Why does salt melt ice? Because dissolved salt lowers the temperature at which liquid water and solid ice can remain in equilibrium. When salt reaches a thin film of liquid water on an icy surface, it dissolves and forms brine. That salty water has a lower freezing point than pure water, so at temperatures where pure water would freeze, some of the ice can melt instead. The effect is called freezing point depression.
This is the chemistry behind road salt, salted ice baths and many winter de-icing systems. It also explains why salt does not work equally well at every temperature. Salt does not generate unlimited heat and it does not magically “burn” through ice. It changes the balance between freezing and melting by changing the composition of the liquid water touching the ice. If the temperature becomes low enough, the particular salt solution can no longer stay liquid and the de-icing effect weakens or stops.
Understanding why salt melts ice connects everyday life with thermodynamics, solutions, phase changes and molecular behaviour. It helps explain why roads are treated before a freeze, why crushed ice works faster than one solid block, why calcium chloride can remain useful at lower temperatures than ordinary sodium chloride, why salt helps make ice cream, and why adding salt to snow can create liquid water even when the air temperature is below 0°C.
The short answer: salt lowers water’s freezing point
Pure water freezes at about 0°C under ordinary atmospheric pressure. But water containing dissolved particles can remain liquid below that temperature. When salt dissolves, it separates into ions. Those dissolved particles make it harder for water molecules to organise themselves into the ordered crystal structure of ice.
The result is a new equilibrium. The salty solution must be cooled to a lower temperature before freezing becomes favourable.
On an icy pavement, a small amount of liquid water is normally present at the surface even when the bulk is solid. Salt dissolves into that water. The resulting brine has a lower freezing point, so more ice melts to dilute the solution until the system reaches a new balance.
That is the essential mechanism.
Ice and water are not completely separate at the surface
An ice cube may look completely solid, but its surface is dynamic. Water molecules are constantly leaving and joining the crystal. Near the melting point, a microscopically thin liquid-like layer can exist at the surface.
This matters because salt needs liquid water in order to dissolve efficiently.
If dry crystals of sodium chloride are scattered onto ice near 0°C, they encounter enough surface moisture to begin dissolving. Once a small amount of brine forms, the brine can contact more ice. Because its freezing point is below the current temperature, additional ice can melt into it.
The process feeds itself until the salt concentration becomes too diluted, the temperature drops too far, or the available salt is exhausted.
What freezing really means
Freezing is not simply “water gets cold and becomes solid.” It is a phase transition.
In liquid water, molecules move and rearrange continuously. In ice, the molecules form a much more ordered crystal structure held together by hydrogen-bonding patterns.
At the freezing point, the liquid and solid phases can coexist in equilibrium. Molecules leave the ice and join the liquid while others join the ice. At equilibrium, there is no net change in the amount of each phase.
Salt shifts that balance.
When ions are mixed into the liquid, the chemical potential of the water in the liquid phase changes. The liquid becomes stable to a lower temperature than pure water. The solid ice is still nearly pure water, so the balance between the two phases moves.
The everyday observation—ice melts—is the visible result of that thermodynamic change.
Why dissolved particles interfere with freezing
It is tempting to say that salt “gets between” water molecules and physically blocks them from freezing. That picture is useful only as a rough starting point.
The deeper explanation is statistical and thermodynamic. A solution has more possible arrangements than pure water because it contains both water molecules and dissolved ions. Forming pure ice removes water molecules from the solution but leaves the salt behind, changing concentration and free energy.
As a result, a lower temperature is needed before the ordered solid phase becomes sufficiently favourable.
This effect is called a colligative property because, for sufficiently dilute ideal solutions, the magnitude depends mainly on the number of dissolved particles rather than their chemical identity.
Different salts still behave differently in real applications because they dissociate into different numbers of ions, dissolve to different extents, interact differently with water and have different practical temperature ranges.
Why sodium chloride splits into ions
Table salt is sodium chloride, NaCl.
In the solid crystal, positively charged sodium ions and negatively charged chloride ions are arranged in a repeating ionic lattice. When the crystal contacts water, polar water molecules interact with those ions and can pull them away from the lattice.
The dissolved salt is therefore not present mainly as intact little NaCl molecules. It exists largely as hydrated sodium ions and chloride ions dispersed through the water.
