A-Level Chemistry: Gaseous State and Gas Behaviour | A-Level 化学：气态特征与气体行为

📚 A-Level Chemistry: Gaseous State and Gas Behaviour | A-Level 化学：气态特征与气体行为

The gaseous state is one of the three classical states of matter. Gases have neither fixed shape nor fixed volume, and their particles are in constant, random motion. Understanding how gases behave under different conditions is essential for solving many A-Level chemistry problems, from calculating molar masses to predicting reaction yields.

气态是物质的三种经典状态之一。气体既没有固定的形状，也没有固定的体积，其粒子处于持续的无规则运动中。理解气体在不同条件下的行为，对于解决许多A-Level化学问题至关重要，从计算摩尔质量到预测反应产率。

1. Characteristics of Gases | 气体的特征

Gases are distinguished from solids and liquids by several observable properties. They expand to fill their container completely, they are highly compressible, and they diffuse and effuse rapidly. Gases also have much lower densities than solids or liquids under normal conditions.

气体与固体、液体的区别在于若干可观察的性质。它们会完全膨胀充满容器，具有高度可压缩性，并且能快速扩散和逸散。在通常情况下，气体的密度远低于固体或液体。

No definite shape or volume – gases adopt the shape and volume of their container.

没有确定的形状或体积 – 气体采用容器的形状和体积。
High compressibility – the large spaces between particles allow gases to be compressed easily.

高度可压缩 – 粒子间距离大，使气体容易被压缩。
Low density – typical gas densities are about 1 kg m⁻³, much lower than liquids (≈1000 kg m⁻³).

密度低 – 典型气体密度约为 1 kg m⁻³，远低于液体（约 1000 kg m⁻³）。
Miscibility – gases mix completely in any proportion because intermolecular forces are negligible.

可混性 – 由于分子间作用力可忽略，气体以任意比例完全混合。

2. Kinetic Molecular Theory | 分子运动理论

The kinetic molecular theory explains gas behaviour using a simple model. It assumes that gases consist of tiny particles separated by large distances, and that these particles are in constant, random motion.

分子运动理论用一个简单模型解释气体行为。它假设气体由相距很远的大量微小粒子组成，且这些粒子处于持续的无规则运动中。

Particles are point masses with negligible volume relative to the gas container.

粒子是质点，相对于气体容器其体积可忽略不计。
No intermolecular forces exist between particles except during brief collisions.

除短暂碰撞外，粒子之间不存在分子间作用力。
Collisions are perfectly elastic – total kinetic energy is conserved.

碰撞是完全弹性的 – 总动能守恒。
Average kinetic energy is proportional to absolute temperature (in kelvin).

平均动能与绝对温度（开尔文）成正比。

The theory leads to the conclusion that pressure arises from collisions of gas particles with the walls of the container. Increasing temperature increases the average kinetic energy, hence the frequency and force of collisions, which increases pressure if volume is constant.

该理论得出结论：压强来源于气体粒子与容器壁的碰撞。升高温度会增加平均动能，从而增加碰撞的频率和力度，若体积不变则压强增大。

3. The Gas Laws | 气体定律

The gas laws describe the relationships between pressure (P), volume (V), temperature (T) and amount of gas (n). At A-Level, you must be able to state these laws and use them in calculations.

气体定律描述了压强（P）、体积（V）、温度（T）和气体物质的量（n）之间的关系。在A-Level考试中，你必须能够陈述这些定律并用于计算。

Law	Statement	Equation
Boyle’s Law 玻意耳定律	At constant n and T, P is inversely proportional to V. 在 n 和 T 恒定时，P 与 V 成反比。	P₁V₁ = P₂V₂
Charles’s Law 查理定律	At constant n and P, V is directly proportional to T (in K). 在 n 和 P 恒定时，V 与 T（K）成正比。	V₁/T₁ = V₂/T₂
Avogadro’s Law 阿伏伽德罗定律	At constant P and T, V is directly proportional to n. 在 P 和 T 恒定时，V 与 n 成正比。	V₁/n₁ = V₂/n₂

These laws are combined into the ideal gas equation, which is the most powerful tool for gas calculations.

