How the table is organised

To tackle this question, start by comprehending how states of matter affect density for a typical material. Density is calculated by the formula: $\text{Density}(\rho )=\frac{\text{mass}}{\text{volume}}$. Generally, in solids, particles are tightly packed, while in gases, they are spread the farthest apart. Thus, the typical density order is: ${\rho }_{\text{solid}}>{\rho }_{\text{liquid}}\gg {\rho }_{\text{gas}}$. We need to identify a transition leading to a reduction in density, moving from a denser to a less dense state.

Examine Solidification: During this process, the state changes from Liquid $\to$ Solid. For most substances, solids have a higher density than liquids. Hence, the density increases instead of decreasing, which contradicts the requirement.

Consider Liquefaction and Sublimation: - Liquefaction: Gas $\to$ Liquid. Liquids are much denser than gases, leading to a density increase. - Sublimation: Gas $\to$ Solid. Solids being denser than gases, results in a substantial density increase. These do not result in a density decrease.

Now, let's look at Evaporation: Here, the change is Liquid $\to$ Gas. As gas particles are significantly spaced compared to liquid ones, the volume experiences a major increase for the same mass. This leads to a substantial reduction in density (${\rho }_{\text{liquid}}\to {\rho }_{\text{gas}}$).

By reviewing the density change for each phase transition, only evaporation results in a density drop. The shift to a gaseous state involves the largest volume expansion, making it the transition with the most pronounced density decrease.

To understand how density alters during state changes, begin with the basic definition of density, $\rho$: $\rho =\frac{m}{V}$. With a constant mass ($m$), density varies inversely with volume ($V$). Kinetic theory states that a substance's volume is influenced by particle spacing, so a density decrease happens when particles spread apart, raising the total volume.

Kinetic theory explains particle arrangements across states: * Solids: Particles are tightly packed in a fixed lattice. * Liquids: Particles are still closely packed but can slide past each other. * Gases: Particles are widely spaced, with much empty space. A notable density drop requires a shift from tightly packed to widely spaced particles.

Exclude options where particles move closer, as these result in a density rise: * Solidification: Liquid $\to$ Solid. Particles form a fixed, packed structure. * Liquefaction: Gas $\to$ Liquid. Particles move from being far apart to being tightly packed. * Sublimation: Gas $\to$ Solid. Particles shift from being widely spaced to a fixed, dense lattice.

Evaporation (Liquid $\to$ Gas) sees particles transitioning from being closely packed to being far apart. In gases, particles are spaced approximately $10$ times more than in liquids, leading to a volume surge of about $1000$ times, resulting in the most notable density decrease.

To solve this, we must consider the Kinetic Molecular Theory, which explains how particles in matter act differently depending on their state. We are analyzing why a liquid like water can "pour" or change form, while a solid like ice remains unchangeable. This distinction is due to the varying strength of particle attractions.

If the attractive forces were absent, particles would not adhere to each other at all, similar to an 'ideal gas'. Since liquids retain a consistent volume and do not disperse like gases, we know significant forces must exist between the particles.

In a solid, the forces are strong enough to keep particles fixed in place. If liquid forces were stronger than those in a solid, the particles would be even more confined and unable to move around each other.

Attractive forces between liquid molecules exist at all times, not just at boiling point. At boiling, particles gain enough energy to break free from these forces and transition into the gas phase, but the forces are always present before that.

In liquids, there is a balance. The forces keep particles close (making liquids hard to compress like solids), but are also relaxed enough to let particles slide over each other. This fluidity lets liquids conform to the shape of their container.

By examining the properties of solids and liquids, we see that the ability of liquids to flow is due to the specific balance of forces found in Option 3.

To determine the answer, we use the particle model (also known as the kinetic theory of matter), as covered in the Matter: States and properties of matter syllabus section. We need to identify the force arrangement that lets liquids be fluid and occupy the shape of their container, while still being a condensed form.

Option 1 suggests forces are "absent". This would mean no attraction among particles, similar to an ideal gas where particles freely fill space. Since liquids stay within a container, forces must be present.

Option 2 implies liquid forces are "stronger" than those in a solid. In solids, particles are fixed in a lattice due to strong forces. If liquids had stronger forces, they would not flow as they do.

Option 4 incorrectly states forces only work at boiling point. At boiling, particles overcome the forces. Attractive forces work throughout the liquid and solid phases, only losing their grip when converting to gas.

After removing the other options, we find that the forces in liquids are strong enough to keep particles together but weak enough to allow movement, defining fluidity.

The query is about the unit of absolute temperature required by the kinetic theory of gases, which states that the average kinetic energy of gas particles is proportional to their temperature.

