AP Chemistry 6.7 Bond Enthalpies (Energies) Study Notes - New Syllabus Effective fall 2024
AP Chemistry 6.7 Bond Enthalpies (Energies) Study Notes.- New syllabus
AP Chemistry 6.7 Bond Enthalpies (Energies) Study Notes – AP Chemistry – per latest AP Chemistry Syllabus.
LEARNING OBJECTIVE
Calculate the enthalpy change of a reaction based on the average bond energies of bonds broken and formed in the reaction.
Key Concepts:
- Enthalpy of Reaction
- Enthalpy & Bond Energies
- Enthalpy of Formation
- Hess’s Law
6.7.A.1 Changes in Potential Energy During a Chemical Reaction:
1. Bond Formation and Breaking:
Bond Breaking: Endothermic (energy is absorbed in bond breaking).
Bond Formation: Exothermic (energy given out as new bonds are being made).
Net energy change for a reaction will depend on energy required to break bonds and energy released when new bonds are formed.
2. Exothermic vs. Endothermic Reactions:
| Property | Exothermic Reactions | Endothermic Reactions |
|---|---|---|
| Energy Change | Energy is released (negative) | Energy is absorbed (positive) |
| Temperature Effect | Surroundings get warmer | Surroundings get cooler |
| Example | Combustion (e.g., burning wood or fuel) | Photosynthesis |
| Energy Flow | Energy flows out of the system | Energy flows into the system |
3. Activation Energy:
– Definition: This is energy used to initiate a chemical reaction. It’s the energy required to break the bonds between the reactants so that fresh bonds can be created between the products.
– Role in Potential Energy: Consider activation energy to be the barrier reactants must climb over to form products. On a potential energy diagram, it’s the summit reactants must scale, acquiring enough energy to make it over to settle into products, which are typically at a lower energy point.
Simply put, activation energy is the original push that gets a reaction started, which enables the reactants to arrive at a brief status (called the transition state) before they are able to convert into products.
4. Enthalpy (ΔH):
– Definition: Enthalpy (ΔH) is the overall change in the potential energy of a system during a chemical reaction.
– Significance:
– ΔH > 0: The reaction absorbs energy (endothermic reaction), if the change in enthalpy is positive.
– ΔH < 0: The reaction releases energy (exothermic reaction), if the change in enthalpy is negative.
– Role: It describes the difference in energy between the products and reactants, showing whether or not energy is released or taken in during the reaction.
In simple terms, ΔH is the total energy change from the beginning to the end of a reaction.
6.7.A.2 Estimating Energy Changes in Reactions: Exothermic vs. Endothermic:
1. Bond Energies:
– Definition: Bond energy is the amount of energy required to break one mole of a specific bond in a molecule into its constituent atoms in the gaseous state.
– How It Helps Estimate Energy:
– To break bonds, you need to supply energy that equals the bond energy of the broken bonds.
– For breaking bonds, energy needs to be equal to the breaking bond energies.
– Evaluating Energy Change in Reactions:
– Energy needed for breaking bonds = Sum of the bond energies of broken bonds.
– Energy on bond formation = Sum of the bond energies of formed bonds.
– Total Energy Change (ΔH) = Bond-breaking energy – Bond-forming energy.
In short, bond energy allows you to estimate the overall change in energy during a reaction based on the amount of energy that must be absorbed to break bonds and the energy released when forming bonds.
2. Exothermic vs. Endothermic Reactions:
| Property | Exothermic Reactions | Endothermic Reactions |
|---|---|---|
| Energy Change | Energy released > Energy required (ΔH < 0) | Energy released < Energy required (ΔH > 0) |
| Reaction | Energy released during bond formation | Energy absorbed to break bonds |
| Example | Combustion (e.g., burning fuel) | Photosynthesis, cooking |
| Effect on Surroundings | Surroundings get warmer | Surroundings get cooler |
3. Calculating Net Energy Change:
To figure out the net energy change of a reaction, you need to look at the energy it takes to break bonds compared to the energy that gets released when new bonds form:
Steps:
1. Energy to Break Bonds: Start by adding up the bond energies of all the bonds in the reactants that you need to break.
2. Energy Released During Bond Formation: Next, sum up the bond energies of the bonds that are formed in the products.
3. Net Energy Change (ΔH):
– ΔH = Energy to break bonds – Energy released in bond formation
### Interpreting the Results:
– Exothermic Reaction: If the energy released during bond formation is greater than the energy required to break the bonds, then ΔH is negative, meaning energy is released.
– Endothermic Reaction: If the energy required to break the bonds is greater than the energy released during bond formation, then ΔH is positive, indicating that energy is absorbed.
Example:
– Reactants: A-B + C-D → Energy to break bonds = 500 kJ
– Products: A-C + B-D → Energy released in bond formation = 600 kJ
ΔH = 500 kJ – 600 kJ = -100 kJ (This is an exothermic reaction, so energy is released.)
OLD Content
Bond Enthalpies (Energies)
- It is the energy stored in a chemical bond → gives information about the strength of a bonding interaction
(sigma = sum, D = bond)- *bond energies listed in textbook/reference