这些定律被合并为理想气体方程，这是气体计算中最强大的工具。

4. The Ideal Gas Equation | 理想气体方程

The ideal gas equation links all four variables: pressure, volume, temperature and amount of substance. It is derived from Boyle’s, Charles’s and Avogadro’s laws.

理想气体方程将压强、体积、温度、物质的量这四个变量联系起来。它由玻意耳定律、查理定律和阿伏伽德罗定律推导而来。

PV = nRT

Where P is pressure, V is volume, n is number of moles, R is the molar gas constant, and T is temperature in kelvin.

其中 P 是压强，V 是体积，n 是摩尔数，R 是摩尔气体常数，T 是开尔文温度。

R = 8.31 J K⁻¹ mol⁻¹ (SI units).

R = 8.31 J K⁻¹ mol⁻¹（国际单位制）。
Pressure must be in pascals (Pa); volume in cubic metres (m³).

压强必须用帕斯卡（Pa）表示；体积用立方米（m³）表示。
Temperature must always be in kelvin (K), where T(K) = T(°C) + 273.15.

温度必须始终用开尔文（K），其中 T(K) = T(°C) + 273.15。

5. Units and Molar Calculations | 单位与摩尔计算

Common exam mistakes come from incorrect unit conversions. You must convert cm³ to m³ and kPa to Pa before substituting into the ideal gas equation.

常见的考试错误来自单位换算不正确。代入理想气体方程前，必须将 cm³ 转换为 m³、kPa 转换为 Pa。

1 m³ = 1000 dm³ = 1 000 000 cm³ | 1 atm = 101 325 Pa | 1 kPa = 1000 Pa

To find molar mass (M) from gas density, use the rearranged form:

要从气体密度求摩尔质量（M），使用重排形式：

m = n × M and ρ = m / V → M = ρRT / P

Example: A gas has density 1.78 kg m⁻³ at 273 K and 1.00 × 10⁵ Pa. Calculate its molar mass.

示例：一种气体在 273 K 和 1.00 × 10⁵ Pa 下密度为 1.78 kg m⁻³。计算其摩尔质量。

M = (1.78 × 8.31 × 273) / (1.00 × 10⁵) = 0.0404 kg mol⁻¹ = 40.4 g mol⁻¹

The gas is likely argon (Ar ≈ 40 g mol⁻¹).

该气体很可能是氩气（Ar ≈ 40 g mol⁻¹）。

6. Molar Volume and Gas Density | 摩尔体积与气体密度

The molar volume of an ideal gas at standard temperature and pressure (STP: 273 K, 1 atm) is 22.4 dm³ mol⁻¹. At room temperature and pressure (RTP: 298 K, 1 atm) it is 24.0 dm³ mol⁻¹.

在标准温度和压力（STP：273 K, 1 atm）下，理想气体的摩尔体积为 22.4 dm³ mol⁻¹。在室温常压（RTP：298 K, 1 atm）下为 24.0 dm³ mol⁻¹。

Condition	Temperature	Pressure	Molar Volume
STP	273 K (0 °C)	1 atm (101.3 kPa)	22.4 dm³ mol⁻¹
RTP	298 K (25 °C)	1 atm (101.3 kPa)	24.0 dm³ mol⁻¹

Density of a gas can be calculated from molar mass and molar volume: ρ = M / molar volume. For example, oxygen (M = 32 g mol⁻¹) has a density at STP of 32 / 22.4 ≈ 1.43 g dm⁻³.

气体密度可由摩尔质量和摩尔体积计算：ρ = M / 摩尔体积。例如，氧气（M = 32 g mol⁻¹）在 STP 下的密度为 32 / 22.4 ≈ 1.43 g dm⁻³。

7. Dalton’s Law of Partial Pressures | 道尔顿分压定律

In a mixture of gases, each gas exerts its own pressure as if it alone occupied the container. The total pressure is the sum of these partial pressures.

在混合气体中，每种气体都施加其自身的压强，就好像它单独占据容器一样。总压强是这些分压之和。

P_total = P₁ + P₂ + P₃ + …

The partial pressure of a gas is related to its mole fraction:

气体的分压与其摩尔分数有关：

P₁ = x₁ × P_total where x₁ = n₁ / n_total

This law is used to calculate the pressure of gases collected over water, where the total pressure includes the saturated vapour pressure of water.

