# 5.2 Physical Properties of Liquids This section explores the key physical properties of liquids, including vapour pressure, boiling point, evaporation, viscosity, and surface tension, with a special focus on how hydrogen bonding influences the unique properties of water. ## 5.2.1 Evaporation *Evaporation* is the spontaneous change of a liquid into its vapour phase at the surface. This process can occur at any temperature. * **Mechanism:** Molecules within a liquid have a range of kinetic energies. Those with sufficiently high kinetic energy can overcome the intermolecular attractive forces at the surface and escape into the gas phase. * **Cooling Effect:** When the highest-energy molecules escape, the average kinetic energy of the remaining liquid molecules decreases. This results in a drop in the liquid's temperature. To regain thermal equilibrium, the liquid absorbs heat from its surroundings, causing the surroundings to cool down. **Factors Affecting Evaporation:** * **Surface Area:** A larger surface area allows more molecules to be at the interface, increasing the rate of evaporation. * **Temperature:** Higher temperatures increase the average kinetic energy of molecules, meaning a larger fraction of them have enough energy to escape, thus increasing the evaporation rate. * **Intermolecular Forces:** Liquids with weaker intermolecular forces (e.g., gasoline, ethyl ether) evaporate faster than those with stronger forces (e.g., water, which has strong hydrogen bonds). > **Real-World Application: Why Does High Air Humidity Make it Feel Hotter?** > Our bodies cool down through the evaporation of sweat. When the air is humid (high concentration of water vapour), the rate of evaporation from our skin slows down significantly. This reduces the body's ability to cool itself, making us feel hotter than the actual air temperature might suggest. ## 5.2.2 Vapour Pressure When a liquid is in a closed container, molecules evaporate from the surface and enter the gas phase. Some of these gas-phase molecules (vapours) will collide with the liquid surface and be recaptured in a process called **condensation**. * **Dynamic Equilibrium:** Initially, the rate of evaporation is greater than the rate of condensation. As more vapour accumulates, the rate of condensation increases until it becomes equal to the rate of evaporation. This state, where the two opposing processes occur at the same rate, is called **dynamic equilibrium**. $$ \text{Liquid} \rightleftharpoons \text{Vapour} $$ * **Definition:** **Vapour pressure** is the pressure exerted by the vapour of a liquid when it is in dynamic equilibrium with the liquid at a given temperature. **Factors Affecting Vapour Pressure:** * **Nature of the Liquid (Intermolecular Forces):** * *Stronger* intermolecular forces hold molecules in the liquid phase more tightly, making it harder for them to evaporate. This results in a **lower** vapour pressure. * *Weaker* intermolecular forces allow molecules to escape more easily, resulting in a **higher** vapour pressure. * **Temperature:** * Increasing the temperature increases the kinetic energy of the molecules, leading to a higher rate of evaporation and thus a **higher** vapour pressure. For example, the vapour pressure of water increases from $4.579 \, \text{torr}$ at $0^{\circ}C$ to $760 \, \text{torr}$ at $100^{\circ}C$. **Note:** Vapour pressure is independent of the amount of liquid, the volume of the container, or the surface area of the liquid. **Table 5.1: Vapour pressure of some liquids at $20^{\circ} \mathrm{C}$.** | Name of Compound | Vapour Pressure at $20^{\circ} \mathrm{C}$ (torr) | | :--- | :--- | | Glycerol | 0.00016 | | Mercury | 0.012 | | Water | 43.9 | | Carbon Tetrachloride | 87 | | Chloroform | 170 | | Ethyl ether | 442.2 | | Isopentane | 580 | ## 5.2.3 Boiling Point The **boiling point** is the temperature at which the vapour pressure of a liquid becomes equal to the external (atmospheric) pressure. At this temperature, bubbles of vapour form throughout the bulk of the liquid, and it rapidly converts into a gas. * **Molar Heat of Vaporization ($\Delta H_{vap}$):** The amount of heat energy required to convert one mole of a liquid into its vapour at its boiling point. For water, $\Delta H_{vap}$ is $40.6 \, \text{kJ/mol}$. This energy is used to overcome intermolecular forces, not to increase the temperature, which remains constant during boiling. **Boiling Point and External Pressure:** The boiling point of a liquid is directly dependent on the external pressure. * **High External Pressure:** More energy is required for the vapour pressure to match the external pressure, so the liquid boils at a **higher temperature**. * *Example:* A pressure cooker increases the pressure inside, raising the boiling point of water to around $120^{\circ}C$ (at 1489 torr), which cooks food faster. * **Low External Pressure:** Less energy is required, so the liquid boils at a **lower temperature**. * *Example:* Atop Mount Everest, where the atmospheric pressure is about 323 torr, water boils