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When you say "why aren't things being destroyed", you presumably mean "why aren't the chemical bonds that hold objects together being broken". Now, we can determine the energy it takes to break a bond - that's called the "bond energy". Let's take, for example, a carbon-carbon bond, since it's a common one in our bodies.

Then bond energy of a carbon-carbon bond is $348\,\rm kJ/mol$, which works out to $5.8 \cdot 10^{-19}\,\rm J$ per bond. If an impacting gas molecule is to break this bond, it must (in a simplified collision scenario) have at least that much energy to break the bond. If the average molecule has that much energy, we can calculate what the temperature of the gas must be:

$$E_\text{average} = k T$$ $$T = \frac{5.8 \cdot 10^{-19}\,\rm J}{1.38 \cdot 10^{-23}\,\rm m^2 kg\, s^{-2} K^{-1}}$$ $$T = 41,580\rm °C$$

That's pretty hot!

Now, even if the average molecule doesn't have that energy, some of the faster-moving ones might. Let's calculate the percentage that have that energy at room temperature using the Boltzmann distribution for particle energy:

$$f_E(E) = \sqrt{\frac{4 E}{\pi (kT)^3}} \exp\left(\frac{-E}{kT} \right)$$

The fraction of particles with energy greater than or equal to that amount should be given by this integral:

$$p(E \ge E_0) = \int_{E_0}^{\infty} f_E(E) dE$$

where $E_0 = 5.8 \cdot 10^{-19}\,\rm J$ yields $p(E \ge E_0) = 1.9 \cdot 10^{-61}$.

So, the fraction of molecules at room temperature with sufficient kinetic energy to break a carbon-carbon bond is $1.9 \cdot 10^{-61}$, an astoundingly small number. To put that in perspective, if you filled a sphere the size of Earth's orbit around the sun with gas at STP, you would need around 16 of those spheres to expect to have even one gas particle with that amount of energy.

So that's why these "torpedoes" don't destroy things generally - they aren't moving fast enough at room temperature to break chemical bonds!

When you say "why aren't things being destroyed", you presumably mean "why aren't the chemical bonds that hold objects together being broken". Now, we can determine the energy it takes to break a bond - that's called the "bond energy". Let's take, for example, a carbon-carbon bond, since it's a common one in our bodies.

The bond energy of a carbon-carbon bond is $348\,\rm kJ/mol$, which works out to $5.8 \cdot 10^{-19}\,\rm J$ per bond. If an impacting gas molecule is to break this bond, it must (in a simplified collision scenario) have at least that much energy to break the bond. If the average molecule has that much energy, we can calculate what the temperature of the gas must be:

$$E_\text{average} = k T$$ $$T = \frac{5.8 \cdot 10^{-19}\,\rm J}{1.38 \cdot 10^{-23}\,\rm m^2 kg\, s^{-2} K^{-1}}$$ $$T = 41,580\rm °C$$

That's pretty hot!

Now, even if the average molecule doesn't have that energy, some of the faster-moving ones might. Let's calculate the percentage that have that energy at room temperature using the Boltzmann distribution for particle energy:

$$f_E(E) = \sqrt{\frac{4 E}{\pi (kT)^3}} \exp\left(\frac{-E}{kT} \right)$$

The fraction of particles with energy greater than or equal to that amount should be given by this integral:

$$p(E \ge E_0) = \int_{E_0}^{\infty} f_E(E) dE$$

In our situation, $E_0 = 5.8 \cdot 10^{-19}\,\rm J$, and this expression yields $p(E \ge E_0) = 1.9 \cdot 10^{-61}$.

So, the fraction of molecules at room temperature with sufficient kinetic energy to break a carbon-carbon bond is $1.9 \cdot 10^{-61}$, an astoundingly small number. To put that in perspective, if you filled a sphere the size of Earth's orbit around the sun with gas at STP, you would need around 16 of those spheres to expect to have even one gas particle with that amount of energy.

