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What is the primary cause of energy changes in chemical reactions?

Which process requires energy input in a chemical reaction?

What term describes the energy needed to break a chemical bond?

Which of the following is true about exothermic reactions?

In an endothermic reaction, how does the energy of the reactants compare to the energy of the products?

What is the symbol used to represent the heat of reaction or enthalpy change?

For an exothermic reaction, what is the sign of $\Delta H$?

Consider the reaction: $A ightarrow B + 200 , kJ \cdot mol^{-1}$. Is this reaction exothermic or endothermic?

Which of the following correctly represents an endothermic reaction in terms of energy and reactants/products?

If a reaction has a $\Delta H = -500 , kJ \cdot mol^{-1}$, how should this value be interpreted?

Consider the reaction: $2H_2O(l) + 572 , kJ ightarrow 2H_2(g) + O_2(g)$. What is the value of $\Delta H$ for this reaction?

In an energy graph for an exothermic reaction, what is the relative position of the reactants and products?

For a reaction to occur, what minimum energy is required?

What is the activated complex in a chemical reaction?

In an energy diagram, the activation energy is represented by the energy difference between:

Which statement is true regarding the activation energy of an exothermic reaction?

For the reaction $H_2(g) + F_2(g) ightarrow 2HF(g)$, which of the following energy changes occurs?

How does the activation energy of an endothermic reaction compare to the overall enthalpy change ($\Delta H$) for the same reaction?

Consider two reactions: Reaction 1 has a lower activation energy than Reaction 2. Assuming other conditions are equal, which reaction will proceed faster?

Which of the following statements accurately describes the energy profile of an endothermic reaction?

If the bond energy of reactant bonds is $X , kJ \cdot mol^{-1}$ and the bond energy of product bonds is $Y , kJ \cdot mol^{-1}$, under what condition will the reaction be exothermic?

Consider the combustion of methane: $CH_4(g) + 2O_2(g) ightarrow CO_2(g) + 2H_2O(g)$. This reaction is known to be exothermic. Which of the following is necessarily true about the bond energies involved?

Which of the following scenarios describes an endothermic process at a molecular level?

If a catalyst is added to a reaction, how does it affect the activation energy and the enthalpy change ($\Delta H$)?

For the reaction $N_2(g) + O_2(g) ightarrow 2NO(g)$, the $\Delta H$ is positive. If we were to reverse the reaction, $2NO(g) ightarrow N_2(g) + O_2(g)$, what would be the sign of $\Delta H$ for the reverse reaction?

What is the fundamental reason for energy changes occurring during a chemical reaction?

Which of the following describes the role of bond energy in a chemical reaction?

If a reaction releases heat into the surroundings, it is classified as:

In an endothermic reaction, how does the total energy absorbed during bond breaking compare to the total energy released during bond formation?

For a reaction where the enthalpy change ($\Delta H$) is positive, which statement is always true?

Consider the reaction: $XY ightarrow X + Y$, with $\Delta H = +200 , kJ \cdot mol^{-1}$. Which of the following is true?

Which of the following correctly represents an exothermic reaction in equation form?

For the reaction $2SO_2(g) + O_2(g) ightarrow 2SO_3(g)$, $\Delta H = -198 , kJ \cdot mol^{-1}$. How much energy is released when 1 mole of $SO_3$ is formed?

In an energy profile diagram of a reaction, the reactants are shown at a higher energy level than the products. This indicates a(n):

What is the minimum energy required to initiate a chemical reaction?

The activated complex is best described as:

On an energy diagram, the activation energy for a forward reaction is represented by the difference in energy between:

Which of the following is true regarding the activation energy of an endothermic reaction?

For the reaction $X_2 + Y_2 ightarrow 2XY$, which energy change must occur initially for the reaction to start?

How does the magnitude of activation energy relate to the rate of a chemical reaction?

