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Toothpaste is what is called a non-newtonian fluid, more specifically toothpaste is a Bingham plastic. This means that the viscosity of the fluid is linearly dependent on the shear stress, but with an offset called the yield stress (see figure below). This yield stress is what makes it hard to say whether it is liquid or solid. The fact that toothpaste is ...

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I recently got a chance to attend a talk by someone who was working on developing analytical instrumentation on Mars. The interesting story is that the initial results by ion-selective electrode was that Mars soil is full of nitrates. Nobody knew on Earth that the nitrate ion selective electrode is far more responsive to perchlorate than nitrate. After ...

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According to Martin Chaplin's Water Dissociation and pH: In ice, where the local hydrogen bonding rarely breaks to separate the constantly forming and re-associating ions, the dissociation constant is much lower (for example at $-4~\mathrm{^\circ C}$, $K_\mathrm{w} = 2 \times 10^{-20}~\mathrm{mol^2~L^{-2}}$). So $[\ce{H+}] = 1.4 \times 10^{-10}~\mathrm{... 46 No such liquid, safe or otherwise, can exist. Evaporation is a strictly endothermic process in all cases. The change in state from liquid to gas is marked by the individual particles gaining enough translational kinetic energy to overcome the mutual attractions present in the liquid phase to "fly free" in the gas phase. It is logically inconsistent for a ... 43 Here's a genchem-level answer for a genchem-level question about the classification of matter: Toothpaste is a sol: a stable suspension of tiny solid particles in a liquid. When the toothpaste dries out you can see what the solid part alone looks like. Mixtures with more than one phase often have interesting properties and behaviors that the components ... 35 First, I think I should make it clear that when water boils, the bonds in the water molecule linking the hydrogen and oxygen atom are not broken. During boiling, the intermolecular bonds in water are the ones that get broken, that is the bonds that link the water molecules together. At room temperature, there is evaporation (I wouldn't call it excitation). ... 35 As your small percentage of molecules with high enough kinetic energy evaporates, the remaining liquid water cools down. But in doing so, it drains heat from its surroundings and thus stays at room temperature (or close to it), so there is still some fraction of molecules that can evaporate, and they do so, and more heat is transferred from the surroundings, ... 33 Interesting question! A few things first: As the ice melts, it cools the water around it. Technically, the ice cube melts because the water cools down. This may sound ridiculous at first, but you must consider the fact that the ice melts because it has drawn "heat" (energy) from its surroundings. The "surroundings" being the air and water that surround ... 30 When you add salt to an ice cube, you end up with an ice cube whose temperature is above its melting point. This ice cube will do what any ice cube above its melting point will do: it will melt. As it melts, it cools down, since energy is being used to break bonds in the solid state. (Note that the above point can be confusing if you're new to thinking ... 28 Good question. Let's assume the container is infinitely strong, non-deformable, and constant in volume. Let's also assume that cooling the water is an equilibrium process -- that way, we won't have any supercooling. At equilibrium, the first tiny bit of ice that freezes will take up more volume than the water it froze from. This will raise the pressure ... 25 As you said, the meaning is exactly the same. Molten reduces the ambiguity, because you emphasize that you know, that it is solid substance at laboratory conditions and you heat it to become liquid (while staying pure substance). As an example, the "electrolysis of liquid sodium chloride" is in principle enough to tell you all about the process, but as we ... 24 Essentially, because the carbon dioxide sublimates from solid (dry ice) to gas at a very low temperature (roughly −78 °C at 1 atm), it causes water vapour in the air to condense, causing a visible fog. Thus what you are seeing is not carbon dioxide, but rather water. When we exhale and it is reasonably warm, the carbon dioxide expelled is roughly body ... 24$\ce{H2}$cannot be liquified at room temperature, whatever the pressure. Generally speaking, all gases can only be liquified when the temperature is under its critical value. 22 Diamond (carbon) does not melt at 1 atm. It sublimes to vapor. Using carbon's theoretical phase diagram below (from Wikimedia), "liquid diamond" could be achieved at about 10 GPa (99 thousand atmospheres) and 5000 K (4700 °C). Edit: In fact, heating up diamonds at 1 atm turns them into graphite first. Only further heating would cause sublimation, as ... 22 Let’s consider the following cases: getting$1\,\mathrm{mol}$of$100\,\mathrm{°C}$water on one’s skin getting$1\,\mathrm{mol}$of$100\,\mathrm{°C}$air on one’s skin getting$1\,\mathrm{mol}$of$100\,\mathrm{°C}$water vapour on one’s skin With the slightly irrealistic assumption that all of these liberate all their thermal energy to the skin while ... 