The transformation of liquid water into solid ice is a common physical process that requires the transfer of energy to occur. The change in state from a liquid to a solid involves a thermodynamic exchange. This energy transfer must happen for the free-flowing water molecules to settle into the structured lattice of ice. Understanding this act of freezing requires a look at how energy moves between the water and its surroundings.
Defining Exothermic and Endothermic Energy Flow
The terms used to describe energy flow in physical and chemical changes are defined by observing the “system” (the water) and the “surroundings” (everything else, like the freezer). Processes that release energy from the system into the surroundings are classified as exothermic. This release often manifests as heat, which is why an exothermic reaction, like burning wood, makes the surroundings feel hotter.
Conversely, an endothermic process is one where the system absorbs energy from its surroundings. This absorption of heat causes the temperature of the surroundings to drop, which is why an endothermic reaction feels cold to the touch. For example, an instant cold pack utilizes an endothermic dissolution reaction to rapidly pull heat from an injured area.
The Answer: Freezing Water is Exothermic
Making ice cubes by freezing liquid water is categorized as an exothermic process. The water, which is the system, must actively release heat energy to its surroundings for the phase change to occur. The freezer’s cooling mechanism is designed to absorb this released heat and pump it away, allowing the water to solidify.
Molecular Energy Exchange During Phase Transitions
The exothermic nature of freezing is explained by the change in the water molecules’ energy state. In the liquid phase, water molecules possess a higher level of potential energy and are in constant, random motion. As the liquid cools, the molecules slow down and begin to align themselves to form the rigid crystal structure of ice.
This transition to a more ordered and stable solid state involves the formation of stronger, more numerous hydrogen bonds. The creation of these new bonds inherently releases energy. This excess potential energy must be expelled as heat into the surroundings, allowing the molecules to lock into the lower-energy solid lattice.
This specific quantity of energy released during the liquid-to-solid transition is known as the latent heat of fusion. For water, this value is approximately 334 kilojoules of energy released for every kilogram that freezes at 0°C. This heat is released even though the temperature remains constant at the freezing point until solidification is complete.
The Reverse Process: Why Melting Requires Energy
In contrast to freezing, the reverse process of melting ice into liquid water is endothermic. When ice melts, it must absorb energy from its surroundings to break the stable hydrogen bonds holding the crystalline structure. By absorbing heat, the molecules gain enough energy to move freely in the higher-energy liquid state.
This required energy is equal in magnitude to the latent heat of fusion released during freezing. This absorption of heat is why melting ice cools down a drink, as the ice draws thermal energy from its surroundings.