According to Le Chatelier's principle, increasing the concentration of products will shift the equilibrium position to the left to counteract the change.

In a zero-order reaction, the rate is constant, leading to a linear decrease in concentration over time.

The rate constant k for a first-order reaction is given by k = 0.693/t1/2. Thus, k = 0.693/10 min = 0.0693 min^-1.

At equilibrium, the rate of the forward reaction equals the rate of the reverse reaction.

The rate-determining step is the slowest step in the mechanism and thus determines the overall reaction rate.

Using the Arrhenius equation and the temperature dependence of the rate constant, Ea can be estimated using the formula Ea = (R * ΔT * ln(2)) / (1/T1 - 1/T2). For a 10°C increase, Ea is approximately 80 kJ/mol.

The saturation concentration is the concentration at which the reaction rate is half of its maximum value.

For a second order reaction, if [A] is doubled, the rate increases by a factor of (2^2) = 4, thus the rate quadruples.

In a zero-order reaction, the rate is constant and the concentration of the reactant decreases linearly with time.

For many reactions, an increase of 10°C can double the reaction rate, so a 20°C increase typically results in the rate doubling.

Increasing temperature raises the kinetic energy of molecules, leading to more frequent and effective collisions, thus increasing the reaction rate.

For a first-order reaction, the unit of the rate constant k is s^-1.