Effect of impurities on melting point

Question

This is a 10th-grade question.

The melting point of a substance decreases with increase in presence of impurities in it. The melting point of ice decreases from 0 °C to -22 °C on mixing salt in it in proper proportion. That is why salt is added to make freezing mixtures to keep ice creams frozen.

I want to ask that suppose ice melts at 0 °C and we are now adding salt to lower melting point but now when we have lowered the melting point the ice will melt more quickly, then what is the purpose of mixing salt?

Answer 1

In the old fashioned process, we used buttermilk sugar and a bit of vanilla. We put them into a churn (which is just a small bucket with a mixing blade in it so you could hand mix the contents) and then added ice around the churn (outside the churn) to cool it down. And then you mixed and mixed and mixed and mixed. But there was a problem. As the ice melted, the water it created was at $\pu{32°F}$ and even though the ice was colder (a commercial freezer is often about $\pu{15°F}$ (but a home freezer is warmer)) the water kept the churn at around (usually above) $\pu{32°F}$ ($\pu{0°C}$). And ice cream (being an "impure" water mixture) is a liquid, not solid, at $\pu{32°F}$. So, this wouldn't work. Enter salt. If you added salt to the ice, then the ice would melt well below $\pu{32°F}$ and be able to cool the churn to the freezing point of the ice cream mixture (I think its about $\pu{25 °F}$, but my memory shouldn't be trusted). So, the salt and ice form a salt water mixture which can be well below $\pu{32°F}$, and so can cool to below $\pu{32°F}$. If you were to add salt to the recipe, then you'd have to cool it even MORE to get it to freeze (although a little salt makes it taste better, imho). So, salt is added to the cooling mixture and (supposedly) not to the recipe's ingredients. I guess I should point out that you need to consider walking over hot rocks vs walking in say an inch of water at the same temperature. Which would burn you faster? There's no question that your feet would have more contact with the water than the solid rocks, so your feet would be more damaged by the water (water at $180$ or $\pu{200 °F}$ WILL burn you). My point is you actually WANT the ice to melt into a salt solution, because that will better cool the churn - as long as there's enough salt so that the liquid is cold enough.

Answer 2

The presence of impurities in a substance results in a lower melting point due to a process called melting point depression. Melting point depression is the reason why adding salt to frozen streets helps to melt the ice.

Melting point depression occurs due to the nature of a material's solid state. Most solids, such as ice, form as crystalline lattices of repeating ions or molecules. This lattice is held together by intermolecular forces that create a strong, stable structure. The solid must be heated to a certain temperature in order for this structure to be disrupted and for the solid to melt. However, the presence of impurities weakens the lattice, making it less stable. As a result, the compound melts at a lower temperature.

When working with organic compounds in a lab, the purity of the compound can be partially determined through the use of a precise measurement of the melting point. If the melting point is within the scientifically accepted range of the material's melting temperature, then the material is presumed to be pure. However, if the observed melting temperatures are outside of this range, then the compound is not pure.

One interesting effect of this process is that sea water freezes at a lower temperature than pure water. Pure water freezes at $\pu{32°F}$ while sea water freezes at $\pu{28.4 °F}$.

I hope you get it. For any problem, leave comment.

Answer 3

In regards to why you do it.

Lets say you live in an environment where the temperature in the winter is -10 °C. In this case pure water at this temperature will be in the solid phase and exist as ice because the surrounding temperature is lower than the melting/freezing point. For driving and general transport, ice on the roads/walkways, etc. is a hazard so this is a concern. One way to deal with it is to physically remove the ice but this is very difficult.

Instead think about what happens if you add an impurity to the water and it reduces the melting point way down to -22 °C as was given in your textbook. Now the temperature at which the water freezes is much lower than the surrounding temperature so the water will still exist as a liquid form even when the temperature is lower than the normal freezing point. Thus ice problem when driving is solved (aside from the environmental concern of all the heavy salts).

Answer 4

I want to ask that suppose ice melts at 0 °C and we are now adding salt to lower melting point but now when we have lowered the melting point the ice will melt more quickly, then what is the purpose of mixing salt?

You want the temperature of the liquid water to be lower. The process of melting the ice cools down the ice/salt water mixture. This is analogous to evaporation cooling down our bodies when our sweat evaporates.

[comments by Buck Thorn] Note that impurities in a solid will affect the melting point, but the freezing point depression effect (a colligative effect) has to do with impurities in the liquid phase, not in the solid. The paragraph is therefore a tad confusing.

For the ice cream making process, we start with pure ice and have the impurity in the liquid. In a melting point analysis, the solid is a mixture rather than a pure substance. For both cases, the change in melting point has to do with homogeneous mixing in the liquid state, though, according to Lisa Nichols:

An impure solid is typically heterogeneous on the microscopic level, with pure regions of each component distributed through the bulk solid much like granite. When an impure solid is warmed, microscopic melting first occurs in a pure region by the component with the lower melting point (compound A in Figure 6.7a). This microscopic melting is not visible to the eye.

The preliminary melting of compound A in Figure 6.7a forms tiny pools of liquid that begin to dissolve compound B from the bulk solid. As compound B is dissolved into the melt (causing it to become more impure), the freezing point of this mixture is depressed. Compound B will continue to dissolve in the melt, until it reaches the eutectic composition (point a in Figure 6.7b), and the system will continue to melt at this composition until the entirety of the minor component (the impurity) is dissolved. Once the minor component is completely dissolved, further melting continues of the bulk component. This increases the purity of the melt, so the melting temperature increases somewhat. The system follows the melting line in Figure 6.7b either to the left or right of the eutectic temperature (depending on which side of the eutectic point is started), adjusting its melting temperature as the bulk component increases its concentration in the melt. This continues until the entire sample is melted.

Source: https://chem.libretexts.org/Bookshelves/Organic_Chemistry/Organic_Chemistry_Lab_Techniques_(Nichols)/06%3A_Miscellaneous_Techniques/6.01%3A_Melting_Point/6.1C%3A__Melting_Point_Theory