One formula unit of sodium chloride can produce two dissolved ionic particles. That helps make salts effective at changing colligative properties such as freezing point.
The real solution is not perfectly ideal, especially at high concentrations, but the ion picture explains why a relatively small mass of salt can have a substantial effect.
Why brine forms before all the ice melts
Salt does not normally turn an icy road instantly into pure liquid water.
It produces brine: a concentrated salt solution.
At first, a few salt crystals dissolve in the available surface moisture. The brine can remain liquid below 0°C, so it encourages more ice to melt. That additional meltwater dilutes the brine. If the solution becomes too dilute for the existing temperature, some water may freeze again unless more salt dissolves.
The system therefore moves toward a concentration-temperature balance.
On a road, the practical goal is often not to melt every gram of snow and ice. It is to break the bond between ice and pavement, prevent new ice from forming, and create conditions in which ploughs and tyres can clear the surface more effectively.
Why salt works better near 0°C than in extreme cold
Freezing point depression has limits.
A sodium chloride–water mixture has a lowest temperature at which a liquid solution can coexist with solid phases at the eutectic composition. For ordinary sodium chloride brine, that limiting temperature is roughly -21°C.
Long before that theoretical limit, road salt can become slow and inefficient because dissolution, melting and heat transfer are all affected by the cold.
Near the freezing point, a modest amount of salt can create enough brine to work effectively. At much lower temperatures, more concentrated brine is required and the process slows.
This is why winter-maintenance agencies may use other chlorides or alternative strategies during severe cold.
Why calcium chloride can work at lower temperatures
Calcium chloride is another common de-icing salt.
When calcium chloride dissolves, one formula unit can produce one calcium ion and two chloride ions. It can create a strong freezing-point-depression effect, and its solutions can remain liquid at lower temperatures than sodium chloride brine over practical operating ranges.
Calcium chloride dissolution is also strongly exothermic: it releases heat as it dissolves. That heat can help start melting under cold conditions.
However, calcium chloride is not automatically the best choice everywhere. It can cost more, attract moisture strongly, affect materials and vegetation, and still contributes chloride to the environment.
De-icing is an engineering decision involving temperature, cost, road conditions, environmental impact and available equipment.
Why magnesium chloride is also used
Magnesium chloride is another chloride salt used in some de-icing programmes.
Like calcium chloride, it can form concentrated brines and operate at lower temperatures than sodium chloride in some conditions. It may be applied as a solid or liquid brine.
Each de-icer has a performance envelope. Engineers consider not only how low the freezing point can be pushed but also how quickly the chemical works, how much must be applied, how it affects surfaces, how easily it is stored and distributed, and what happens after the meltwater enters soil and waterways.
The chemistry is only one part of the operational choice.
Why road crews spread salt before a storm
Pre-treatment can be more effective than waiting for a thick layer of ice to form.
If brine is applied to a road before freezing precipitation arrives, it can make it harder for ice to bond strongly to the pavement.
This process is called anti-icing.
A thin brine layer lowers the local freezing point at the road surface. Snow or freezing water may still accumulate, but the bond can be weaker and subsequent mechanical removal easier.
Timing matters. Rain can wash brine away. Very low temperatures can reduce effectiveness. Traffic can redistribute material.
Winter road treatment is therefore based on forecasts, pavement temperature, precipitation type and expected traffic, not simply air temperature.
Why pavement temperature matters more than the number on a weather app
The surface being treated has its own temperature.
Road pavement can be warmer or colder than the surrounding air depending on sunlight, cloud cover, wind, ground heat, time of day and material properties.
A de-icer acts where the brine contacts the pavement and ice. The relevant question is therefore whether that local system is within the chemical’s effective range.
This is why professional winter-maintenance operations monitor road-surface conditions directly where possible.
The same idea applies at home. A shaded step, metal surface and dark paved driveway can freeze differently even under the same reported air temperature.
Why salt does not melt ice by making it hotter
One common misconception is that salt melts ice because the salt itself is warm or because the reaction always produces heat.
For ordinary sodium chloride, the main de-icing mechanism is freezing point depression, not a large release of heat.
Melting ice actually requires energy. Ice absorbs latent heat as it changes into liquid water. That energy comes from the surroundings: the pavement, air, sunlight, water and other nearby material.
Some salts, especially calcium chloride, release heat when they dissolve, which can assist the process. Others may absorb heat during dissolution.
