specific heat

Examples of specific heat in the following topics:

• Specific Heat and Heat Capacity

  • the specific heat capacity, often simply called specific heat, which is the heat capacity per unit mass of a pure substance.
  • Given the molar heat capacity or the specific heat for a pure substance, it is possible to calculate the amount of heat required to raise/lower that substance's temperature by a given amount.
  • In these equations, m is the substance's mass in grams (used when calculating with specific heat), and n is the number of moles of substance (used when calculating with molar heat capacity).
  • Specific heat capacity is the measure of the heat energy required to raise the temperature of a given quantity of a substance by one kelvin.
  • The above simulation demonstrates the specific heat and the latent heat.

• Constant-Pressure Calorimetry

  • We already know our equation relating heat (q), specific heat capacity (C), and the change in observed temperature ($\Delta T$) :
  • We will now illustrate how to use this equation to calculate the specific heat capacity of a substance.
  • What is the specific heat of the unknown metal?
  • (The specific heat of water is 4.18 $\frac {J} {g^\circ C}$)
  • The specific heat capacity of the unknown metal is 0.166 $\frac {J} {g ^\circ C}$ .

• Heating Curve for Water

  • A heating curve shows how the temperature changes as a substance is heated up at a constant rate.
  • A constant rate of heating is assumed, so that one can also think of the x-axis as the amount of time that goes by as a substance is heated.
  • The amount of heat added, q, can be computed by: $q=m\cdot C_{H_2O(s)}\cdot \Delta T$ , where m is the mass of the sample of water, C is the specific heat capacity of solid water, or ice, and $\Delta T$ is the change in temperature during the process.
  • Note that the specific heat capacity of liquid water is different than that of ice.
  • Note that the specific heat capacity of gaseous water is different than that of ice or liquid water.

• Constant-Volume Calorimetry

  • The total heat given off in the reaction will be equal to the heat gained by the water and the calorimeter:
  • Keep in mind that the heat gained by the calorimeter is the sum of the heat gained by the water, as well as the calorimeter itself.
  • where Cwater denotes the specific heat capacity of the water ($1 \frac{cal}{g ^{\circ}C}$), and Ccal is the heat capacity of the calorimeter (typically in $\frac{cal}{^{\circ}C}$).
  • The sample is ignited by an iron wire ignition coil that glows when heated.
  • From the change in temperature, the heat of reaction can be calculated.

• Variation of Physical Properties Within a Group

  • Metallic elements are shiny, usually gray or silver in color, and conductive of heat and electricity.

• Solutions and Heats of Hydration

  • When ions dissolve in water, the stabilizing interactions that result release energy called the "heat of hydration."
  • Ionic solids are insoluble in the majority of non-aqueous solvents, but they tend to have high solubility specifically in water.
  • M^+ (g) + X^-(g) \to M^+ (aq) + X^-(aq)$ [heat of hydration]
  • A hot solution results when the heat of hydration is much greater than the lattice energy of the solute.
  • Predict whether a given ionic solid will dissolve in water given the lattice energy and heat of hydration

• Heat and Work

  • Heat transfer by convection occurs through a medium.
  • Lastly, heat can also be transferred by radiation; a hot object can convey heat to anything in its surroundings via electromagnetic radiation.
  • Like heat, the unit measurement for work is joules (J).
  • Heat and work are related.
  • Work can be completely converted into heat, but the reverse is not true: heat energy cannot be wholly transformed into work energy.

• The Three Laws of Thermodynamics

  • This law says that there are two kinds of processes, heat and work, that can lead to a change in the internal energy of a system.
  • Since both heat and work can be measured and quantified, this is the same as saying that any change in the energy of a system must result in a corresponding change in the energy of the surroundings outside the system.
  • If heat flows into a system or the surroundings do work on it, the internal energy increases and the sign of q and w are positive.
  • Conversely, heat flow out of the system or work done by the system (on the surroundings) will be at the expense of the internal energy, and q and w will therefore be negative.
  • Specifically, the entropy of a pure crystalline substance (perfect order) at absolute zero temperature is zero.

• Changes in Energy

  • The concept of entropy can be described qualitatively as a measure of energy dispersal at a specific temperature.
  • This is because some energy is expended as heat, limiting the amount of work a system can do.
  • The state function has the important property that in any process where the system gives up energy ΔE, and its entropy falls by ΔS, a quantity at least TR ΔS of that energy must be given up to the system's surroundings as unusable heat (TR is the temperature of the system's external surroundings).

• Historical Background

  • In an experiment designed to prepare ammonium cyanate from silver cyanate, he heated the latter with ammonium chloride expecting the outcome shown below.
  • Thus, strongly heating organic substances such as carbohydrates and proteins yielded water, ammonia and carbonaceous solids (all inorganic), with loss of the vial essence.
  • First, carbon disulfide, obtained by reaction of carbon with sulfur, was converted to carbon tetrachloride by heating with chlorine, and the simultaneous pyrolysis of CCl4 yielded a mixture of products which included tetrachloroethene, presumably formed from dichlorocarbene (:CCl2).
  • CS2 + Cl2 + heat —> CCl4 + Cl2C=CCl2 + many other products
  • That the atoms of allyltoluidine should, in the course of one reaction, selectively reorganize and combine in this specific fashion is beyond all reasonable probability.