The rate of a chemical reaction is highly sensitive to changes in temperature. This temperature dependence is mathematically explained using the Arrhenius equation; which expresses the relationship between the rate constant, the absolute temperature, the frequency factor, and the activation energy. The activation energy and frequency factor can also be determined graphically by converting the Arrhenius equation into a non-exponential form. Using the natural logarithms on both sides, an equation for a linear function is generated. The slope value corresponds to the negative value of activation energy over gas constant, and the y-intercept corresponds to the natural log of the frequency factor. This equation can be used to generate a graph called the Arrhenius plot, in which the natural log of the rate constant is denoted as a function of the inverse of temperature in kelvin. Kinetic data of experiments and reactions can be illustrated and analyzed using this Arrhenius plot. In this example, the graph yields a straight line. The slope value given in kelvin is set equal to the negative value of activation energy over R. After assigning the value for the gas constant and solving for the activation energy, a value of 93.1 kJ/mol is obtained. Besides, the y-intercept of 26.8 is equal to the natural log of the frequency factor. Thus, solving for A gives the value of 4.36 × 1011 with the unit one-over-molarity-seconds — the same unit as the rate constant. In cases of limited kinetic data or difficulties with graphical representation, a two-point form of the Arrhenius equation can be utilized to calculate the activation energy in a non-graphical manner. In such cases, the non-exponential form of the Arrhenius equation is modified to include rate constants at two different temperatures. Subsequent subtraction and rearrangement of the expression yield the two-point form of the Arrhenius equation, which is used to calculate the activation energy from experimentally-generated rate constants at two different temperatures. By substituting the values, the activation energy for this reaction is calculated to be 145 kJ/mol.