Cherryl's Chem Test #3

The given answer correctly calculates the concentration of HI at equilibrium using the equation provided. It substitutes the values into the equation [HI] = 0.275 - 2(0.02748) / (0.275)^2, which simplifies to [HI] = 0.22 M. This value represents the concentration of HI at equilibrium in the reaction vessel.

The equilibrium constant, Kc, is calculated by dividing the concentration of N2O4 by the square of the concentration of NO2. In this case, the concentration of N2O4 at equilibrium is given as 0.0750 M. Therefore, to find Kc, we divide 0.0750 by the square of the concentration of NO2. The concentration of NO2 is not explicitly given in the question, but it can be determined by using the stoichiometry of the balanced equation. Since 2 moles of NO2 are produced for every mole of N2O4 that decomposes, the initial moles of NO2 can be calculated as (0.625 mol of N2O4) x (2 mol of NO2 / 1 mol of N2O4) = 1.25 mol of NO2. The initial concentration of NO2 can then be calculated as (1.25 mol of NO2) / (5.00 L of vessel) = 0.25 M. Taking the square of this concentration and dividing it into 0.0750 gives us a Kc value of 7.5.

The rate constant for the first order reaction A--B can be determined using the equation K = ln([A]t/[A]0) / t, where [A]t is the concentration of A at time t, [A]0 is the initial concentration of A, and t is the time elapsed. Using the given data, we can calculate the rate constant as follows: K = ln(0.88/1.76) / 6 = ln(0.5) / 6 = -0.693 / 6 = -0.116 s^-1. Therefore, the rate constant for the first order reaction A--B is -0.116 s^-1.

The given graphs show a constant rate of reaction, regardless of the changes in concentration. This indicates that the reaction rate is independent of the concentration of the reactants. Therefore, this reaction is zero-order, meaning that the rate of reaction is constant and not influenced by the concentration of the reactants.

The rate of the reaction is directly proportional to the square of the concentration of A and the square of the concentration of B. This can be determined by comparing the initial rates of the reaction in the different experiments. In Experiment 1, doubling the concentration of A and B results in an increase in the initial rate by a factor of 4. In Experiment 2, doubling the concentration of A results in an increase in the initial rate by a factor of 4, while keeping the concentration of B constant. In Experiment 3, doubling the concentration of B results in an increase in the initial rate by a factor of 4, while keeping the concentration of A constant. Therefore, the rate law that most nearly accounts for these data is k[A]2[B]2.

The rate constant of a reaction is dependent on the activation energy (Ea) and the temperature. According to the Arrhenius equation, the rate constant (k) is proportional to the exponential of (-Ea/RT), where R is the gas constant and T is the temperature in Kelvin.

In this case, we are given the rate constant (k1) at 15 degrees C and we need to find the rate constant (k2) at 37 degrees C. Since the activation energy (Ea) is constant, we can use the Arrhenius equation to find the ratio of k2/k1.

By plugging in the values of Ea, R, T1 (15 degrees C + 273 = 288 K), and T2 (37 degrees C + 273 = 310 K), we can solve for the ratio of k2/k1.

The ratio k2/k1 is equal to 0.104, which means the rate constant at 37 degrees C (k2) is 0.104 times the rate constant at 15 degrees C (k1). Therefore, the correct answer is K2=0.104.

The given reaction is shown to be second order with respect to NO2, which means that the rate of the reaction is directly proportional to the square of the concentration of NO2. This is represented by the rate=k[NO2]2. The other options, rate=k[NO2][CO]2, rate=k[CO]2, and rate=k[NO][CO2], do not accurately represent the given information about the reaction order.

The answer K2=(K1)-2 is correct because it follows the relationship between the two equilibrium constants K1 and K2. According to the given equilibria, the forward reaction of the first equilibrium is the reverse reaction of the second equilibrium. Therefore, the equilibrium constant for the second equilibrium (K2) is the reciprocal of the equilibrium constant for the first equilibrium (K1) raised to the power of -2. This can be mathematically represented as K2=(K1)-2.

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The given correct answer states that the forward reaction is exothermic. This means that the forward reaction releases heat energy into the surroundings. In an exothermic reaction, the products have lower energy than the reactants, resulting in a negative change in enthalpy (ΔH). This can be inferred from the diagram provided, where the forward reaction is shown to release energy.

The equilibrium constant expression for a reaction is determined by the stoichiometry of the reaction. In this reaction, the stoichiometry is 2H2S(g) --> H2(g) + S2(g). This means that two molecules of H2S react to form one molecule each of H2 and S2. Therefore, the concentration of H2S is squared in the equilibrium constant expression to account for this stoichiometry. The concentrations of H2 and S2 are not squared because they appear as single molecules in the balanced equation. Therefore, the correct equilibrium constant expression is Kc=[H2][S2] divided by [H2S]2.

Based on the given information, the reaction is not at equilibrium. The equilibrium constant (Kc) is a large value, indicating that the reaction favors the formation of products (SO3). However, the concentrations of reactants (SO2 and O2) are higher than the concentration of the product (SO3), suggesting that the reaction has not yet reached equilibrium. Therefore, the reaction will occur in the direction from right to left in order to reach equilibrium and increase the concentration of the product.

