Is Sublimation Endothermic Or Exothermic Process?

Sublimation is a process where there is a thermal energy modification in the origin and ending energy states of the molecules or atoms associated. Many such processes reside in physics and chemistry that associate the input or output of thermal energy. If a process like sublimation were to give off heat, it will be hailed as an exothermic process. If the process were to take in heat, it will be hailed as an endothermic process.

This modification in energy state from the start to the end states is related to the total energy situation of the molecules or atoms, so knowing whether it is an exothermic or endothermic process gives a researcher an expression of the final energy state.

Is Sublimation Endothermic Or Exothermic?

Sublimation is the changeover of a substance straight from the solid to the gas state, without passing through the liquid state. Sublimation is an endothermic process that appears at temperatures and pressures below a substance’s triple point in its phase diagram, which correlates to the lowest pressure at which the substance can occur as a liquid. The reverse process of sublimation is deposition or desublimation, in which a substance passes straight from a gas to a solid phase.

Sublimation has also been applied as a generic term to describe a solid-to-gas transition (sublimation) pursued by a gas-to-solid transition (deposition). While vaporization from liquid to gas appears as evaporation from the surface if it occurs below the boiling point of the liquid, and as boiling with the construction of bubbles in the interior of the liquid if it appears at the boiling point, there is no such characteristic for the solid-to-gas transition which always appears as sublimation from the surface.

At normal pressures, most chemical aggregates and elements possess three disparate states at distinct temperatures. In these facts, the transition from the solid to the gaseous state expects an intermediate liquid state. The pressure assigned to is the partial pressure of the substance, not the total (e.g. atmospheric) pressure of the full system. So, all solids that possess a detectable vapor pressure at a certain temperature commonly can sublime in the air (e.g. water ice just below 0 °C). For some substances, such as carbon and arsenic, sublimation is much clearer than evaporation from the melt, because the pressure of their triple point is very high, and it is challenging to obtain them as liquids.

The term sublimation indicates a physical change of state and is not used to define the transformation of a solid to a gas in a chemical reaction. For example, the disengagement on heating of solid ammonium chloride into hydrogen chloride and ammonia is not sublimation but a chemical reaction. Correspondingly, the combustion of candles, containing paraffin wax, to carbon dioxide and water vapor is not sublimation but a chemical reaction with oxygen.

Sublimation is created by the absorption of heat which provides enough energy for some molecules to overcome the attractive forces of their neighbors and escape into the vapor phase. Since the process requires additional energy, it is an endothermic change. The enthalpy of sublimation (also called heat of sublimation) can be calculated by summing the enthalpy of fusion and the enthalpy of vaporization.

Energy Change Accompanies All Phase Changes

The question then pursues, what kind of change in energy accompanies each phase change? To learn this, think about the flow of the particles in each phase. You also need to realize how attracted the molecules are to each other within the phase.

Solids incorporate particles that aren’t moving so much as compared to a liquid or a gas. They have some thermal motion, but definitely not the same amount as a liquid or a gas. Only after boosting energy (or heat) do these particles initiate to move around faster.

Think about a piece of ice. The particles of water in the piece of ice are not moving much until the water starts to melt. What allows the water to melt? Well, it’s a supplement of heat.

What about when you boil water? You need to bring the water over a flame in order to add heat to the arrangement and have the water boil in order to make water vapor.

This input of energy is also sufficient to overcome the attractive forces that hold the particles together. Water is a good example of a substance that has consequential intermolecular forces carrying it together. Water likes to stick to itself through hydrogen bonding. Thus the energy that is input must be sufficient to have the molecules stop sticking to themselves so much.

This implies that as you move from solid to liquid to gas, all accompanying phase modifications require the input of heat. Thus, these phase changes are an example of an endothermic reaction.

On the other hand, changing from gas to liquid to solid requires the opposite: Heat must be released. These phase changes are hailed as exothermic reactions.

In order to form liquid water into ice, you must put the water into a cold environment so that heat leaves the water. Only then will the water freeze.

When your hand touches steam, you feel heat because the steam immediately condenses upon touching your skin. The release of energy is felt as heat as the water vapor goes to water.