Chemistry Lab: Gas Laws Experiment

Gas laws are a set of rules that describe the behaviour of gases under various conditions of temperature, pressure, and volume. The first experiments to quantify the relationship between temperature and volume were carried out by French chemist Jacques Alexandre César Charles in 1783, and later by Joseph-Louis Gay-Lussac. These experiments laid the foundation for what is now known as Charles' Law. Another important figure in the study of gas laws is Robert Boyle, who conducted early experiments to determine the relationship between the pressure and volume of a gas, resulting in what is now known as Boyle's Law. Gas laws can be studied through various experiments, such as observing the behaviour of gases in a balloon or a bottle, and they have practical applications, like explaining why car manufacturers recommend adjusting tyre pressure before a long trip.

Charles' Law

Charles's Law, also known as the Law of Volumes, is an experimental gas law that describes how gases tend to expand when heated. It was formulated by Jacques Charles in the 1780s and presented to the French National Institute in 1802 by Joseph Louis Gay-Lussac, who credited Charles for the discovery.

Charles's Law states that the volume of a gas is directly proportional to its absolute temperature, provided that the pressure and the amount of gas remain constant. This relationship can be expressed mathematically as V1/T1=V2/T2, where V represents volume and T represents temperature in Kelvin. This means that as the temperature of a gas increases, its volume will also increase proportionally, and conversely, a decrease in temperature will lead to a decrease in volume.

The law can be demonstrated through experiments that involve collecting gas in a balloon or an inverted graduated cylinder. By determining the volume of the balloon and the volume of water displaced in the cylinder, along with using the ideal gas law, the number of moles of gas can be calculated. This allows for an evaluation of the efficiency of the experimental methods.

Charles's Law is a special case of the general gas law and can be derived from the kinetic theory of gases, which relates the macroscopic properties of gases (such as pressure and volume) to the microscopic properties of the gas molecules (such as mass and speed). It is important to note that Charles's Law implies that the volume of a gas will decrease to zero at a certain temperature, known as absolute zero, which is −266.66 °C or −273.15 °C.

Gay-Lussac's experiments

Gay-Lussac's Law, discovered by French chemist Joseph Gay-Lussac in 1808, states that the pressure exerted by a gas (of a given mass and kept at a constant volume) varies directly with the absolute temperature of the gas. In other words, the pressure exerted by a gas is proportional to the temperature of the gas when the mass is fixed and the volume is constant.

Gay-Lussac's Law can be used to solve for unknown pressure. The temperatures must first be converted to Kelvin, and then the equation must be rearranged algebraically to solve for the unknown pressure variable. The known quantities are then substituted into the equation and solved.

Gay-Lussac's Law is a variant of the ideal gas law, where the volume of gas is held constant. The pressure of a gas is directly proportional to its temperature while the volume is kept constant. This can be expressed as P / T = constant or Pi / Ti = Pf / Tf. This means that raising a gas's temperature induces a relative increase in its pressure, assuming that the volume does not change. Conversely, reducing the temperature allows the pressure to decrease proportionally.

Gay-Lussac's Law is very similar to Charles's Law, with the only difference being the type of container. The container in a Charles's Law experiment is flexible, while it is rigid in a Gay-Lussac's Law experiment. Gay-Lussac attributed his findings to Jacques Charles because he used much of Charles's unpublished data from the 1780s.

Boyle's Law

Additionally, Boyle's Law can be applied to understand the process of breathing. When we inhale, we increase the volume of our chest cavity, creating low pressure in our lungs, which causes air to be sucked into them. Conversely, when we exhale, we decrease the volume of our chest cavity, creating high pressure, and air leaves our lungs.

Amontons' Law

$$ P \propto T \quad \text{or} \quad P = \text{constant} \times T \quad \text{or} \quad P = k \times T$$

Where ∝ means “is proportional to”, and k is a proportionality constant that depends on the identity, amount, and volume of the gas.

Dalton's Law of Partial Pressures

Gas laws are a set of rules that define the physical behaviour and properties of gases under various conditions of temperature, pressure, and volume. Dalton's Law of Partial Pressures, formulated by John Dalton in 1801 and published in 1802, is one such law.

Dalton's Law states that in a mixture of non-reacting gases, the total pressure exerted is equal to the sum of the partial pressures of the individual gases. In other words, each gas in a mixture contributes its own pressure to the overall pressure of the system. This law is particularly useful for understanding and predicting the behaviour of gas mixtures, such as air, which is composed of several different gases.

Mathematically, the pressure of a mixture of non-reactive gases can be expressed as the sum of the partial pressures of each component:

> {\displaystyle p_{\text{total}}=\sum _{i=1}^{n}p_{i}=p_{1}+p_{2}+p_{3}+\cdots +p_{n}}

Where p1, p2, ..., pn represent the partial pressures of each individual gas in the mixture. The variable xi represents the mole fraction of the ith component in the total mixture of n components.

It is important to note that Dalton's Law assumes ideal gas behaviour, which means that the gases have no intermolecular forces and occupy zero volume. In reality, gases do not always behave ideally, especially at high pressures and low temperatures. Therefore, Dalton's Law is not strictly followed by real gases, and the deviation from the law increases with pressure.

In experimental settings, Dalton's Law of Partial Pressures can be applied to determine the pressure of a gas in a mixture. For example, in a chemistry experiment, a gas may be collected in a balloon or an inverted graduated cylinder. By measuring the volume of the balloon or the volume of water displaced in the cylinder, along with other variables such as temperature and pressure, the ideal gas law can be used to calculate the number of moles of gas present. This allows for a comparison between the expected and actual yields, providing an evaluation of the efficiency of the experimental methods.

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