During a reaction, the enthalpy change due to breaking bonds is typically positive, while the enthalpy change due to forming bonds is typically negative. Which represents the net enthalpy change?

During a reaction, the enthalpy change due to breaking bonds is typically positive, while the enthalpy change due to forming bonds is typically negative. Which represents the net enthalpy change?

• A. Net enthalpy change is always zero
• B. Net enthalpy change is equal to ∆H formed minus ∆H broken
• C. Net enthalpy change is ∆H broken minus ∆H formed
• D. Net enthalpy change depends only on temperature

Answer: (C)

The net enthalpy change comes from two competing processes in a reaction: breaking bonds in the reactants (which requires energy) and forming new bonds in the products (which releases energy). To get the overall ΔH, you compare how much energy you had to input to break bonds with how much energy you recover by forming bonds. In bond-energy terms, you sum the energy needed to break all bonds and then subtract the energy released when new bonds form. If you denote the energy to break bonds as ΔH broken (a positive value) and the energy released on forming bonds as ΔH formed (taken as the amount released, a positive magnitude), the net enthalpy change is ΔH broken minus ΔH formed. This matches the standard way of balancing the two contributions: energy in minus energy out. For example, if breaking bonds costs 120 kJ and forming bonds releases 90 kJ, the net enthalpy change is 120 − 90 = 30 kJ (endothermic). Note: if you keep track of signs with ΔH formed as the actual signed enthalpy change (negative for formation), you would write the net as ΔHbroken + ΔHformed. Both ways describe the same energy balance; the subtraction form using magnitudes is the common teaching convention.

The net enthalpy change comes from two competing processes in a reaction: breaking bonds in the reactants (which requires energy) and forming new bonds in the products (which releases energy). To get the overall ΔH, you compare how much energy you had to input to break bonds with how much energy you recover by forming bonds.

In bond-energy terms, you sum the energy needed to break all bonds and then subtract the energy released when new bonds form. If you denote the energy to break bonds as ΔH broken (a positive value) and the energy released on forming bonds as ΔH formed (taken as the amount released, a positive magnitude), the net enthalpy change is ΔH broken minus ΔH formed. This matches the standard way of balancing the two contributions: energy in minus energy out.

For example, if breaking bonds costs 120 kJ and forming bonds releases 90 kJ, the net enthalpy change is 120 − 90 = 30 kJ (endothermic).

Note: if you keep track of signs with ΔH formed as the actual signed enthalpy change (negative for formation), you would write the net as ΔHbroken + ΔHformed. Both ways describe the same energy balance; the subtraction form using magnitudes is the common teaching convention.