Specific Heat Lab and Thermodynamics Review

Specific Heat Lab - Analysis Questions

• Include data and examples in your analysis questions to support your conclusions.
• Use data to demonstrate the differences in specific heat between substances.

Showing Work and Rounding

• Show your work clearly, either on the lab sheet or attached to it.
• Include units in your calculations.
• Round answers reasonably, avoiding excessive decimal places.

Percent Error

• Expectations for percent error are typically around 10%, unless otherwise specified.
• Be cautious about temperature changes in the water; this is a critical area for error.

Repeating Trials and Checking Masses

• If a trial yields a high percent error (e.g., 17%), consider running it again.
• Double-check the mass of the calorimeter cup.
• Verify the masses of the metal samples (copper, iron, aluminum) if redoing trials, as they may have been mixed up between tables.

Technique for Temperature Measurement:

• Instead of using a stirring rod, swirl the cup while keeping the temperature probe inside to more effectively distribute temperature evenly.

  • Swirling is more aggressive and effective with small water volumes.
  • Avoid probe placement too close or too far from the metal.

• Use smaller amounts of water to achieve a larger \Delta T for the water, which will reduce the percentage of error.

• Focus on precise initial and final water temperatures.

Upcoming Lab & Specific Heat Problems Discussion

• The next lab involves adding ice to water to cool it.
• This process is quicker than the specific heat lab.
• Remaining class time can be used to rework specific heat problems.
• Complete textbook problems (8, 10, 12, 14, 16) to prepare for the test.
• A similar type of problem will be on the test, which will require understanding of calorimeter heat capacity.
• Resource availability: Solution manuals provide work for odd-numbered problems; even-numbered solutions may also be available.

Thermodynamics: Heat of Reaction

• Heat of reaction ($\Delta H$): positive for endothermic reactions (energy increases from reactants to products) and negative for exothermic reactions (energy decreases).

Thermodynamic Table Data Provided

• A data table will be provided with three values for each substance:
  • Enthalpy ($\Delta H$): heat energy in a reaction.
  • Entropy (S): measure of disorder or chaos within a system.
  • Gibbs Free Energy (G): used to analyze the spontaneity of chemical reactions.

Entropy

• Entropy: Degree of chaos; more organized = less entropy. Gases > Liquids > Solids in terms of entropy.
• The universe trends towards increased entropy.
• The combination of entropy and whether a reaction is exothermic or endothermic determines Gibbs Free Energy.

Enthalpy

• For potential energy diagrams, the difference between the energy of products and reactants yields \Delta H .
• Heats of formation: energy required to form a substance (positive value) or energy released when forming a substance (negative value).

Key Pattern: Pure Elements

• The heat of formation of a pure element in its natural state at 25°C and 100 kPa is zero.

  • Examples: Aluminum (Al), Carbon (C), Copper (Cu), Fluorine (F).
  • This applies to Gibbs Free Energy as well, but not to entropy, which always has a value > 0.

Units

• Enthalpy is measured in kilojoules per mole (kJ/mol); pay attention to the amounts (moles) in balanced equations.

Calculating Heat of Reaction

• The heat of reaction can be calculated by taking the sum of the heats of formation of the products minus the sum of the heats of formation of the reactants.

  \Delta H{reaction} = \sum \Delta H{products} - \sum \Delta H_{reactants}

Example: Formation of Liquid Water

• 2H2(g) + O2(g) \rightarrow 2H_2O(l)
• If the heat of formation for liquid water ($\H_2O(l)$) is -285.8 kJ/mol, then the \Delta H for the reaction is 2 \times -285.8 \text{ kJ} = -571.6 \text{ kJ} .
• Since hydrogen and oxygen are pure elements in their natural state, their heats of formation are zero.
• Heat of formation is always expressed per mole.
• When forming water exothermically, the reaction is exothermic regardless; heat of formation of produced water is negative.

Practice Problems

• Problem 1: Formation of Fe2O3

  • Given: 2Fe(s) + 1.5O2(g) \rightarrow Fe2O_3(s), \Delta H = -1643 \text{ kJ}
  • Since iron (Fe) and oxygen (O_2) are pure elements, their heats of formation are zero.
  • Therefore, \Delta H = 2 \times x - (4 \times 0 + 3 \times 0) implies 2x = -1643 \text{ kJ} , so x = -821.5 \text{ kJ/mol}

• Problem 2: Decomposition of Mercury Oxide

  • Given: 2HgO(s) + 182 \text{ kJ} \rightarrow 2Hg(l) + O_2(g)
  • The forward reaction (decomposition) is endothermic (energy is added to the reactant side).
  • To find the heat of formation of mercury oxide, flip the reaction: 2Hg(l) + O_2(g) \rightarrow 2HgO(s)
  • The reverse reaction (formation) will be exothermic, so the energy is released/lost.
  • Calculation: If \Delta H = 82 \text{ kJ} for 2 moles of HgO, then x = -91 \text{ kJ/mol} for one mole by making and dividing.

Lab: Cooling Water with Ice

• The mass of ice that melts into water is critical.
• Sources of error can be discussed in lab write-ups.

Procedure

• Weigh the empty cup; note the data and ensure this value is used through all calculations for this run.
• Add water to the cup (ideally around 100 mL to avoid exceeding balance capacity).
• Record the initial temperature of the water.
• Add ice to find out how much energy it takes to melt the ice.

Technique: Adding Ice

• Goal: Cool water from room temperature down to approximately 0°C; the ice will melt and the temperature of the water will lower.
• Avoid extreme slowness to prevent the air from interfering in the procedure.
• Avoid extreme haste which can also skew results.
• Add multiple pieces of ice at once (e.g. four or five).
• Swirl the mixture continuously.
• Add more ice chips to maintain a steady temperature decrease (a couple of degrees every few seconds).

Removing Excess Ice

• Remove excess ice with your hands.
• Shake access water
• Simultaneous balancing of errors occurs by the nature of this process.

Calculations and Considerations

• Water that started at room temperature will change temperature.
• Final mass will tell you the amount of ice melted.
• The mass of water that's changing temperature is the intial mass, prior of adding the ice.
• Calculate energy loss using q = mc\Delta T , where:
  • m is the initial mass of the water.
  • \Delta T is the temperature change of the water (from room temperature to 0°C).
• Calculate the mass difference to determine how much ice melted, using this to determine the mass of ice that changed state (melted).

Exam Information

• The exam begins and ends on Friday the 30th and Monday the 2nd.
• The exam is broken into three parts.
• The first two parts will be on Friday.