Kinetics and Energy - AP Chemistry

The integrated rate law shown is for a reaction rate law with a first-order dependence on reactact A. If such a rate law accurately describes the kinetics of the reaction, then ln\[A\] will vary linearly with respect to time.

Since reactant molecules collide and interact to break old bond and form new ones, any factors affecting collision and interactions will affect the reaction rate. Thus, increasing temperature will increase both the frequency of collisions and the average kinetic energy of the molecules. Enzymes bring molecules close to each other and orient them in a way that facilitates reactions.

If the activation energy of the forward reaction is greater than that of the reverse reaction, the products must have a higher enthalpy than the reactants. The net enthalpy change is therefore positive, meaning that it is endothermic.

If the activation energy of a forward reaction is greater than that of the reverse reaction, thenthe products must have a higher enthalpy than the reactants (draw a potential energy diagram to visualize).

The rate-determining step in a reaction mechanism is a kinetic bottleneck, in that it prevents the overall reaction from proceeding; thus, it is what determines how quickly the overall reaction can proceed.

Activation energy is the energy barrier that needs to be overcome for a reaction to proceed; the higher the activation energy, the slower the reaction. Activation energy can only be altered via a catalyst. Catalysts are chemical substances that lower the activation energy, allowing reactions to proceed faster. Other physical quantities such as temperature and pressure don’t alter the activation energy.

Reaction rate is a direct measure of the speed of a reaction. As mentioned, increasing activation energy will increase the barrier and, therefore, slow down the reaction.

Catalysts are chemical substances that decrease the activation energy, thereby increasing the reaction rate. They are commonly used in chemical reactions to drastically speed up reactions that might otherwise take hours or days to complete. Enzymes are biological catalysts that facilitate most of the biological reactions happening in our bodies.

This question is asking about the relationship between activation energy of a reaction (kinetics) and the amount of products produced (equilibrium). Remember that altering the speed of a reaction (kinetics) does not change the equilibrium of the reaction. Increasing or decreasing activation energy (which alters the speed of reaction) will simply allow for the reaction to proceed slower or faster, respectively. It will not change the amount of products produced at the end.

Activation energy is the difference between the energy of the transition state and the energy of the reactant. Recall that activation energy is the energy barrier that needs to be overcome by a reaction. The transition state is a higher-energy, intermediary molecule that lies in between the reactants and products. It is created when reactants are in the process of becoming products. Therefore, to get to products, it is necessary to go through this high-energy transition state.

Activation energy is the energy needed to climb this energy “hill” (energy needed to go from reactants to transition state); therefore, the activation energy is the energy of the transition state (top of the energy hill/barrier) minus the energy of the reactant.

Catalysts are substances that increase reaction rates without being consumed in the reaction. They decrease the activation energy needed, and they do not always need to be in the same phase as the reactants. In heterogeneous catalysis, the catalyst is in a different phase than the reactants. Equilibrium concentrations of both reactants and products are unchanged by the addition of a catalyst.

All of the choices are true, except that catalysts can be in distinct phases than the reactants. These are known as heterogenous catalysts.

A catalyst is a substance that increases the rate of a reaction, typically by lowering the activation energy required to initiate a reaction. The catalyst does not affect the equilibrium of a reaction, and is not consumed during the reaction.

A catalyst has no effect on the relative stability of the reactants or products, nor does it effect the temperature of a reaction.

Instead, catalysts lower the energy of transition states, increasing their stability, to lower the overall activation energy of the reaction. When the reaction requires less energy, it proceeds at a faster rate.

A catalyst will not be consumed during a reaction, so the catalyst will be whichever chemical is found both on the reactant side of the equation and on the product side.

For equation 1 that is compound Z; for equation 2 that is compound A.

Catalysts increase the reaction rate without being consumed during the reaction. They don't cause the reaction to make more product, but since the catalyst won't be used up in the reaction.

Enzymes function by forming complexes with their substrates at active sites. This interaction is often thought of as a lock and key mechanism, in the sense that the active site is shaped to fit a substrate.

Enzymes are biological catalysts and therefore they work to lower the energy barrier or activation energy that prevents a reaction from proceeding to equilibrium. In other words: enzymes and catalysts in general make a reaction reach equilibrium faster.

Therefore enzymes do not change the equilibrium product concentration, just the time it takes to get to equilibrium.

Note that when catalyst decreases the activation energy (Ea), will not be affected. The step in a reaction with the largest activation energy usually is the slow step, which catalysts facilitate. Catalysts do not affect the thermodynamic quantities . Since catalysts are not consumed in the reaction, they do not appear in the net equation of the reaction.

For a first order reaction, the ln \[A\]t is linear with t.