The unifying principle remains that the dissolved ions lower the equilibrium freezing temperature of the liquid phase.
Why melting ice can make the mixture colder
This effect is familiar from making ice cream in an ice-and-salt bath.
When salt causes some ice to melt, the melting process absorbs latent heat. If the system is reasonably insulated, that energy can come from the surrounding mixture, lowering its temperature.
The salt allows liquid water to exist below 0°C, while the melting ice absorbs heat. Together, those effects can produce a bath significantly colder than ordinary melting ice.
That cold bath draws heat from the ice-cream mixture, causing the water in the food to freeze while the mixture is stirred.
The same chemistry that clears a path can therefore be used to create a colder refrigeration bath.
Why salt and ice are used in traditional ice-cream making
Pure melting ice tends to remain near 0°C while both solid and liquid water coexist at ordinary pressure. That is not cold enough to freeze an ice-cream mixture rapidly and controllably.
Adding salt creates brine that can remain liquid below 0°C. As ice continues melting, the bath becomes colder.
The food mixture is kept separate from the salty brine but transfers heat to it through the container wall.
Stirring the food mixture disrupts the growth of large ice crystals, helping produce a smoother texture.
What looks like a kitchen trick is a practical demonstration of phase equilibrium, heat transfer and colligative properties.
Why adding more salt does not help forever
If some salt is useful, it may seem that unlimited salt should melt unlimited ice at any temperature.
That is not how the phase diagram works.
There is a maximum useful concentration range for a given salt-water system. At the eutectic point, the solution reaches the lowest temperature at which liquid brine can exist in equilibrium with the relevant solids.
Below that temperature, adding more of the same salt cannot maintain a liquid solution indefinitely.
Excess salt can also remain undissolved if there is not enough liquid water available.
Good de-icing therefore uses enough chemical for the conditions, not the largest possible pile.
Why crushed salt and liquid brine can act faster
Reaction and dissolution rates depend on contact.
Smaller salt particles provide more surface area relative to their mass. They can contact thin films of water at more points and begin dissolving quickly.
Liquid brine is already dissolved. It can spread across a surface immediately instead of waiting for solid crystals to find enough moisture.
This is one reason pre-wetted salt and brine spraying are common winter-maintenance techniques.
However, a very fine material can also be scattered or diluted easily. Application method is part of the engineering problem.
Why sand does not melt ice
Sand is often spread on icy surfaces, but it works differently from salt.
Ordinary sand does not dissolve into enough particles to create a useful freezing point depression.
Its purpose is mechanical.
Sand can increase friction between tyres or shoes and an icy surface. It may improve traction even when the ice remains frozen.
In very cold conditions where salt is less effective, abrasives can be valuable.
De-icing and traction are different objectives. Salt changes phase behaviour. Sand changes surface grip.
Why sugar can also lower the freezing point
Freezing point depression is not unique to salt.
Sugar dissolves in water and also lowers its freezing point. This is important in frozen desserts, where dissolved sugar helps control how much water freezes and influences texture.
But sugar is not a practical road de-icer. It is more expensive, sticky, biologically active and less efficient per unit mass than common ionic salts for this purpose.
The comparison is useful because it shows that the mechanism is a property of solutions rather than a mysterious special power belonging only to table salt.
Why seawater freezes below 0°C
Seawater contains dissolved salts, so its freezing point is lower than that of pure freshwater.
Typical ocean water begins freezing around -1.8°C, although the exact value depends on salinity and pressure.
As sea ice forms, much of the salt is excluded from the growing ice crystals. Concentrated brine remains in channels and pockets and can drain into the water below.
This brine rejection increases the salinity and density of nearby seawater and contributes to important ocean processes.
The same basic principle—dissolved solute lowering freezing point—operates at planetary scale.
Why salt can damage plants
Road salt does not vanish after winter.
Chloride and sodium ions can move into soil and water. High salt concentrations can make it harder for plant roots to take up water and can disturb nutrient balance.
Salt spray from roads can also contact leaves and stems directly.
The severity depends on concentration, species, drainage, soil properties and repeated exposure.
This environmental cost is one reason road agencies try to improve application efficiency. Using the right amount at the right time can reduce both expense and unnecessary chloride loading.
Why salt can accelerate corrosion
Salt water conducts electricity better than pure water because it contains mobile ions.
