Category: water steam and ice

I understand that water changes to its gaseous state when boiled at 212 degrees F. But why does that same water, in the form of falling rain, change to its gaseous state at a much lower temperature after it strikes a hot surface, such as a paved road? — JA, North Carolina

Liquid water can evaporate to form gaseous water (i.e., steam) at any temperature, not just at its boiling temperature of 212 F. The difference between normal evaporation and boiling is that, below water’s boiling temperature, evaporation occurs primarily at the surface of the liquid water whereas at or above water’s boiling temperature, bubbles of pure steam become stable within the liquid and water can evaporate especially rapidly into those bubbles. So boiling is a just a rapid form of evaporation.

What you are actually seeing when raindrops land on warm surfaces is tiny water droplets in the air, a mist of condensation. Those droplets happen in a couple of steps. First, the surface warms a raindrop and speeds up its evaporation. Second, a small portion of warm, especially moist air rises upward from the evaporating raindrop. Third, that portion of warm moist air cools as it encounters air well above the warmed surface. The sudden drop in temperature causes the moist air to become supersaturated with moisture—it now contains more water vapor than it can retain at equilibrium. The excess moisture condenses to form tiny water droplets that you see as a mist.

This effect is particularly noticeable when it’s raining because the humidity in the air is already very near 100%. The extra humidity added when the warmed raindrops evaporate is able to remain gaseous only in warmed air. Once that air cools back to the ambient temperature, the moisture must condense back out of it, producing the mist.

Why does ice float on water? — Y, Laguna

Solid ice is less dense than liquid water, meaning that a liter of ice has less mass (and weighs less) than a liter of water. Any object that is less dense than water will float at the surface of water, so ice floats.

That lower-density objects float on water is a consequence of Archimedes’ principle: when an object displaces a fluid, it experiences an upward buoyant force equal in amount to the weight of the displaced fluid. If you submerge a piece of ice completely in water, that piece of ice will experience an upward buoyant force that exceeds the ice’s weight because the water it displaces weighs more than the ice itself. The ice then experiences two forces: its downward weight and the upward buoyant force from the water. Since the upward force is stronger than the downward force, the ice accelerates upward. It rises to the surface of the water, bobs up and down a couple of times, and then settles at equilibrium.

At that equilibrium, the ice is displacing a mixture of water and air. Amazingly enough, that mixture weighs exactly as much as the ice itself, so the ice now experiences zero net force. That’s why its at equilibrium and why it can remain stationary. It has settled at just the right height to displace its weight in water and air.

As for why ice is less dense than water, that has to do with the crystal structure of solid ice and the more complicated structure of liquid water. Ice’s crystal structure is unusually spacious and it gives the ice crystals their surprisingly low density. Water’s structure is more compact and dense. This arrangement, with solid water less dense than liquid water, is almost unique in nature. Most solids are denser than their liquids, so that they sink in their liquids.

Why do washed clothes dry faster in open air than in a closed room? — A, Aizawl, India

What thrills me about your question is that while we’ve all noticed this effect, we’re never taught why it happens. Let me ask your question in another way: we know that opening a window makes the clothes dry faster, but how do the clothes know that the window is open? Who tells them?

The explanation is both simple and interesting: the rate at which water molecules leave the cloths doesn’t depend on whether the window is open or closed, but the rate at which water molecules return to the cloths certainly does. That return rate depends on the air’s moisture content and can range from zero in dry air to extremely fast in damp air. Air’s moisture content is usually characterized by its relative humidity, with 100% relative humidity meaning that air’s water molecules land on surfaces exactly as fast as water molecules in liquid water leave its surface. When you expose a glass of water to air at 100% relative humidity, the glass will neither lose nor gain water molecules because the rates at which water molecules leave the water and land on the water are equal. Below 100% relative humidity, the glass will gradually empty due to evaporation because leaving will outpace landing. Above 100% relative humidity, the glass will gradually fill due to condensation because landing will outpace leaving.

The same story holds true for wet clothes. The higher the air’s relative humidity, the harder it becomes for water to evaporate from the cloths. Landing is just too frequent in the humid air. At 100% relative humidity the clothes won’t dry at all, and above 100% relative humidity they’ll actually become damper with time.

When you dry clothes in a room with the window open and the relative humidity of the outdoor air is less than 100%, water molecules will leave the clothes more often than they’ll return, so the clothes will dry. But when the window is closed, the leaving water molecules will remain trapped in the room and will gradually increase the room air’s relative humidity. The drying process will slow down as the water-molecule return rate increases. When the room air’s relative humidity reaches 100%, drying will cease altogether.

Would ice in the freezer absorb the smell in the freezer? — ML, Auckland NZ

Despite the freezer’s low temperature and the motionlessness of all the frozen foods inside it, there is still plenty of microscopic motion going on. Every surface inside the freezer is active, with individual molecules landing and leaving all the time. Whenever a molecule on the surface of a piece of food manages to gather enough thermal energy from its neighbors, it will break free of the surface and zip off into the air as a vapor molecule. And whenever a vapor molecule in the air collides with the surface of another piece of food, it may stick to that surface and remain there indefinitely.

Since the freezer has a nearly airtight seal, the air it contains remains inside it for a long time. That means that the odor molecules that occasionally break free of a pungent casserole at one end of the freezer have every opportunity to land on and stick to an ice cube at the other end. With time, the ice cube acquires the scent of the casserole and becomes unappealing.