So that's why these "torpedoes" don't destroy things generally - they aren't moving fast enough at room temperature to break chemical bonds!

When you say "why aren't things being destroyed", you presumably mean "why aren't the chemical bonds that hold objects together being broken". Now, we can determine the energy it takes to break a bond - that's called the "bond energy". Let's take, for example, a carbon-carbon bond, since it's a common one in our bodies.

Then bond energy of a carbon-carbon bond is $348\,\rm kJ/mol$, which works out to $5.8 \cdot 10^{-19}\,\rm J$ per bond. If an impacting gas molecule is to break this bond, it must (in a simplified collision scenario) have at least that much energy to break the bond. If the average molecule has that much energy, we can calculate what the temperature of the gas must be:

$$E_\text{average} = k T$$ $$T = \frac{5.8 \cdot 10^{-19}\,\rm J}{1.38 \cdot 10^{-23}\,\rm m^2 kg\, s^{-2} K^{-1}}$$ $$T = 41,580\rm °C$$

That's pretty hot!

Now, even if the average molecule doesn't have that energy, some of the faster-moving ones might. Let's calculate the percentage that have that energy at room temperature using the Boltzmann distribution for particle energy:

$$f_E(E) = \sqrt{\frac{4 E}{\pi (kT)^3}} \exp\left(\frac{-E}{kT} \right)$$

The fraction of particles with energy greater than or equal to that amount should be given by this integral:

$$\int_{E_0}^{\infty} f_E(E) dE$$

where $E_0 = 5.8 \cdot 10^{-19}\,\rm J$ yields $p(E \ge E_0) = 1.9 \cdot 10^{-61}$.

So, the fraction of molecules at room temperature with sufficient kinetic energy to break a carbon-carbon bond is $1.9 \cdot 10^{-61}$, an astoundingly small number. To put that in perspective, if you filled a sphere the size of Earth's orbit around the sun with gas at STP, you would need around 16 of those spheres to expect to have even one gas particle with that amount of energy.

So that's why these "torpedoes" don't destroy things generally - they aren't moving fast enough at room temperature to break chemical bonds!

When you say "why aren't things being destroyed", you presumably mean "why aren't the chemical bonds that hold objects together being broken". Now, we can determine the energy it takes to break a bond - that's called the "bond energy". Let's take, for example, a carbon-carbon bond, since it's a common one in our bodies.

Then bond energy of a carbon-carbon bond is $348\,\rm kJ/mol$, which works out to $5.8 \cdot 10^{-19}\,\rm J$ per bond. If an impacting gas molecule is to break this bond, it must (in a simplified collision scenario) have at least that much energy to break the bond. If the average molecule has that much energy, we can calculate what the temperature of the gas must be:

$$E_\text{average} = k T$$ $$T = \frac{5.8 \cdot 10^{-19}\,\rm J}{1.38 \cdot 10^{-23}\,\rm m^2 kg\, s^{-2} K^{-1}}$$ $$T = 41,580\rm °C$$

That's pretty hot!

Now, even if the average molecule doesn't have that energy, some of the faster-moving ones might. Let's calculate the percentage that have that energy at room temperature using the Boltzmann distribution for particle energy:

$$f_E(E) = \sqrt{\frac{4 E}{\pi (kT)^3}} \exp\left(\frac{-E}{kT} \right)$$

The fraction of particles with energy greater than or equal to that amount should be given by this integral:

$$p(E \ge E_0) = \int_{E_0}^{\infty} f_E(E) dE$$

where $E_0 = 5.8 \cdot 10^{-19}\,\rm J$ yields $p(E \ge E_0) = 1.9 \cdot 10^{-61}$.

So, the fraction of molecules at room temperature with sufficient kinetic energy to break a carbon-carbon bond is $1.9 \cdot 10^{-61}$, an astoundingly small number. To put that in perspective, if you filled a sphere the size of Earth's orbit around the sun with gas at STP, you would need around 16 of those spheres to expect to have even one gas particle with that amount of energy.

