the rate of reaction is -(Delta[A])/(Deltat) . What is the significance of the 'negative ' sign in the expression -(Delta[A])/(Deltat) ?

For a chemical reaction xP rarr yQ In [-(d[P])/(dt)] = In[(d[Q])/(dt)]+6.909 The ratio of x to y is approximately 10^n ,n = ?

For xA+yB rArr zC, - (d[A])/(dt) = (d[B])/(dt) = 1.5 (d[C])/(dt) then x,y and z are respectively.

What are the signs of DeltaG and E^@ cell for spontaneous reaction?

In the reaction 3A to 2B , rate of reaction +('d(B)')/(dt) is equal to

For the first order reaction Ararr2B , rate can be expressed any one of the following rate equations. (1)-(d[A])/(dt)=k_(1) [A] (2) (d[B])/(dt)=k_(2) [A] Find the relation between k_(1) and k_(2) .

Rate law for the reaction A + B to product is rate= k[A]^2{B] . What is the rate constant, if rate of reaction at a given temperature is 0.22 Ms^-1, when [A] =1M and (B]= 0.25 M?

aArarr bB is a first order reaction with respect to A. The rate of this reaction can be determined from any of the following equations : -(d[A])/(dt)=k_(1)[A] and (d[B])/(dt)=k_(2)[A] . Deduce the relation between k_(1) and k_(2) .