Energy Changes in Reactions
Chapter 2 | Section 3: Exothermic and Endothermic Changes
1. Energy in Chemical Reactions
In every chemical reaction, bonds in reactants are broken (requires energy) and new bonds in products are formed (releases energy). The difference determines the energy change.
Exothermic Reactions
Chemical reactions which proceed with the release of heat energy.
- Temperature of surroundings increases.
- Energy of Reactants > Energy of Products.
$A + B \rightarrow C + D + \text{Heat}$
Endothermic Reactions
Chemical reactions which proceed with the absorption of heat energy.
- Temperature of surroundings decreases.
- Energy of Products > Energy of Reactants.
$A + B + \text{Heat} \rightarrow C + D$
2. Key Examples
Exothermic Examples:
- Respiration: $C_6H_{12}O_6 + 6O_2 \rightarrow 6CO_2 + 6H_2O + \text{Energy}$
- Burning of Coal: $C + O_2 \rightarrow CO_2 + \Delta$
- Reaction of Water with Quicklime: $CaO + H_2O \rightarrow Ca(OH)_2 + \text{Heat}$
Endothermic Examples:
- Photosynthesis: $6CO_2 + 6H_2O + \text{Light} \rightarrow C_6H_{12}O_6 + 6O_2$
- Formation of Nitric Oxide: $N_2 + O_2 \xrightarrow{\Delta} 2NO$
- Decomposition of Limestone: $CaCO_3 \xrightarrow{\Delta} CaO + CO_2$
3. Activation Energy
The minimum amount of energy required to start a chemical reaction is called **Activation Energy**. Even exothermic reactions (like burning a matchstick) need a little "spark" to get started.
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