Examples of heat in the following topics:
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- the molar heat capacity, which is the heat capacity per mole of a pure substance.
- the specific heat capacity, often simply called specific heat, which is the heat capacity per unit mass of a pure substance.
- Now we can plug our values into the formula that relates heat and heat capacity:
- Latent heat of melting describes tœhe amount of heat required to melt a solid.
- The above simulation demonstrates the specific heat and the latent heat.
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- 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.
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- Water transitions from ice to liquid to water vapor as heat is added to it.
- 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.
- Instead, use the heat of fusion ($\Delta H_{fusion}$ ) to calculate how much heat was involved in that process: $q=m\cdot \Delta H_{fusion}$, where m is the mass of the sample of water.
- Use the heat of vaporization ($\Delta H_{vap}$ ) to calculate how much heat was absorbed in this process: $q=m\cdot C_{H_2O(g)}\cdot \Delta T$, where m is the mass of the sample of water.
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- 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.
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- The heat transferred to/from the solution in order for the reaction to occur is equal to the change in enthalpy ($\Delta H = q_P$), and a constant-pressure calorimeter thus measures this heat of reaction.
- We already know our equation relating heat (q), specific heat capacity (C), and the change in observed temperature ($\Delta T$) :
- What is the specific heat of the unknown metal?
- (The specific heat of water is 4.18 $\frac {J} {g^\circ C}$)
- The number of joules of heat released into each gram of the solution is calculated from the product of the rise in temperature and the specific heat capacity of water (assuming that the solution is dilute enough so that its specific heat capacity is the same as that of pure water's).
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- In an endothermic system, the $\Delta H$ value is positive, so the reaction absorbs heat into the system.
- Notice that in an endothermic reaction like the one depicted above, we can think of heat as being a reactant, just like A and B.
- In an exothermic system, the $\Delta H$ value is negative, so heat is given off by the reaction.
- $A + B \rightarrow C + heat,\: \Delta H = -$
- Notice that here, we can think of heat as being a product in the reaction.
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- At constant pressure, the change in enthalpy is equal to the heat given off, or the heat absorbed, in a given chemical reaction:
- Due to this relation, the change in enthalpy, $\Delta H$, is often referred to as the "heat of reaction."
- In order to melt the ice cube, heat is required, so the process is endothermic.
- Paul Andersen explains how heat can be absorbed in endothermic or released in exothermic reactions.
- Therefore, the change in enthalpy is negative, and heat is released to the surroundings.
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- This stored chemical energy, or heat content, of the system is known as its enthalpy.
- Exothermic reactions release heat and light into their surroundings.
- Excess energy from the reaction is released as heat and light.
- Endothermic reactions, on the other hand, absorb heat and/or light from their surroundings.
- Significant heat energy is needed for this reaction to proceed, so the reaction is endothermic.
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- Heat of solution refers to the change in enthalpy when a solute is dissolved into a solvent.
- The heat of solution, like all enthalpy changes, is expressed in kJ/mol for a reaction taking place at standard conditions (298.15 K and 1 bar).
- The heat of solution can be regarded as the sum of the enthalpy changes of three intermediate steps:
- The value of the overall heat of solution, $\Delta H^\circ_{sol}$, is the sum of these individual steps.
- Depending on the relative signs and magnitudes of each step, the overall heat of solution can be either positive or negative, and therefore either endothermic or exothermic.
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- CH3-CH2-CH3 + 5 O2 —> 3 CO2 + 4 H2O + heat
- Precise heats of combustion measurements can provide useful iinformation about the structure of molecules.
- CH3-CH2-CH3 + 4 O2 —> CO2 + 2 CO + 4 H2O + heat
- From the previous discussion, we might expect isomers to have identical heats of combustion.
- Thus, the heat of combustion of pentane is –782 kcal/mole, but that of its 2,2-dimethylpropane (neopentane) isomer is –777 kcal/mole.