A change in state, also known as a phase change, refers to when a substance changes between state, such as from liquid to solid, liquid to gas, and so on. Chocolate melting is one example of a phase change.
Recall from earlier sections that with sufficient energy, bonds will break. Consider the implications of that and think about what happens to ice when its taken out of the freezer and left at room temperature.
From everyday knowledge, we know that the ice will melt to become a puddle of liquid water. This is due to the change in energy, in the form of heat, being sufficient to break some of the bonds in the ice structure. Not all of the bonds are broken, because not enough energy is provided. If you can visualize it, a solid with some broken bonds is just liquid! Similarly, what happens when liquid water is heated up? The water becomes gaseous water vapor, since the increased temperature provides sufficient energy to break all of the water bonds, which results in a gas. We can sum up these observations with these two statements:
Changes in energy can cause changes in state.
Changes in state result from the breaking or forming of bonds
For most examples in everyday life, phase changes are due solely to changes in temperature. We'll explain this later on, but phase changes can be caused through changes in pressure as well. That usually catches people off guard, but think about it: what will increasing/decreasing pressure due to the bonds of a material? What does increasing pressure due to the internal energy of a system (think thermodynamics)?
The changes in state have the following titles:
| Initial State | Final State | Process |
Solid |
Liquid |
Melting/Fusion |
Solid |
Gas |
Sublimation |
Liquid |
Gas |
Vaporization |
Liquid |
Solid |
Solidification |
Gas |
Solid |
Deposition |
Gas |
Liquid |
Condensation |
The melting point for a substance is the temperature at which the substance changes from solid to liquid. Likewise, the boiling point for a substance is the temperature at which the substance changes from liquid to gas. Finally, the freezing point for a substance is the temperature at which the substance will change from liquid to solid.
Each of these phase changes corresponds to a specify enthalpy. Whenever a substance undergoes a phase change, it either absorbs or releases a specific amount of energy.
For vaporization, this enthalpy is called the heat of vaporization (`DeltaH_"vaporization"`). The heat of vaporization is unique to each compound and is always a positive quantity. Think about it: in order to break the bonds of the liquid state to become the solid state, energy must be absorbed.
Acetone and liquid iron have heats of vaporization equal to `(31.3 "kJ")/"mol"` and `(340 "kJ")/"mol"` . Since the heat of vaporization of liquid iron is more than 10 times the heat of vaporization of acetone, it would take 10 times the energy to vaporize liquid iron than acetone. If you've ever worked with acetone, you'll know that acetone evaporates just by being left out whereas liquid iron is likely to stay as liquid iron. The heats of vaporization explain that difference in behavior!
For melting, the energy is called the heat of fusion (`DeltaH_"fusion"`). Just like with the heat of vaporization, the heat of fusion is a positive quantity.
The `DeltaH_"fusion"` of ice is `(6 "kJ")/"mol"` which means that, in order to melt, ice has to absorb `6 kJ` of energy for every mole of ice.
From everyday experiences, you'll know that in order to melt a solid or vaporize a liquid, you have to heat up the solid or liquid. The purpose of the heat is to provide sufficient energy to overcome the heats of fusion/vaporization.
What about the enthalpies for liquids to solids or gases to liquid? For the reverse processes, the enthalpies are simply the negative of the original process! In order for a gas to condense, it has to release the corresponding `DeltaH_"vaporization"`. Likewise, in order for a liquid to solidify, it has to release the corresponding `DeltaH_"fusion"`.
If this doesn't make sense, think of it this way:
Whenever bonds are broken (`"solid" rArr "liquid" rArr "gas"`), energy must be absorbed by the bonds and therefore the `DeltaH` must be positive.
Whenever bonds are formed (`"gas" rArr "liquid" rArr "solid"`), energy must be released and thus the `DeltaH` must be negative.
This table should help:
| Initial State | Final State | Enthalpy |
Solid |
Liquid |
`DeltaH_"fusion"` |
Liquid |
Solid |
`-DeltaH_"fusion"` |
Liquid |
Gas |
`DeltaH_"vaporization"` |
Gas |
Liquid |
`-DeltaH_"vaporization"` |
Colligative properties are properties of a solid that only depend on the amount of the solute, not the identity. Properties such as reactivity, color, etc are non-colligative properties because they vary depending on what the solute is.
Recall that, when a salt is dissolved in solute, the salt dissolves into its respective ions. `NaCl` for example:
`NaCl + H_2O_((l)) rArr Na^+ + Cl^- + H_2O_((l))`
And `MgBr_2` :
`MgBr_2 + H_2O_((l)) rArr Mg^(2+) + 2Br^- + H_2O_((l))`
Colligative properties in the context of phase changes are properties that are dependent on the number of ions in solution. For `NaCl`, there are 2 ions, `Na^+ and Cl^-`. For `MgBr_2`, there are 3 ions, `Mg^(2+) and 2Br^-`. The identities of the ions do not matter, only the number.