So that's why these "torpedoes" don't destroy things generally - they aren't moving fast enough at room temperature to break chemical bonds!

When you say "why aren't things being destroyed", you presumably mean "why aren't the chemical bonds that hold objects together being broken". Now, we can determine the energy it takes to break a bond - that's called the "bond energy". Let's take, for example, a carbon-carbon bond, since it's a common one in our bodies.

The bond energy of a carbon-carbon bond is $348~kJ/mol$, which works out to $5.8 \cdot 10^{-19} J$ per bond. If an impacting gas molecule is to break this bond, it must (in a simplified collision scenario) have at least that much energy to break the bond. If the average molecule has that much energy, we can calculate what the temperature of the gas must be:

$$E_{average} = k T$$ $$T = \frac{5.8 \cdot 10^{-19} J}{1.38 \cdot 10^{-23} m^2 kg s^{-2} K^{-1}}$$ $$T = 41,580 °C$$

That's pretty hot!

Now, even if the average molecule doesn't have that energy, some of the faster-moving ones might. Let's calculate the percentage that have that energy at room temperature using the Boltzmann distribution for particle energy:

$$f_E(E) = \sqrt{\frac{4 E}{\pi (kT)^3}} \exp\left(\frac{-E}{kT} \right)$$

The fraction of particles with energy greater than or equal to that amount should be given by this integral:

$$\int_{E_0}^{\infty} f_E(E) dE$$

where $E_0 = 5.8 \cdot 10^{-19}~J$ yields $p(E \ge E_0) = 1.9 \cdot 10^{-61}$.

So, the fraction of molecules at room temperature with sufficient kinetic energy to break a carbon-carbon bond is $1.9 \cdot 10^{-61}$, an astoundingly small number. To put that in perspective, if you filled a sphere the size of Earth's orbit around the sun with gas at STP, you would need around 16 of those spheres to expect to have even one gas particle with that amount of energy.

So that's why these "torpedoes" don't destroy things generally - they aren't moving fast enough at room temperature to break chemical bonds!

When you say "why aren't things being destroyed", you presumably mean "why aren't the chemical bonds that hold objects together being broken". Now, we can determine the energy it takes to break a bond - that's called the "bond energy". Let's take, for example, a carbon-carbon bond, since it's a common one in our bodies.

Then bond energy of a carbon-carbon bond is $348\,\rm kJ/mol$, which works out to $5.8 \cdot 10^{-19}\,\rm J$ per bond. If an impacting gas molecule is to break this bond, it must (in a simplified collision scenario) have at least that much energy to break the bond. If the average molecule has that much energy, we can calculate what the temperature of the gas must be:

$$E_\text{average} = k T$$ $$T = \frac{5.8 \cdot 10^{-19}\,\rm J}{1.38 \cdot 10^{-23}\,\rm m^2 kg\, s^{-2} K^{-1}}$$ $$T = 41,580\rm °C$$

That's pretty hot!

Now, even if the average molecule doesn't have that energy, some of the faster-moving ones might. Let's calculate the percentage that have that energy at room temperature using the Boltzmann distribution for particle energy:

$$f_E(E) = \sqrt{\frac{4 E}{\pi (kT)^3}} \exp\left(\frac{-E}{kT} \right)$$

The fraction of particles with energy greater than or equal to that amount should be given by this integral:

$$\int_{E_0}^{\infty} f_E(E) dE$$

where $E_0 = 5.8 \cdot 10^{-19}\,\rm J$ yields $p(E \ge E_0) = 1.9 \cdot 10^{-61}$.

So, the fraction of molecules at room temperature with sufficient kinetic energy to break a carbon-carbon bond is $1.9 \cdot 10^{-61}$, an astoundingly small number. To put that in perspective, if you filled a sphere the size of Earth's orbit around the sun with gas at STP, you would need around 16 of those spheres to expect to have even one gas particle with that amount of energy.

So that's why these "torpedoes" don't destroy things generally - they aren't moving fast enough at room temperature to break chemical bonds!