该定律用于计算在水面上收集的气体的压强，此时总压强包括水的饱和蒸气压。

Example: 0.20 mol N₂ and 0.30 mol O₂ are placed in a 5.0 dm³ container at 300 K. Calculate the partial pressures.

示例：0.20 mol N₂ 和 0.30 mol O₂ 置于 5.0 dm³ 容器中，温度 300 K。计算分压。

n_total = 0.50 mol; x(N₂) = 0.20/0.50 = 0.40 | P_total = nRT/V = (0.50 × 8.31 × 300) / 0.005 = 249 300 Pa

P(N₂) = 0.40 × 249 300 ≈ 99 700 Pa | P(O₂) = 0.60 × 249 300 ≈ 149 600 Pa

8. Graham’s Law of Effusion and Diffusion | 格雷厄姆逸散与扩散定律

Graham’s law states that the rate of effusion (or diffusion) of a gas is inversely proportional to the square root of its molar mass.

格雷厄姆定律指出：气体的逸散（或扩散）速率与其摩尔质量的平方根成反比。

rate₁ / rate₂ = √(M₂ / M₁)

Lighter gases move faster than heavier gases at the same temperature. For example, hydrogen (M = 2 g mol⁻¹) effuses about four times faster than oxygen (M = 32 g mol⁻¹), since √(32/2) = √16 = 4.

在相同温度下，较轻的气体比较重的气体运动得更快。例如，氢气（M = 2 g mol⁻¹）逸散速率约为氧气（M = 32 g mol⁻¹）的四倍，因为 √(32/2) = √16 = 4。

This law can be used to separate isotopes or to identify an unknown gas by comparing its rate with a known gas.

该定律可用于分离同位素，或通过比较未知气体与已知气体的速率来鉴别未知气体。

9. Deviations from Ideal Behaviour | 对理想行为的偏离

Real gases do not obey the ideal gas equation perfectly. The ideal gas model assumes point particles with no intermolecular forces, but real gas particles have finite volume and attract each other.

真实气体并不能完美地遵守理想气体方程。理想气体模型假设粒子为质点且无分子间作用力，但真实气体粒子具有有限体积并相互吸引。

At high pressure, the volume of the particles becomes significant, so the actual volume available is less than the container volume. At low temperature, intermolecular attractions become stronger, reducing the impact of collisions and causing the observed pressure to be lower than ideal.

在高压下，粒子自身体积变得显著，因此实际可用的体积小于容器体积。在低温下，分子间吸引力变强，减弱了碰撞的冲击力，导致实际压强低于理想值。

Condition	Main cause of deviation	Effect on PV/RT
High pressure 高压	Particle volume not negligible 粒子体积不可忽略	PV/RT increases PV/RT 增大
Low temperature 低温	Intermolecular attractions 分子间吸引力	PV/RT decreases PV/RT 减小

The van der Waals equation is one correction to the ideal gas law, adding terms for particle volume (b) and intermolecular forces (a).

范德华方程是对理想气体定律的一种修正，引入了粒子体积项（b）和分子间作用力项（a）。

(P + an²/V²) (V – nb) = nRT

10. Conditions for Ideal Behaviour | 有利于理想行为的条件

Real gases approximate ideal behaviour when the assumptions of the kinetic theory hold most closely. This occurs at low pressure and high temperature.

当分子运动理论的假设最接近成立时，真实气体近似于理想行为。这发生在低压和高温条件下。

Low pressure – particles are far apart, so their own volume is negligible and intermolecular forces are minimal.

低压 – 粒子相距远，因此粒子自身体积可忽略，分子间作用力极小。
High temperature – kinetic energy overcomes intermolecular attractions, so collisions are effectively elastic.

高温 – 动能克服分子间吸引力，因此碰撞近似完全弹性。
Noble gases (He, Ne, Ar) behave most ideally because they have weak dispersion forces and small, non-polar atoms.

稀有气体（He、Ne、Ar）行为最接近理想，因为它们的色散力弱且原子小、无极性。

Helium is nearly ideal even at moderate pressures because its atoms are tiny and show almost no attraction for each other.

氦气即使在中等压强下也接近理想，因为其原子极小且彼此几乎不表现吸引力。

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A-Level Chemistry: Gaseous State and Gas Behaviour | A-Level 化学：气态特征与气体行为