That can accelerate electrochemical corrosion processes on metals such as steel, especially where protective coatings are damaged.
Vehicles, bridges and other infrastructure exposed repeatedly to chloride-containing spray therefore require corrosion management.
Designers may use coatings, drainage, corrosion-resistant materials, inspection and washing to reduce damage.
This is an important reminder that a material can solve one engineering problem while creating another. De-icing improves winter safety, but chloride exposure has costs that must be managed.
Why salt can damage concrete indirectly
Concrete is porous. Water and dissolved salts can enter cracks and pores.
Repeated freezing and thawing can stress concrete because water expands when it freezes. De-icing chemicals can increase the number of freeze-thaw cycles around the surface by keeping some water liquid at lower temperatures and then allowing refreezing when conditions change.
Chlorides can also reach reinforcing steel inside concrete and contribute to corrosion. Rust products occupy more volume than the original steel and can crack or spall surrounding concrete.
The exact deterioration process depends on concrete quality, air entrainment, drainage, exposure and salt type.
De-icing chemistry therefore belongs to civil engineering as much as to classroom chemistry.
Why salt works differently on fresh snow and compacted ice
Fresh snow contains many air spaces and a large surface area. Compacted ice is denser and may be strongly bonded to the pavement.
Salt can penetrate and form brine more easily in some loose snow conditions, while a thick ice layer can require time for brine channels to develop.
Mechanical removal is therefore important.
Ploughing removes bulk snow so that chemical treatment does not have to melt an enormous mass. De-icer can then work on the thinner residual layer and the bond at the pavement surface.
Trying to dissolve an entire deep snowfall with salt alone would be slow, expensive and environmentally wasteful.
Why sunlight can make salted ice disappear faster
Melting requires energy.
Sunlight can warm pavement and ice, providing some of the latent heat required for melting. Dark surfaces often absorb more solar energy than bright snow.
If salt has already lowered the freezing point, solar heating can push the system further into the liquid region.
This is why the same amount of salt can seem much more effective on a bright afternoon than during a cold, cloudy night.
De-icing performance is a combined result of chemistry and heat flow.
Why salt can make refreezing dangerous later
Salted ice can melt into brine during warmer conditions, but the solution can later refreeze if the temperature falls below its current freezing point.
This can happen after dilution by additional snow, rain or meltwater.
A surface that looked wet can become icy again as temperature changes.
The presence of salt lowers the freezing point; it does not make refreezing impossible.
This is why winter safety depends on repeated monitoring rather than a one-time application.
A simple classroom experiment
A safe demonstration can compare two cups containing similar amounts of crushed ice.
Leave one cup untreated and mix salt into the other. Measure temperature with suitable thermometers and observe the amount of liquid water over time.
The salted mixture can fall below 0°C while containing liquid brine. That single observation challenges the common idea that liquid water cannot exist below the normal freezing point.
Students can then vary salt mass, ice particle size or starting temperature while keeping other variables controlled.
The experiment becomes a lesson not only in freezing point depression but also in fair testing, measurement uncertainty and energy transfer.
What the freezing-point equation tells us
For dilute solutions, freezing point depression can be approximated by a relationship in which the temperature change depends on the solvent’s cryoscopic constant, the concentration of dissolved particles and a factor describing how many particles the solute produces.
The equation is often written as ΔTf = iKfm.
Here, ΔTf is the decrease in freezing point, Kf is a property of the solvent, m is molality and i is the van ’t Hoff factor describing the effective number of dissolved particles.
The simple equation works best for dilute solutions. Road brines can be concentrated enough that real behaviour deviates from ideal assumptions, so engineers use measured phase data rather than relying only on the introductory formula.
Still, the equation captures the central principle: more dissolved particles generally push the freezing point lower.
Why the word “melt” can be slightly misleading
When people say “salt melts ice,” the phrase sounds as if salt directly supplies the energy required for melting.
A more precise description is that salt lowers the equilibrium freezing temperature, allowing some ice to melt under conditions where pure water would otherwise remain frozen.
The energy for the phase change still has to come from somewhere.
This distinction matters because it explains why de-icing speed depends on the thermal environment. A chemical can make melting thermodynamically favourable but the process can still be slow if little heat is available.