at only $69^{\circ}C$. In Murree (700 torr), water boils at $98^{\circ}C$. * **Vacuum Distillation:** By reducing the pressure artificially, liquids can be made to boil at much lower temperatures. This is useful for distilling heat-sensitive compounds like glycerine, which would decompose at its normal boiling point ($290^{\circ}C$) but can be safely distilled at $210^{\circ}C$ under a vacuum (50 torr). ## 5.2.4 Viscosity **Viscosity** is a measure of a liquid's resistance to flow. A liquid with high viscosity (e.g., honey) flows slowly, while one with low viscosity (e.g., water) flows easily. * **Cause:** Viscosity arises from the internal friction between adjacent layers of molecules as they slide past one another. Stronger intermolecular forces lead to greater internal friction and thus higher viscosity. * **Factors Affecting Viscosity:** * **Intermolecular Forces:** Stronger forces (like hydrogen bonding) increase viscosity. * **Molecular Size and Shape:** Large, irregular molecules (like those in honey) can become entangled, increasing viscosity. Small, compact molecules (like in acetone) lead to low viscosity. * **Temperature:** Increasing the temperature increases the kinetic energy of molecules, which helps them overcome intermolecular forces more easily. Therefore, **viscosity decreases as temperature increases**. **Units:** * **SI Unit:** Pascal-second ($Pa \cdot s$) or $kg \, m^{-1} \, s^{-1}$. * **Common Unit:** Poise (P). * **Conversion:** $1 \, Pa \cdot s = 10 \, \text{poise}$ ## 5.2.5 Surface Tension **Surface tension** is a property of a liquid that causes its surface to behave like a stretched elastic membrane. * **Cause:** Molecules in the bulk of the liquid are attracted equally in all directions by neighboring molecules. However, molecules at the surface experience a net inward pull because there are no molecules above them. This inward force pulls the surface molecules together, creating tension. * **Effect:** To minimize this tension, liquids tend to adopt a shape with the smallest possible surface area. This is why small liquid droplets are spherical. * **Factors Affecting Surface Tension:** * **Intermolecular Forces:** Stronger intermolecular forces result in a stronger inward pull on the surface molecules, leading to **higher surface tension**. Water has a high surface tension due to its strong hydrogen bonds. * **Temperature:** Increasing temperature increases the kinetic energy of molecules, weakening the intermolecular forces and thus **decreasing surface tension**. **Units:** * **SI Units:** Joules per square meter ($J \, m^{-2}$) or Newtons per meter ($N \, m^{-1}$). ## 5.2.6 Hydrogen Bonding and the Properties of Water The strong hydrogen bonds between water molecules are responsible for many of its unique and vital physical properties. 1. **High Surface Tension:** The strong hydrogen bonding in water creates a powerful net inward pull on surface molecules, resulting in an unusually high surface tension compared to other liquids. | Solvent | Surface Tension ($\gamma \times 10^{-2}$) ($N m^{-1}$) | | :--- | :--- | | **Water** | **7.275** | | Methanol | 2.26 | | Ethanol | 2.28 | | Benzene | 2.888 | | Hexane | 1.84 | | $CCl_4$ | 2.70 | 2. **Low Vapour Pressure:** The strong attractive forces make it difficult for water molecules to escape into the vapour phase, resulting in a relatively low vapour pressure for its molar mass. 3. **High Heat of Vaporization:** A large amount of energy ($40.6 \, \text{kJ/mol}$) is required to break the hydrogen bonds and convert liquid water into steam. This property is crucial for regulating temperature in biological systems and the environment. 4. **High Boiling Point:** Because of its low vapour pressure, water must be heated to a high temperature ($100^{\circ}C$ at standard pressure) for its vapour pressure to equal the atmospheric pressure. ## Possible Questions/Answers * **Q:** Why does water boil at a lower temperature on a mountain than at sea level? **A:** Atmospheric pressure is lower at higher altitudes. A liquid boils when its vapour pressure equals the external atmospheric pressure. Since the external pressure is lower on a mountain, water's vapour pressure needs to reach a lower value to initiate boiling, which occurs at a lower temperature. * **Q:** What is the relationship between intermolecular forces and vapour pressure? **A:** They are inversely related. Stronger intermolecular forces hold molecules in the liquid phase more tightly, making evaporation more difficult and thus lowering the vapour pressure. Weaker forces allow for easier evaporation and a higher vapour pressure. * **Q:** Why is honey more viscous than water at the same temperature? **A:** Honey consists of large sugar molecules with strong intermolecular forces, which cause significant internal friction and resistance to flow. Water molecules are small, and although they have hydrogen bonds, the overall resistance to flow is much lower than in honey.

5.2 Physical Properties Of Liquids | States And Phases Of Matter | Chemistry 11