Increasing the temperature will change the equilibrium constant, K, because it affects the balance between the reactants and products in a chemical reaction. According to Le Chatelier's principle, when the temperature is increased, the system will shift in the direction that absorbs heat, which is the endothermic reaction. This will cause the equilibrium position to shift towards the products or reactants, depending on whether the reaction is exothermic or endothermic. Consequently, the concentrations of the reactants and products will change, altering the value of the equilibrium constant, K.

Increasing the temperature would alter the rate constant (k) for the reaction. This is because temperature is directly proportional to the rate constant according to the Arrhenius equation. As temperature increases, the kinetic energy of the reactant molecules also increases, leading to more frequent and energetic collisions. This results in a higher rate of reaction and therefore a larger value of the rate constant.

Increasing the temperature would alter the rate constant (k) for the reaction. This is because increasing the temperature increases the kinetic energy of the reactant molecules, leading to more frequent and energetic collisions. This results in an increase in the rate of the reaction, causing the rate constant (k) to increase.

The correct answer is 2. The tendency for a reaction to proceed to completion can be determined by the value of the equilibrium constant (Kp). The higher the value of Kp, the greater the tendency for the reaction to proceed to completion. In this case, the reactions are given with their respective Kp values. Comparing the Kp values, we can see that the reaction with the lowest Kp value (1.3 x 10 -5) has the least tendency to proceed to completion, while the reaction with the highest Kp value (5.9 x 10 -5) has the greatest tendency to proceed to completion. Therefore, the correct order of increasing tendency to proceed to completion is 2, 1, 3.

Removing Br2 from the reaction will shift the equilibrium position to favor more products. This is because according to Le Chatelier's principle, when a reactant is removed from a system, the equilibrium will shift in the direction that produces more of that reactant. In this case, removing Br2 will cause the reaction to produce more NO and Br2, favoring the formation of products.

Chemical equilibrium is a state in a chemical reaction where the forward and reverse reactions occur at the same rate. This means that the rate at which reactants are converted into products is equal to the rate at which products are converted back into reactants. At this point, there is no net change in the concentrations of reactants and products, indicating that equilibrium has been reached.

A catalyst speeds up a reaction by decreasing the activation energy for the reaction. Activation energy is the minimum amount of energy required for a reaction to occur. By lowering this energy barrier, a catalyst allows the reaction to proceed more readily, increasing the rate of the reaction. The catalyst itself is not consumed in the reaction and can be used repeatedly.

The correct answer is "Z must react in a step after the rate determining step." This means that the formation of Z does not directly affect the rate of the reaction. The rate determining step is the slowest step in a reaction and determines the overall rate. Since Z is formed after this step, its concentration does not affect the rate.

The correct answer is rate=1/2 delta C divided by delta temp. This is because the reaction given in the question is A+3B--2C. The rate of the reaction is determined by the change in concentration of the reactants or products with respect to time. In this case, the rate is proportional to the change in concentration of C divided by the change in temperature. Therefore, the correct expression for the rate is rate=1/2 delta C divided by delta temp.

When the concentration of B is doubled, the rate of reaction will increase by a factor of 8. This is because the rate law for the reaction is rate=k[A][B]3, which means that the rate is directly proportional to the concentration of B cubed. When the concentration of B is doubled, it will be raised to the power of 3, resulting in an increase of 2^3=8 in the rate of reaction.

The given equation represents a straight line in the form of y = mx + c, where y represents the natural logarithm of the rate constant (ln k), x represents the inverse temperature (1/T), and c represents the y-intercept. By comparing the equation with the given equation (y = -2.09 x 10^4*x + 343.3), we can see that the slope (m) is equal to -2.09 x 10^4. The activation energy (Ea) can be calculated using the formula Ea = -m * R, where R is the gas constant. Plugging in the values, we get Ea = -(-2.09 x 10^4) * 8.314 = 17.4 kJ/mol. Therefore, the activation energy is 17.4 kJ/mol.

The answer [NO]2[Cl2]=Kc[NOCl]-2 is always true because it represents the equilibrium constant expression for the given reaction. The equilibrium constant (Kc) is defined as the ratio of the concentrations of the products (raised to their stoichiometric coefficients) to the concentrations of the reactants (raised to their stoichiometric coefficients), with each concentration term raised to the power of its coefficient in the balanced equation. In this case, the coefficient of NOCl is -2, so it is raised to the power of -2 in the equilibrium constant expression. Therefore, the given answer is the correct representation of the equilibrium constant expression for the reaction.

Compressing the gas mixture into a smaller volume would cause the Br2 concentration to decrease. According to Le Chatelier's principle, when the volume of a gas is decreased, the system will try to counteract the change by shifting the equilibrium towards the side with fewer moles of gas. In this reaction, compressing the gas mixture would decrease the volume, causing the equilibrium to shift towards the side with fewer moles of gas, which is the side without Br2. As a result, the concentration of Br2 would decrease.

When the temperature increased from 10 degrees C to 20 degrees C, it caused the molecules to possess at least the minimum energy required for the reaction. This means that more molecules were able to overcome the activation energy barrier and participate in the reaction. As a result, the proportion of molecules possessing the minimum energy required for the reaction doubled.