Heat capacity is the ability of a material to absorb heat without directly reflecting all of it as a rise in temperature. You should read the sections on heat and temperature as background, and the water section would help, too.
As heat is added uniformly to like quantities of different substances, their temperatures can rise at different rates. For example, metals,
good conductors of heat, show fast temperature rises when heated. It is relatively easy to heat a metal until it glows red. On the other hand, water can absorb a lot of heat with a relatively small rise in temperature. Insulating materials (insulators) are very poor conductors of heat, and are used to isolate materials that need to be kept at different temperatures — like the inside of your house from the outside.
This graph shows the rise in temperature as heat is added at the same rate to equal masses of aluminium (Al) and water (H2O). The temperature of water rises much more slowly than that of Al.
In the metal, Al atoms only have translational kinetic energy (although that motion is coupled strongly to neighbor atoms). Water, on the other hand, can rotate and vibrate as well. These degrees of freedom of motion can absorb kinetic energy without reflecting it as a rise in temperature of the substance.
Most substances obey the law of equipartition of energy over a broad range of temperatures. The law says that energy tends to be distributed evenly among all of the degrees of freedom of a molecule — translation, rotation and vibration. This has consequences for substances with more or fewer atoms. In the diagram below, each container represents a degree of freedom. The situations for a 3-atom and a 10-atom
molecule are shown. If the same total amount of heat energy is added to each molecule, the 3-atom molecule ends up with more energy in its translational degrees of freedom. Because the 10-atom molecule has more vibrational modes in which to store kinetic energy, less is available to go into the translational modes, and it is mostly the translational energy that we measure as temperature.
There's one more refinement left to make to heat capacity. Obviously, the amount of heat required to raise the temperature of a large quantity of a substance is greater than the amount required for a small amount of the same substance.
To control for the amount, we generally measure and report heat capacities as specific heat, the heat capacity per unit mass.
Specific heats of a great many substances have been measured under a variety of conditions. They are tabulated in books an on-line.
We generally choose units of J/gram or KJ/Kg. The specific heat of liquid water is 4.184 J/g, which is also 4.184 KJ/Kg. The calorie is a unit of heat defined as the amount of heat required to raise the temperature of 1 cm3 of water by 1˚C.
Specific heat is the heat capacity per unit mass.
The specific heat of water is 1 cal/g˚C = 4.184J/g˚C
The heat, q, required to raise the temperature of a mass, m, of a substance by an amount ΔT is
q = mCΔT = mC(Tf - Ti)
where C is the specific heat and Tf and Ti are the final and initial temperatures.
The slope of a graph of temperature vs. heat added to a unit mass is just 1/C.
Using this formula, it's relatively easy to calculate heat added, final or initial temperature or the specific heat itself (that's how it's measured) if the other variables are known.
The heat q added or evolved for a temperature change of a mass m of a substance with specific heat C is
The units of specific heat are usually J/mol·K (J·mol-1K-1) or J/g·K (J·g-1·K-1). Remember that it's OK to swap ˚C for K because the size of the Celsius degree and the Kelvin are the same.
(Use the table below to look up missing specific heats.)
Phase changes are a big source or sink of heat. Here, for example, is the heating curve of water.
It shows the rise in temperature as heat is added at a constant rate to water. Here's what's going on in regions A-E:
A. Heat is added to solid water (ice) below 0˚C, and its temperature rises at a constant rate.
B. Solid ice is melted to liquid water. During the addition of the latent heat of fusion (ΔHf), no temperature rise is observed, but hydrogen bonds holding the ice together break.
C. Heat is added to liquid water above 0˚C, and its temperature rises at a constant rate until the boiling point at 100˚C.
D. Water at 100˚C absorbs a great deal of heat energy at 100˚C as it undergoes a phase transition from liquid to gas. This is the latent heat of vaporization, ΔHv, the energy it takes for water to have no more cohesive force.
E. Finally, gaseous water above 100˚C absorbs heat, increasing its temperature at a constant rate. Water has no more phase transitions after this.
The relatively large attractive intermolecular forces between water molecules gives water very high heats of fusion and vaporization. Compared to most other substances, it takes a large amount of heat to melt water ice and to boil or evaporate water.
Enthalpies of fusion and vaporization are tabulated and can be looked up. The Wikipedia page of a compound is usually a good place to find them. Below we'll do an example of a heat calculation as the temperature of a substance rises through a phase change.
Cohesive forces are forces that hold a substance together. When water hits a waxy or hydrophobic surface, it forms small sphere-like drops – "beads." These beads of water minimize the contact with the surface and with the air, and maximize the contact of water with itself. Liquid water is very cohesive. It forms intermittent, but relatively strong bonds with itself.
Other substances like CO2 lack such strong intermolecular attractions, and don't form liquids or solids unless very cold or at very high pressure.
The heat absorbed or released upon a phase transition is calculated by multiplying the enthalpy of vaporization, ΔHv, or the enthalpy of fusion, ΔHf by the number of moles of substance:
The enthalpy of fusion is often called the "latent heat of fusion" and the enthalpy of vaporization is often called the "latent heat of vaporization."
The units of ΔHf and ΔHv are Joules/mole (J·mol-1) or J/g (J·g-1).
Solution: There is a phase transition of water in this temperature range, so this problem will comprise three steps:
Here are the calculations for each of our steps:
Step 1:The amount of heat required to raise the temperature of ice (before it melts) by 20˚C is:
Note that we've converted the Celsius temperatures to Kelvin.
Step 2: The amount of heat required to melt 18 g of ice is:
Step 3: The amount of heat required to raise the temperature of liquid water by 25˚C is:
Adding all of these energies up, we get the total, q = 2642 J,
Now let's compare this to a similar calculation, but this time we'll heat liquid water through its boiling point to a gas.
Solution: This is also a three-step problem, but this time we're vaproizing water. Here are the steps:
Here are the calculations for each of our steps:
Step 1:The amount of heat required to raise the temperature of water (before it vaporizes) from 80˚C to 100˚C is:
(Use the table below to look up missing specific heats; heats of fusion or vaporization are given in the problems.)
|water (ice)||2.11||methanol (CH3OH)||2.14|
|water (liquid)||4.184||ethanol (C2H5OH)||2.44|
|water (steam)||2.08||ethylene glycol (C2O2H6)||2.2|
|aluminum (s)||0.897||hydrogen (H2) gas||14.267|
|copper (s)||0.385||benzene (C6H6)||1.750|
|iron (s)||0.450||wood (typical)||1.674|
|lead (s)||0.129||glass (typical)||0.867|
Until now, in order to keep things simple, I've been referring to specific heat as C. But heat capacity (specific heat if it's per mole or per gram) can change depending on whether the thermodynamic variables of pressure or temperature are held constant during heating or cooling.
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