Common myths about salt and ice
Myth: salt always melts ice instantly
Melting takes time because the salt must dissolve, brine must spread and heat must enter the melting system.
Myth: salt works at any temperature
Every salt-water system has practical and thermodynamic temperature limits.
Myth: salt melts ice because it is hot
The main mechanism is freezing point depression. Some de-icers release heat when dissolving, but that is not the universal explanation.
Myth: more salt is always better
Excess material can remain undissolved and increases environmental and infrastructure costs without proportional benefit.
Myth: sand and salt do the same job
Salt changes freezing behaviour. Sand mainly improves traction.
Frequently asked questions
Does table salt work on ice?
Yes. Ordinary sodium chloride can lower the freezing point and form brine, although road-grade salt is often cheaper and more practical for large surfaces.
Why does salt make ice colder?
Salt allows ice to melt below 0°C. Melting absorbs latent heat, which can cool the remaining mixture if that energy comes from the surrounding ice-brine system.
Can salt prevent water from freezing completely?
No. Salt lowers the freezing point, but sufficiently cold conditions will still freeze the solution or produce a mixture of solid phases.
Why is salt spread on roads before snow?
Pre-applied brine can reduce the strength of the bond between new ice and pavement, making later removal easier.
Does salt make ice safer immediately?
Not always. A partially melted surface can remain slippery, and brine can refreeze if temperature falls or the solution becomes diluted.
Why is salt used around ice-cream containers?
The salt lowers the freezing point of the ice-water bath so it can become colder than 0°C and draw heat from the ice-cream mixture.
The deeper answer: salt changes the phase boundary
The deepest answer to why salt melts ice is not that salt attacks the ice crystal or produces a mysterious warming reaction. Salt changes the conditions under which liquid water is stable.
Pure water near 0°C sits close to the boundary between liquid and solid. Add dissolved ions, and the liquid solution becomes thermodynamically favourable over a wider temperature range. Ice can then melt into that solution until composition, temperature and available energy establish a new equilibrium.
That single principle explains a remarkable range of observations: salty roads staying wet below freezing, seawater freezing below 0°C, ice-and-salt baths becoming colder, different de-icers having different temperature ranges, and excess salt eventually reaching diminishing returns.
The visible event is simple: a crystal of ice turns into liquid.
The system underneath is a balance of molecular order, dissolved particles, chemical potential, temperature and heat transfer.
Salt melts ice because it changes that balance.
Why the phase diagram is the clearest map of salt and ice
A phase diagram turns the salt-and-ice story into a map. One axis represents temperature and another represents salt concentration. Different regions show where liquid brine, ice, solid salt-containing phases or mixtures are stable. Instead of asking only whether the temperature is above or below 0°C, the diagram asks a better question: at this exact concentration and temperature, what phases can exist together?
Start with pure water at 0°C. Move horizontally by adding salt and the liquid can remain stable at lower temperatures. Continue increasing concentration and the freezing boundary drops until the eutectic point is reached. Beyond that special composition, adding still more salt no longer keeps pushing the liquid boundary down in the same simple way.
This map explains why road crews care about both temperature and concentration. A weak brine may be suitable for a mild frost but not for a deep freeze. A strong brine may remain liquid lower, but concentration, cost and material effects still matter. The useful state of the road surface is therefore determined by where the real mixture sits on the phase diagram.
Why rock salt and table salt follow the same chemistry
Rock salt and ordinary table salt are both based mainly on sodium chloride. Their de-icing chemistry is therefore fundamentally the same: sodium and chloride ions dissolve into water and lower the solution’s freezing point.
The practical products differ in purity, crystal size, additives and cost. Food-grade table salt is refined for consumption. Road salt is manufactured and graded for large-scale winter maintenance, where perfect culinary purity would add expense without improving the main freezing-point effect.
Particle size also matters. Very large crystals may persist longer and provide a reservoir of salt as moisture develops. Finer particles dissolve faster but can be scattered or diluted more readily. Operators choose a material and application method suited to the expected weather and road condition.
Why water is necessary for dry salt to begin working
Salt must enter solution before it can lower the freezing point of water effectively. A completely dry salt crystal sitting on a completely dry cold surface cannot create brine by itself.
In real winter conditions, however, ice surfaces contain liquid-like layers and snow often contains enough moisture for dissolution to begin. Road agencies may also pre-wet solid salt with brine so that the chemical starts acting sooner and adheres better to the pavement instead of bouncing away under traffic.