In thermodynamics, absolute temperature is measured on a scale beginning at absolute zero ($0\text{ K}$), the point where particles theoretically cease all movement. This is crucial for understanding gas behavior.

Option A, Joules ($J$), is a unit used for measuring energy, like kinetic energy itself, and not temperature. Thus, it is not suitable for measuring absolute temperature.

Options B (Celsius) and D (Fahrenheit) are considered relative temperature scales. They are not 'absolute' because their starting points don't represent a lack of thermal energy. For instance, $0^{\circ }\text{C}$ is the freezing point of water.

The Kelvin scale is the SI standard for thermodynamic temperature. For the kinetic theory of gases, the proportional relationship is expressed as: $\text{Average }{E}_k\propto T$ where $T$ is in Kelvin. This makes Kelvin the right choice for absolute temperature.

According to the kinetic theory of gases, the average kinetic energy ($\overline{E_k}$) of particles is proportional to the absolute temperature ($T$). This is shown in the equation: $\overline{E_k}\propto T$ When discussing the Characteristics of gases in Matter, we need a scale where zero signifies no kinetic energy.

Looking at Option A, Joules ($J$) is the unit for energy. Although the question involves kinetic energy, it requires the unit of the temperature that energy corresponds to. Measuring temperature in Joules is like timing with a ruler!

Consider Option B and Option D. Both Celsius (${}^{\circ }\text{C}$) and Fahrenheit (${}^{\circ }\text{F}$) are relative scales based on arbitrary points. In kinetic theory, when temperature is $0$, kinetic energy should also be $0$, which isn't true at $0^{\circ }\text{C}$ or $0^{\circ }\text{F}$.

The Kelvin scale is the absolute scale used in science. At $0\text{ K}$ (Absolute Zero), molecular motion theoretically stops, indicating zero kinetic energy. Thus, Kelvin is the correct unit for the relationship in the kinetic theory of gases.

According to the Kinetic Molecular Theory (KMT), the absolute temperature of a material reflects the average kinetic energy ($E_k$) of its particles.

This relationship is shown as: $T\propto \text{Average }{E}_k$

This idea is fundamental to the Matter: States and properties of matter topic. If you need a refresher on particle behavior in different states, that's a great section to revisit!

Kinetic energy is the energy due to motion. For a particle with a certain mass ($m$) and speed ($v$), its kinetic energy is described by:

$E_k=\frac{1}{2}m{v}^2$

As the mass remains unchanged, a rise in kinetic energy (due to increased temperature) leads to a rise in the particle's speed ($v$). Simply put: More heat = More kinetic energy = Faster particles.

Using the relationship we established ($T\uparrow ⟹v\uparrow$), let's review the options. Option D claims temperature doesn't affect speed, which contradicts the particle model's definition of temperature. Option C states particles slow down as temperature rises, which is the opposite of what science tells us.

Option B maintains that particles only move at temperatures above $100{}^{\circ }\text{C}$. This is inaccurate because $100{}^{\circ }\text{C}$ is only the boiling point of water at standard pressure, not a limit for particle movement. As per KMT, particles are constantly in motion at any temperature above absolute zero ($-273.15{}^{\circ }\text{C}$).

Given the direct relationship between absolute temperature and average kinetic energy, the logical conclusion is that particles move faster as the temperature increases.

The particle model (or kinetic molecular theory) explains particle behavior. A crucial part of this model is recognizing how thermal energy affects particle motion. We're identifying the assumption that connects temperature with movement.

Does temperature have no effect? Consider everyday observations: aromas spread faster when cooking, and sugar dissolves quicker in hot drinks. These are because heat affects particles. If temperature had no impact, these processes wouldn't speed up.

This choice claims that particles only move above $100{}^{\circ }\text{C}$. However, we know particles in liquid water at room temperature ($25{}^{\circ }\text{C}$) or in ice ($0{}^{\circ }\text{C}$) are still moving or vibrating. Particle movement only theoretically stops at absolute zero ($-273.15{}^{\circ }\text{C}$), which is much lower than $100{}^{\circ }\text{C}$.

If particles moved slower as temperature increased, heating substances would slow down reactions and diffusion, which isn't the case. In physics, temperature is tied to the average kinetic energy of particles. Given the kinetic energy formula:

$E_k=\frac{1}{2}m{v}^2$

An increase in temperature (and thus kinetic energy) must lead to faster speeds ($v$), not slower.

Having ruled out the other options, we're left with the essential assumption of the particle model: as temperature rises, particles gain kinetic energy and move faster.