Photosynthetic plants use the following reaction to produce glucose, cellulose, and so forth: $$6 \mathrm{CO}_{2}(g)+6 \mathrm{H}_{2} \mathrm{O}(l) \frac{\text { Sunlight }}{\longrightarrow} \mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6}(s)+6 \mathrm{O}_{2}(g)$$ How might extensive destruction of forests exacerbate the greenhouse effect?

Short Answer

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Extensive destruction of forests exacerbates the greenhouse effect by reducing the rate of photosynthesis due to the loss of photosynthesizing surface, releasing stored carbon from plant biomass back into the atmosphere as CO₂, and disrupting the water cycle. This leads to increased atmospheric CO₂ levels and an increased greenhouse effect, causing global temperatures to rise and resulting in negative consequences such as extreme weather events, floods, sea-level rise, and loss of biodiversity.

Step by step solution

01

Understanding Photosynthesis

Photosynthesis is the process by which plants, algae, and some bacteria convert carbon dioxide (CO₂) and water (H₂O) into glucose molecules (C₆H₁₂O₆) and oxygen (O₂) using sunlight as an energy source. The chemical equation for this process can be represented as: \[6 CO_{2}(g) + 6 H_{2}O(l) \xrightarrow{Sunlight} C_{6}H_{12}O_{6}(s) + 6 O_{2}(g)\]
02

Role of Forests in Photosynthesis

Forests are responsible for a significant proportion of the Earth's photosynthesis process. They provide a large surface area for sunlight absorption, have a vast number of plant species, and have a high rate of water transpiration, which aids in the process. As plants perform photosynthesis, they consume CO₂, removing it from the atmosphere, and release O₂ as a byproduct.
03

Forests and Carbon Sequestration

Forests not only help in the photosynthesis process but also play a vital role in storing carbon. This process, known as carbon sequestration, occurs when plants use CO₂ to produce glucose and other organic compounds during photosynthesis, helping to store carbon in large quantities in the form of plant biomass (wood, leaves, and roots).
04

Effects of Forest Destruction

Extensive forest destruction through activities such as deforestation, forest fires, or clearing for agriculture or development results in a significant reduction of the Earth's photosynthesizing surface. The destruction of forests results in the following consequences: 1. Reduced rate of photosynthesis: With fewer plants available to consume CO₂ and produce O₂, the overall rate of photosynthesis decreases. 2. Release of stored carbon: When forests are destroyed, the stored carbon in the plant biomass is released back into the atmosphere as CO₂, either through decomposition or combustion. 3. Disruption of the water cycle: The destruction of forests can also lead to a disruption in the water cycle, further affecting the photosynthesis process.
05

Exacerbating the Greenhouse Effect

By extensively destroying forests, we will see an increase in atmospheric CO₂ levels due to reduced photosynthesis and the release of stored carbon. The consequent increase in greenhouse gases will trap more heat in the Earth's atmosphere, causing global temperatures to rise and exacerbate the greenhouse effect. This can lead to various negative consequences such as extreme weather events, floods, sea-level rise, and loss of biodiversity.

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Most popular questions from this chapter

Calculate \(w\) and \(\Delta E\) when 1 mole of a liquid is vaporized at its boiling point \(\left(80 .^{\circ} \mathrm{C}\right)\) and 1.00 atm pressure. \(\Delta H_{\text { vap }}\) for the liquid is 30.7 \(\mathrm{kJ} / \mathrm{mol}\) at \(80 .^{\circ} \mathrm{C} .\)

Combustion of table sugar produces \(\mathrm{CO}_{2}(g)\) and $\mathrm{H}_{2} \mathrm{O}(l) .$ When 1.46 \(\mathrm{g}\) table sugar is combusted in a constant-volume (bomb) calorimeter, 24.00 \(\mathrm{kJ}\) of heat is liberated. a. Assuming that table sugar is pure sucrose, $\mathrm{C}_{12} \mathrm{H}_{22} \mathrm{O}_{11}(s)$ write the balanced equation for the combustion reaction. b. Calculate \(\Delta E\) in $\mathrm{kJ} / \mathrm{mol} \mathrm{C}_{12} \mathrm{H}_{22} \mathrm{O}_{11}$ for the combustion reaction of sucrose.

You have a 1.00 -mole sample of water at \(-30 .^{\circ} \mathrm{C}\) and you heat it until you have gaseous water at \(140 .^{\circ} \mathrm{C}\) . Calculate \(q\) for the entire process. Use the following data. $$ \begin{aligned} \text { Specific heat capacity of ice } &=2.03 \mathrm{J} /^{\circ} \mathrm{C} \cdot \mathrm{g} \\ \text { Specific heat capacity of water } &=4.18 \mathrm{J} /^{\circ} \mathrm{C} \cdot \mathrm{g} \\ \text { Specific heat capacity of steam } &=2.02 \mathrm{J} /^{\circ} \mathrm{C} \cdot \mathrm{g} \end{aligned} $$ $$ \mathrm{H}_{2} \mathrm{O}(s) \longrightarrow \mathrm{H}_{2} \mathrm{O}(l) \quad \Delta H_{\mathrm{fision}}=6.02 \mathrm{kJ} / \mathrm{mol}\left(\mathrm{at} 0^{\circ} \mathrm{C}\right) $$ $$ \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{H}_{2} \mathrm{O}(g) \quad \Delta H_{\mathrm{vaporization}}=40.7 \mathrm{kJ} / \mathrm{mol}\left(\mathrm{at} 100 .^{\circ} \mathrm{C}\right) $$

The reaction $$ \mathrm{SO}_{3}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{H}_{2} \mathrm{SO}_{4}(a q) $$ is the last step in the commercial production of sulfuric acid. The enthalpy change for this reaction is \(-227 \mathrm{kJ} .\) In designing a sulfuric acid plant, is it necessary to provide for heating or cooling of the reaction mixture? Explain.

Nitromethane, \(\mathrm{CH}_{3} \mathrm{NO}_{2},\) can be used as a fuel. When the liquid is burned, the (unbalanced) reaction is mainly $$ \mathrm{CH}_{3} \mathrm{NO}_{2}(l)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{CO}_{2}(g)+\mathrm{N}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(g) $$ a. The standard enthalpy change of reaction $\left(\Delta H_{\mathrm{rxn}}^{\circ}\right)$ for the balanced reaction (with lowest whole-number coefficients \()\) is $-1288.5 \mathrm{kJ} .\( Calculate \)\Delta H_{\mathrm{f}}^{\circ}$ for nitromethane. b. A 15.0 -L flask containing a sample of nitromethane is filled with \(\mathrm{O}_{2}\) and the flask is heated to \(100 .^{\circ} \mathrm{C}\) . At this temperature, and after the reaction is complete, the total pressure of all the gases inside the flask is 950 . torr. If the mole fraction of nitrogen \(\left(\chi_{\text { nitrogen }}\right)\) is 0.134 after the reaction is complete, what mass of nitrogen was produced?

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