What happens to the melting and boiling points of a solution when a salt is added to the solution? It turns out that the melting and boiling points will actually shift depending on the number of ions present. A solution of pure water and a solution of water + salt will both have different melting and boiling points! This behavior is referred to as freezing point depression and boiling point elevation.
We can calculate how much the freezing point depresses through the following equation:
`-i*K_f*m=DeltaT_f`
`i` is the Van't Hoff factor and is the number of ions in solution
`K_f` is a constant that depends on the identity of the solution
`m` is the molality of the solution
`DeltaT_f` is the change in freezing point
Notice that `m` stands for molality and not molarity (`M`). Molality is defined as `"moles solute"/"mass of solvent"` . For example, The molality of a solution with `5 "moles" NaCl` in `1000 g "water"` is:
`(5 "moles" NaCl)/(1 "kg" "water") = 5 m NaCl`
The boiling point elevation equation is essentially the same:
`-i*K_b*m=DeltaT_b`
`i` is the Van't Hoff factor and is the number of ions in solution
`K_b` is a constant that depends on the identity of the solution
`m` is the molality of the solution
`DeltaT_b` is the change in boiling point.
Let's do a problem to illustrate how these formulas work.
#1. Calculate the molality of `30 "g" "KBr"` dissolved in `550 "mL" "of water"`
The equation for molality is `"moles solute"/"mass of solvent"`. First, convert the solute to moles and the solvent to kg:
`(30 "g" "KBr")((1 "mol")/(119 "g" "KBr"))=0.252 "mol"`
Water has a density of `(1 "g")/"mL"`
`(550 "mL" "water")((1 "g")/"mL")((1 "kg")/(1000 "g"))=0.55 "kg"`
`"molality"=(0.252 "mol solute")/(0.55 "kg solvent") = 0.46 m`
Answer: `0.46 m "KBr"`
#2. Calculate the freezing point of water after `35 "g"` of an unknown salt with molecular weight `17.5 "g"/"mol"` is dissolved in `750 "mL"` of water. The `K_f` of water is `(1.86°C)/"m"` and the Van't Hoff factor is `3`.
As with #1, first convert the solute to moles and the solvent to kg.
`(35 "g")((1 "mol")/(17.5 "g"))=2 "mol"`
`(750 "mL water")((1 "g")/(1 "mL"))((1 "kg")/(1000 "g"))=0.75 kg`
`"molality"=(2 "mol solute")/(0.75 "kg solvent")=2.67 m`
`DeltaT_f=-i*K_f*m`
`DeltaT_f=-(3)((1.86°C)/"m")(2.67 m)=14.9°C`
The normal freezing point of water is `0°C`. This will be our `T_i`
`T_f - T_i=-14.9°C`
`T_f=-14.9°C+T_i`
`T_f=-14.9°C`
By adding the salt, the freezing point of water decreases by almost `15°C`! Isn't that cool!?
Answer: `T_f=-14.9°C`
1. When a substance changes its state, we call it a phase change.
2. Phase changes are caused by changes in energy. This energy goes to the breaking or formation of bonds.
3. Substances have specific temperatures at which they undergo phase changes. For example, the melting point of a substance is the temperature at which the substance will melt.
4. Whenever a substance undergoes a phase change, the substance will absorb or release a specific amount of heat.
5. Colligative properties are properties that do not depend on the identity of the solute, only the amount of the solute.
6. The melting and boiling points of a solution can be changed through the introduction of ions into the solution.
1. Melting Point Apparatus
The melting point of a compound is usually distinct and well categorized enough such that we can use the melting point to test the identity of a compound. Below is an image of a melting point apparatus, a device that heats up to high temperature and contains a viewing window for the user to determine the melting point of a substance.
By putting a small sample of a compound into a tube and heating up the compound until it melts, chemists are able to determine the melting point of the compound, thus provind g them with an educated guess at the identity of a compound.
2. Electrolytes
Electrolytes are just ions! Anytime you drink a sports drink for electrolytes, you're drinking a solution with ions in them - most commonly `Na^+ and K^+`.
3. Salt on roads
In places where roads ice over, large amounts of salt are dumped onto the road in order to melt the ice. Why does this work?
As we learned in this section, adding ions to a solution lowers the freezing point of the solution. By adding salt onto the frozen road, the freezing point of ice is lowered. Eventually, the freezing point is lowered so much that the ice no longer freezes and returns to water, thereby "de-frosting" the road.
4. Freezing Point Elevation
Water is one of few solutions that exhibit a freezing point depression. Most substances exhibit freezing point elevations.