Consider an energy profile diagram where the activated complex is only slightly higher in energy than the reactants. This suggests:

For a reaction to be exothermic, which relationship between bond energies must be true?

Which scenario describes an endothermic process at a molecular level?

If reaction 1 has an activation energy of $100 , kJ \cdot mol^{-1}$ and reaction 2 has an activation energy of $50 , kJ \cdot mol^{-1}$, and both reactions are exothermic with the same $\Delta H$, which reaction will proceed faster at the same temperature?

For the reaction $A ightarrow B$, the activation energy is $E_a$. If the reaction is reversed ($B ightarrow A$), what is the activation energy for the reverse reaction in terms of $E_a$ and $\Delta H_{forward}$?

Consider the combustion of methane ($CH_4 + 2O_2 ightarrow CO_2 + 2H_2O$), an exothermic reaction. Which of the following must be true regarding the energy changes?

In an energy diagram for an endothermic reaction, where is the activated complex located relative to the reactants and products?

For the reaction $N_2(g) + 3H_2(g) ightarrow 2NH_3(g)$, $\Delta H = -92 , kJ \cdot mol^{-1}$. What is the enthalpy change for the reaction $NH_3(g) ightarrow 1/2 N_2(g) + 3/2 H_2(g)$?

Which statement best describes the relationship between activation energy and the enthalpy change ($\Delta H$) for a reaction?

Imagine two different endothermic reactions, Reaction A and Reaction B. Reaction A has a larger activation energy than Reaction B. Assuming all other factors are equal, which reaction is likely to be slower?

What is the term for the energy required to break the chemical bond between two atoms?

When new bonds are formed during a chemical reaction, what happens to energy?

In an exothermic reaction, how does the energy of the products compare to the energy of the reactants?

Which of the following must be overcome for a chemical reaction to occur?

What is the term for the high-energy, unstable state that occurs during a chemical reaction where bonds are breaking and forming?

Which of the following statements correctly describes the enthalpy change ($\Delta H$) for an endothermic reaction?

Consider the reaction: $\text{Reactants} + ext{Energy} \rightarrow ext{Products}$. What type of reaction is this?

In terms of energy required vs. released, which scenario describes an exothermic reaction?

Which equation correctly includes the heat of reaction for an exothermic process?

On an energy diagram, what represents the activation energy for a reaction?

Why is energy required to break bonds in reactant molecules?

How is the heat of reaction ($\Delta H$) calculated?

What is the role of the 'activated complex' in a chemical reaction?

If a reaction has a positive $\Delta H$, what does this indicate about the energy change?

If for a given reaction, the energy of the products is much lower than the energy of the reactants, what can be inferred about the activation energy?

What is the impact on the energy graph when comparing an exothermic to an endothermic reaction?

Consider the energy diagram of a reaction. Which part of the diagram indicates the presence of an activated complex?

If $\Delta H = -100 kJ/mol$ for a reaction, and the activation energy is $50 kJ/mol$, what can be concluded about the reverse reaction?

Given the reaction: $\text{A} + ext{B} \rightarrow ext{C}$, sketch the energy profile if the activation energy is very high compared to an alternative reaction $\text{X} + ext{Y} \rightarrow ext{Z}$ that has a very low activation energy. What is the MOST likely observation?

If a reaction proceeds via multiple steps, how does the step with the highest activation energy affect the overall reaction rate?

When is the enthalpy change ($\Delta H$) most crucial in determining the spontaneity of a reaction, especially if we are considering whether a reaction will proceed without continuous external energy input?

The hypothetical element 'X' is able to form diatomic and triatomic allotropes (X$\text{_2}$ and X$\text{_3}$, respectively). Given that the formation of X$\text{_2}$ from individual X atoms is highly exothermic, but the formation of X$\text{_3}$ requires significant energy input, what can you infer about the stability and bond energies in these allotropes?

Which of the following methods could be used to calculate reliable heat of reaction ($\Delta H$) of complex reaction mechanism where multiple intermediate products are formed?