To stop this migration of molecules, you should seal each item in the freezer in its own container. That way, any molecules that leave the food’s surface will eventually return to it. Since ice cubes are normally exposed to the air in the freezer, keeping the odor molecules trapped in their own sealed containers keeps the freezer air fresh and the ice cubes odor-free.

If you have a deck that is snow covered with a very light, fluffy snow, and no one touches it, but in the next few days, from the sun, or whatever, the snow becomes “heavier” to move, does it actually weigh more? — PP

As the snow settles and becomes denser, it may feel “heavier”, but its total weight doesn’t change much. The same water molecules are simply packing themselves into a smaller space. So while each shovel-full of the dense stuff really does weigh more than a shovel-full of the light stuff, the total number of water molecules present on your deck and their associated weight is still the same.

In actually, some of the water molecules have almost certainly left via a form of solid-to-gas evaporation known technically as “sublimation.” You have seen this conversion of ice into gas when you have noticed that old ice cubes in your freezer are smaller than they used to be or when you see that the snow outside during a cold spell seems to vanish gradually without ever melting. Sublimation is also the cause of “freezer burn” for frozen foods left without proper wrapping.

My mother-in-law feels that by shaking a partially consumed bottle of carbonated beverage after re-sealing it, it will re-pressurize keeping the carbonation better than just resealing it. I believe that, since the amount of CO₂ in the beverage and the container will stay constant, that either re-sealing or re-sealing and shaking will have the same net effect when it comes to maintaining carbonation. Is she right? – JK, New Mexico

No, you are right. In the long run, the number of CO₂ molecules left in the bottle when you close it is all that matters. Those molecules will drift in and out of the liquid and gas phases until they reach equilibrium. At the equilibrium point, there will be enough molecules in the gas phase to pressurize the bottle and enough in the liquid phase to give the beverage a reasonable amount of bite.

By giving the sealed bottle a shake, your mother-in-law is simply speeding up the approach to equilibrium. She is helping the CO₂ molecules leave the beverage and enter the gas phase. The bottle then pressurizes faster, but at the expense of dissolved molecules in the beverage itself. If there is any chance that you’ll drink more before equilibrium has been reached, you do best not to shake the bottle. That way, the equilibration process will be delayed as much as possible and you may still be able to drink a few more of those CO₂ molecules rather than breathing them.

Incidentally, shaking a new bottle of soda just before you open it also speeds up the equilibration process. For an open bottle, equilibrium is reached when essentially all the CO₂ molecules have left and are in the gas phase (since the gas phase extends over the whole atmosphere). That’s not what you want at all. Instead, you try not to shake the beverage so that it stays away from equilibrium (and flatness) as long as possible. For most opened beverages, equilibrium is not a tasty situation.

Does ice melt faster in air or in water? — BP

Ice will melt fastest in whatever delivers heat to it fastest. In general that will be water because water conducts heat and carries heat better than air. But extremely hot air, such as that from a torch, will beat out very cold water, such as ice water, in melting the ice.

I always thought that pure water cannot exceed 100° Celsius at atmospheric pressure without first turning into its gaseous state. How is it that the water heated in the microwave oven can superheat and exceed 100° Celsius? — AC

The relative stabilities of liquid and gaseous water depend on both temperature and pressure. To understand this, consider what is going on at the surface of a glass of water. Water molecules in the liquid water are leaving the water’s surface to become gas above it and water molecules in the gas are landing and joining the liquid water below. It’s like a busy airport, with lots of take-offs and landings. If the glass of water is sitting in an enclosed space, the arrangement will eventually reach equilibrium—the point at which there is no net transfer of molecules between the liquid in the glass and the gas above it. In that case, there will be enough water molecules in the gas to ensure that they land as often as they leave.

The leaving rate (the rate at which molecules break free from the liquid water) depends on the temperature. The hotter the water is, the more frequently water molecules will be able to break away from their buddies and float off into the gas. The landing rate (the rate at which molecules land on the water’s surface and stick) depends on the density of molecules in the gas. The more dense the water vapor, the more frequently water molecules will bump into the liquid’s surface and land.

As you raise the temperature of the water in your glass, the leaving rate increases and the equilibrium shifts toward higher vapor density and less liquid water. By the time you reach 100° Celsius, the equilibrium vapor pressure is atmospheric pressure, which is why water tends to boil at this temperature (it can form and sustain steam bubbles). Above this temperature the equilibrium vapor pressure exceeds atmospheric pressure. The liquid water and the gas above it can reach equilibrium, but only if you allow the pressure in your enclosed system to exceed atmospheric pressure. However, if you open up your enclosed system, the water vapor will spread out into the atmosphere as a whole and there will be a never-ending stream of gaseous water molecules leaving the glass. Above 100° C, liquid water can’t exist in equilibrium with atmospheric pressure gas, even if that gas is pure water vapor.

So how can you superheat water? Don’t wait for equilibrium! The road to equilibrium may be slow; it may take minutes or hours for the liquid water to evaporate away to nothing. In the meantime, the system will be out of equilibrium, but that’s ok. It happens all the time: a snowman can’t exist in equilibrium on a hot summer day, but that doesn’t mean that you can’t have a snowman at the beach… for a while. Superheated water isn’t in equilibrium and, if you’re patient, something will change. But in the short run, you can have strange arrangements like this without any problem.