22 Hydrogen critical temperature is$\pu{32.938 K, resp. -240.21 ^{\circ}C}$. Above this temperature, it cannot be liquified. So to answer your question, you can get as high pressure as you can produce and the container can withstand, as there is no condensation reducing the pressure. WARNING: An accidental explosive container rupture can easily cause severe ... 21 This is actually an active area of research for water clusters. In principle, for$\ce{(H2O)_{n}}$there should be a "melting" phase transition, much like for ice to liquid water. So, in principle, if you had an accurate enough theoretical method, you could do molecular dynamics and see when the melting point of the cluster matches bulk water. ... 21 Here's the boring answer: Toothpaste is a mixture of some solids and some liquids. The question "is it solid or liquid?" makes sense when you're talking about a substance or a mixture with a single phase—that is, a substance or mixture that's pretty much completely uniform throughout space. Examples of single-phase materials include pure water (a liquid), a ... 21 When the temperature of helium gas is decreased to about 5.2 K, a phase transition to ordinary liquid helium ($\ce{He}$I) occurs. The behavior of this liquid phase is normal and identical to any other liquid phase. As temperature is decreased more (at moderate pressures), helium does not solidify. In fact, it undergoes a phase change to a second liquid ... 20 The critical point is a phase transition of phase transitions. It's a different beast than your everyday run-of-the-mill phase transition like water boiling to vapor or ice melting to water. The diagram you included is very misleading in its use of color. The "yellow" supercritical regime doesn't really exist. Beyond the critical point, gas is ... 19 Your Question: Which "exotic salt" can lower water's freezing point by$\pu{-70 ^\circ C}$? Here is your "exotic compound" although it is not a salt by definition. It is a base: Aqua ammonia, also called ammoniacal liquor, ammonia liquor, or ammonia water, is produced by dissolving ammonia gas ($\ce{NH3}$) in water. The proper chemical name of aqua ammonia ... 17 Does the activity of a solid or liquid change over the course of a reaction? The density of a solid or liquid reactant doesn't change over the course of a reaction. The mass and volume do as it is consumed, but the ratio of the two is constant. If the reaction causes a temperature change then there are small changes in density, but that would also alter the ... 16$\mathrm{pH}$is the aqueous concentration of$\ce{H3O+}$or$\ce{H+}$ions in soution. I would not say that ice lacks$\ce{H3O+}$and$\ce{OH-}$ions as ice's structure would allow for such, however, since the ions are not in aqueous solution, the material cannot rightfully have a "$\mathrm{pH}$" as we know it. 16 The critical point is a point of convergence of all state properties of the respective liquid and gas. It can be considered as the degeneration point, where there is no difference between gas and liquid and this distinguishing does not make sense any more. It can be also said the supercritical fluid near the critical point is neither gas neither liquid. It ... 15 I believe the ice cracked due to residual strains from freezing. Since ice freezes from the outside inward and it expands as it freezes, that as the inner water freezes, it imparts a tensile force on the surrounding ice (like the opposite effect of tempering glass). As the warm liquid removes ice, the cross sectional area under tension decreases while the ... 15 Yes they will boil all right. Sure, there might be some kinetic impediment to it if you let the liquids to settle in layers, but if you stir them so as to expose their surfaces, they will boil†. This is what steam distillation is all about. As for the first law, it will hold just fine. You burn your firewood, you get the heat, but it is not for free: ... 14 From Changes of Phase (or State): ... So, how could there be a change in heat during a state change without a change in temperature? "During a change in state the heat energy is used to change the bonding between the molecules. In the case of melting, added energy is used to break the bonds between the molecules. In the case of freezing, energy is ... 14 The enthalpy of vaporization of$\ce{HCN}$is higher than for$\ce{NH3}$, which suggests that$\ce{HCN}$molecules interact more strongly than$\ce{NH3}$molecules.$\ce{C-H}$bonds are not usually considered good hydrogen bond donors, but$\ce{HCN}$is unusual. For example$\ce{HCN}$has a$\mathrm pK_\mathrm a$value of 9.2, indicating that the$\ce{CN}\$ ...

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Why do we have water vapors when our body temperature is also <100°C in the first place? At normal pressure, water boils at 100°C, meaning that bubbles of pure steam form under water. At lower temperatures, water molecules reversibly move from the liquid to the gas phase and back. The higher the temperature, the higher the vapor pressure, and the higher ...

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