Heat transfer is all around! The episode starts by describing the heat transfer on a particle level.
Heat transfer is all around! The episode starts by describing the heat transfer on a particle level (0:58). Heat transferred can be quantified using the heat transfer equation (1:42). The heat transfer equation takes into account the mass of the substance (2:32), the specific heat capacity (4:12) as well as the change in temperature (6:05). It can be measured in an experimental set up by using calorimetry (6:18).
Question (8:01): What is the molar heat capacity of water in J/mol K?
Thank you for listening to The APsolute RecAP: Chemistry Edition!
(AP is a registered trademark of the College Board and is not affiliated with The APsolute RecAP. Copyright 2020 - The APsolute RecAP, LLC. All rights reserved.)
Website:
EMAIL:
Follow Us:
Hi and welcome to the APsolute Recap: Chemistry Edition. Today’s episode will recap heat capacity and calorimetry.
Lets Zoom Out:
Unit 6 - Thermodynamics
Topic 6.4 - Heat Capacity and Calorimetry
Big idea - Energy
An ice cube melting in your warm hand or in a glass of water, frozen food thawing or the hot sand under our feet at the beach: heat is being transferred from a hotter object to a cooler object. But what happens when heat is being transferred and how can we quantify the heat transfer?
Let’s zoom in:
What happens when we transfer heat? From our gas law episodes, we already know that temperature is proportional to average kinetic energy. The particles in hot water therefore have a higher kinetic energy than the particles in cold water. If we are pouring them together, the particles collide and the contact can result in a transfer of thermal energy - from the particle with higher average kinetic energy to the particle with lower average kinetic energy. Eventually, and obviously I am using this term rather loosely, since it can happen rather fast, thermal equilibrium is reached. All particles have the same average kinetic energy and therefore the same temperature. Tada! You can now enjoy your shower - without burning up or freezing!
But let’s quantify our heat transfer using q equals m c delta T. With this equation we can calculate the amount of energy, q, being released or absorbed by a system during heat transfer. The amount of energy that is released or absorbed has to stay constant, according to the first law of thermodynamics. It states that energy is conserved in chemical and physical processes. So if the system loses 2.7 kJ, the surroundings have to absorb 2.7 kJ. As you can see in the equation, the amount of energy is determined by the mass, the specific heat capacity and the actual change in temperature. Let’s take a closer look at these variables in a systematic approach: We are keeping two constant and are changing one of them to determine the effect it has on our dependent variable, the heat q.
Let’s make some hot water for tea. We want to change the temperature of water from 20 degrees Celsius, which is room temperature to 85 degrees Celsius. Quick life hack: tea water should never boil! It changes the taste of the steeped tea. In both setups, we use water, therefore the specific heat capacity is the same, and in both instances we want to raise the temperature by 65 degrees Celsius. But: we are heating different amounts of water: a lovely cup versus an entire pot. Which one will take more energy? Well, you’ve certainly experienced it, especially when using a pot and not one of those electric kettles: the more water you want to heat, the more energy you need.
The second variable we can isolate is the specific heat capacity, c. What’s that? It is the amount of energy required to raise the temperature of 1 g of a substance by 1 K (or 1 °C). And it is SPECIFIC to the substance. Water has a specific heat capacity of 4.184 J - which is, fun fact, 1 calorie! That means, it takes 4.184 J to raise the temperature of 1g of water by 1 degree Celsius. These values are tabulated, no need to memorize them! In comparison, copper has a specific heat capacity of 0.385 J. It takes much less energy to raise the temperature of 1 g of copper! Generally speaking - and as we know Chemistry is the Science of exceptions - metals have lower specific heat capacities than nonmetals. And, as usual: Water IS special with an unusually high specific heat capacity - THANK YOU hydrogen bonding!
To test this in an experimental approach, we could determine the amount of energy it takes to heat 100 g of water by 20 degrees Celsius vs the amount of energy it takes to heat 100 g of copper by 20 degrees. For water, this would be 8,368 J, for copper 770 J. Therefore, the higher the specific heat capacity of a substance, the more energy is needed to raise its temperature.
On a quick side note: Sometimes, you can also encounter the molar heat capacity, which is defined as the amount of energy required to raise the temperature of ONE MOLE of a substance by 1 degree Celsius. Be careful - if this refers to mole, you have to convert your mass to moles as well! Hello Molar Mass!
The third variable we can change is, of course, the amount of temperature. This is pretty straight forward: It takes more energy to raise the temperature of 100g of water by 20 degrees Celsius than by 5 degrees Celsius.
To measure the heat transfer, scientists use calorimetry. In the scope of AP Chemistry, we are focusing on coffee cup calorimetry, which measures the amount of heat released or absorbed when creating aqueous solutions - for example like our hand warmer or cooling packs. The set up is pretty simple: We have two styrofoam cups nested together to create a somewhat isolated system. To measure the change in temperature, we use a thermometer. Often, we also add a magnetic stirrer to make sure the reaction goes to completion. Knowing the specific heat capacity of water, namely 4.184 J, as well as the amount of substance and measuring the change in temperature during the dissolution, we can calculate the amount of heat being released in exothermic reactions and absorbed in endothermic reactions. One tip for AP Chemistry: your calculations refer to the solution, so you use the mass of water PLUS the mass of the compound you are dissolving as mass in the equation or, if you are combining two aqueous solutions, use the combined mass.
To recap:
During heat transfer, energy is being transferred from substances with higher kinetic energy to substances with lower kinetic energy until thermal equilibrium is reached. To quantify the amount of heat transferred, we can use q = m c delta T where m is the mass in grams, c is the specific heat capacity and delta T is the change in temperature. The greater the mass and the greater the change in temperature, the greater the amount of heat being released or absorbed. The specific heat capacity is specific to a substance. Calorimetry is an experimental approach to measure the transfer of heat.
Today’s Question of the day is about molar heat capacity.
Question:
What is the molar heat capacity of water in J/mol K?