This requirement explains why timing matters. Salt applied to a moist surface just before freezing can begin distributing through a thin film. The same dry crystals scattered on extremely cold, dry, hard ice may act much more slowly.
Why temperature and concentration must be considered together
It is easy to discuss de-icing as if there were one universal “freezing point of salt water.” There is not. The freezing temperature depends on how much salt is dissolved.
A lightly salted solution may freeze only a little below 0°C. A more concentrated sodium chloride solution can remain liquid much lower. But the relationship is not unlimited and not perfectly linear across all concentrations.
Now consider what happens when salty ice melts. The newly melted water dilutes the brine. Dilution raises its freezing point again. If the surrounding temperature remains low, that changing concentration can push the solution back toward refreezing. More salt may then need to dissolve to maintain the liquid state.
The system is continually adjusting. That is why a pile of salt on a road is not a one-time chemical switch. It participates in a changing mixture of ice, water, dissolved ions and heat.
Why bridges can freeze before nearby roads
Bridges and elevated road decks are exposed to cold air above and below. Ordinary ground-level roads exchange heat with the soil beneath them, which can sometimes keep the pavement warmer for longer during a cooling event.
Because a bridge can lose heat from more surfaces, its pavement temperature may fall below freezing sooner even when the nearby road remains merely wet. This is why warning signs about bridges freezing first are based on a real heat-transfer difference.
For de-icing, the lesson is important: air temperature alone cannot describe risk. Surface construction, exposure and stored heat affect whether water freezes and how quickly a salt treatment can work.
Why de-icing is partly a transport problem inside the ice
Once brine forms, it has to reach the interface between ice and the surface underneath. Channels and cracks help the liquid move. Mechanical stress from tyres, footsteps and ploughing can break the ice and expose new surfaces.
This is why the physical condition of the ice matters. A thin porous frost can respond quickly. A thick polished layer can take longer because the chemical has less access to the lower boundary and more total ice must be affected.
Chemistry sets the equilibrium, but transport controls how rapidly the system approaches that equilibrium. Dissolution, diffusion, liquid flow and heat transfer all influence the visible speed of melting.
Why salt use is an optimisation problem, not a maximum-dose problem
Professional winter maintenance tries to achieve a safe road condition with the smallest effective material use. Too little salt may fail to prevent bonding or create enough brine. Too much wastes money, increases chloride runoff and can accelerate corrosion without producing a proportional safety gain.
The correct application depends on pavement temperature, expected precipitation, traffic, residual salt already on the surface, treatment timing and the chemical being used. Forecasting therefore becomes part of chemistry. A road crew is not merely spreading a substance; it is managing a changing physical system before and during a weather event.
The same principle applies on a smaller scale at home. A thin icy step does not require an enormous mound of salt. Mechanical removal, a modest appropriate treatment and later clearing of slush may be more effective than repeated heavy application.
Why alternative de-icers exist
Chloride salts are effective and relatively inexpensive, but they are not the only substances capable of lowering water’s freezing point. Acetate- and formate-based de-icers, urea in some specialised uses, and proprietary mixtures may be chosen where corrosion, environmental sensitivity or operating temperature changes the trade-off.
No alternative is impact-free. A chemical that is less corrosive may cost more. An organic de-icer may increase biological oxygen demand in receiving water. A product may work well on an airport surface but be too expensive for an entire highway network.
The existence of alternatives reinforces the main scientific idea: de-icing is not about a unique magical property of sodium chloride. It is about changing solution properties while balancing engineering constraints.
Why the lesson matters beyond winter roads
Freezing point depression is one example of a larger principle: mixing substances can change the physical properties of a liquid. Similar solution effects appear in boiling-point elevation, osmotic pressure and vapour-pressure lowering.
That makes the salt-and-ice question a useful bridge between everyday observation and formal chemistry. A student can begin with a slippery footpath and end up thinking about ions, entropy, phase diagrams, latent heat, concentration and environmental engineering.
The best general-knowledge questions often work this way. They start with something ordinary and reveal an entire system underneath. Salt melts ice because dissolved particles change the equilibrium between solid water and liquid solution—but following that sentence carefully opens a much larger story about how matter changes phase and how engineers use those changes in the real world.