Chemical transformations · Chemical thermodynamics
Energy, heat and work
Crack a hand-warmer open at a Green Bay Packers game in January and it heats up in your pocket. Press a cold pack on a sore ankle and it chills the skin. Energy moves in both cases, but where does it come from, and how can you measure how much moved? This page shows how to track energy with a thermometer and a balance, and how to account for it when a gas pushes a piston.
Energy, heat and temperature
Energy is never created or destroyed. It is only stored in a system or transferred between a system and its surroundings. This is the conservation of energy, and it underpins every energy calculation in thermochemistry.
Heat is a transfer, not something a body contains. A hot block does not hold a store of "heat". It holds internal energy (the kinetic and potential energy of its particles). When the block is placed in cold water, energy flows out of it as heat. The same idea applies in a reaction: if a reaction releases energy, the solution around it receives that energy and its temperature rises.
Calorimetry: measuring heat with a thermometer
Drop a hot metal nut into a beaker of cool water. The water warms up and the metal cools until they reach the same temperature. How much heat moved? You only need the masses, the temperatures and one number for each material.
In an insulated container (a calorimeter, for example a coffee-cup calorimeter: two nested foam cups with a lid), no heat escapes. All the heat lost by the hot object is gained by the cooler one, so $$q_{\text{lost}} = q_{\text{gained}} \quad\text{or}\quad \sum q = 0.$$
The same method gives $\Delta H$ for a reaction in solution. In a neutralization lab, the reaction happens in the calorimeter and the temperature rise of the solution gives the heat released.
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In a lab experiment, an aluminum block of mass $m = 150.0\,\text{g}$ is heated from an initial temperature of $18.0\,^{\circ}\text{C}$ to $68.0\,^{\circ}\text{C}$. Given…
In a lab experiment, an aluminum block of mass $m = 150.0\,\text{g}$ is heated from an initial temperature of $18.0\,^{\circ}\text{C}$ to $68.0\,^{\circ}\text{C}$. Given that the specific heat capacity of aluminum over this range is $c = 0.897\,\text{J}\times\text{g}^{-1}\times\text{K}^{-1}$ and assuming that heat losses to the surroundings are negligible:
a) Calculate the quantity of heat absorbed by the metal block, expressed in joules and in kilojoules.
b) Justify conceptually why the heat absorbed is not an intrinsic property possessed by the block, clearly distinguishing between heat and temperature.
In a constant-pressure mixing calorimeter whose own heat capacity is negligible, $250.0\,\text{g}$ of liquid water at a temperature of $20.0\,^{\circ}\text{C}$ are…
In a constant-pressure mixing calorimeter whose own heat capacity is negligible, $250.0\,\text{g}$ of liquid water at a temperature of $20.0\,^{\circ}\text{C}$ are introduced. Next, a metal piece of mass $80.0\,\text{g}$, taken from an oven at $95.0\,^{\circ}\text{C}$, is added. The mixture is stirred until a thermal equilibrium temperature of $24.2\,^{\circ}\text{C}$ is reached. Data: specific heat capacity of liquid water $c_{\text{agua}} = 4,184\,\text{J}\times\text{g}^{-1}\times\text{K}^{-1}$. Assume that there are no heat losses to the surroundings.
a) Set up the calorimetry balance and determine the specific heat capacity of the metal, expressed in $\text{J}\times\text{g}^{-1}\times\text{K}^{-1}$.
b) If a small heat loss to the surroundings had occurred during the experiment, explain whether the experimental value obtained for the specific heat capacity of the metal would be greater or less than the true value.
Thermodynamics rests on four laws. You can see each one in a kitchen or a gym.
Zeroth law: if A is in thermal equilibrium with B, and B with C, then A is in equilibrium with C. This is why a thermometer works: it reaches the same temperature as your skin or a cup of soup, so its reading tells you that temperature.
First law: energy is conserved. Energy is transferred as heat or work, never created or destroyed. A car engine turns the chemical energy of gasoline into motion and heat.
Second law: heat flows spontaneously from hot to cold, never the reverse, and disorder (entropy) of the universe increases. A hot slice of pizza cools to room temperature, and it never heats itself back up.
Third law: as a substance approaches absolute zero ($0\ \text{K}$), its entropy approaches a minimum. You can get closer and closer to $0\ \text{K}$, but never reach it.
Work and the first law
Heat is only one way to transfer energy. A gas can also transfer energy by pushing something, which is work. Physics students meet work as $W = Fs$.
Work of expansion: a sealed gas syringe is warmed gently and the plunger moves out, pushing back the air around it. The gas has done work on the surroundings.
Energy is conserved, so the internal energy $U$ of a closed system changes only through heat and work:
Constant volume, constant pressure and ΔH
In the neutralization lab the cup is open to the air, so the reaction happens at constant pressure. The heat measured is the enthalpy change: $q_P = \Delta H$. That is why a coffee-cup calorimeter gives you $\Delta H$ directly.
Summary
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