Pressure Units: Videos & Practice Problems

The SI unit for pressure is the Pascal (Pa), named after the French mathematician Blaise Pascal. Pressure can be understood as the force exerted by gas molecules on the walls of their container. In a gas, molecules move randomly in straight lines, colliding with each other and the container's surfaces. These collisions are what create pressure.

To quantify pressure, we can use the formula:

\( P = \frac{F}{A} \)

where \( P \) represents pressure in Pascals, \( F \) is the force in newtons (N), and \( A \) is the area in square meters (m²). This relationship highlights that pressure is the force applied per unit area. Understanding this concept is crucial as we delve into the behavior of ideal gases and later, non-ideal gases, which will further illustrate how gas molecules interact within a confined space.

In chemistry, pressure is a crucial concept, and several non-SI units are commonly used to express it. The most significant units include atmospheres (atm), millimeters of mercury (mmHg), and torr. Understanding the relationships between these units is essential, especially since they often appear in exams without being explicitly provided.

Key pressure values to memorize are:

• 1 atm = 760 mmHg
• 1 atm = 760 torr

While these units are vital, other units like pascals (Pa), kilopascals (kPa), bars, and pounds per square inch (psi) are also used, though they are less frequently required to be memorized. Their values are as follows:

• 1 atm = 101.325 kPa = 101325 Pa
• 1 bar = 100 kPa = 100000 Pa
• 1 psi = 14.696 psi

These relationships allow chemists to convert between different pressure units easily. For example, knowing that 1 atm equals 14.696 psi can help in various calculations. While the first three units (atm, mmHg, and torr) are the most commonly used in chemistry, familiarity with the others can enhance understanding and application in different contexts.

With this foundational knowledge of pressure units, one can confidently approach problems and examples involving pressure in chemical contexts.

To convert the average pressure in Denver, Colorado, from inches of mercury to millimeters of mercury and then to atmospheres, we can utilize dimensional analysis. The average pressure is given as 24.9 inches of mercury. The first step is to convert inches of mercury to centimeters of mercury, using the conversion factor that 1 inch equals 2.54 centimeters.

Starting with the pressure:

24.9 \text{ inches of mercury} \times \frac{2.54 \text{ cm}}{1 \text{ inch}} = 63.246 \text{ cm of mercury}

Next, we convert centimeters of mercury to millimeters of mercury, knowing that 1 centimeter equals 10 millimeters:

63.246 \text{ cm of mercury} \times \frac{10 \text{ mm}}{1 \text{ cm}} = 632.46 \text{ mm of mercury}

Since the original value has three significant figures, we round this to:

632 \text{ mm of mercury}

Now, to convert millimeters of mercury to atmospheres, we use the conversion factor that 1 atmosphere equals 760 millimeters of mercury:

632.46 \text{ mm of mercury} \times \frac{1 \text{ atm}}{760 \text{ mm}} \approx 0.832 \text{ atm}

Thus, the final results are:

Pressure in millimeters of mercury: 632 mm of mercury

Pressure in atmospheres: 0.832 atm

This process illustrates how dimensional analysis can effectively convert between different units of pressure.