Module 4: Drivers of Reactions

1. Energy Changes in Chemical Reactions

1.1. Temperature and Heat Energy

• When heat energy is added to any substance, the particles begin to move faster
  • In solids, the particles vibrate more quickly, while in liquids and gases, the particles’ velocities are higher
• Temperature is just a measurement of the average kinetic energy (or movement) of the particles in a substance
• Note: “coldness” is just a lack of heat energy. For example, 0°C is still a lot of heat energy, which is why particles still have about 273° lower to go before they hit absolute zero

1.1.1. Celcius and Kelvin

• The Celcius scale is a temperature scale based on the triple point of water

  A substance’s triple point is the pressure and temperature at which the three phases (solid, liquid, gas) coexist in thermodynamic equilibrium.

  Until 2019, the triple point of water was, by definition, 0°C. Now it’s about 0.01C, because science can’t give me a break.

• While the Celcius (or centigrade) scale is practical for everyday situations, it is limited in that it has negative values. In other words, it is not absolute.
• You might remember from Module 3 that there is another temperature scale which uses the same unit size as Celcius, but is actually absolute: the Kelvin scale
• A substance which is 0 Kelvin (0K) contains 0 heat energy, so all of the particles would cease to move.

NOTE: This has never actually been done, although we’ve gotten pretty close. The current record is about 100 pK $(100pK=0.00000000045K=4.5x\cdot10^{-10}K)$

1.1.2. Specific Heat Capacity

• Specific Heat Capacity is a measure of how much heat energy (in Joules) is required to change the temperature of 1kg of a substance by 1K
• The SI units of SHC are Joules per Kilogram per Kelvin $(Jkg^{-1}K^{-1})$
• In formulae, SHC is represented as $c$
• For example, if the same amount of heat energy is added to the same mass of copper and water, copper will heat up about 13x more than the water. This is because water has a higher specific heat capacity than copper

NOTE: The strength of Hydrogen Bonds between water molecules means that water has one of the highest SHCs of any common substance.

1.1.3. Calculating Specific Heat Capacity

$Q=mc\Delta T$

$Q=$ amount of heat energy involved $(J)$

$m=$ mass of substance $(kg)$

$c=$ specific heat capacity of the substance $(Jkg^{-1}K^{-1})$

$\Delta T=$ change in temperature $(K)$

1.1.3.1. Comparison of common SHCs

1.2. Calorimetry

• Calorimetry is a technique used to measure the energy change occuring during chemical processes

• The term is derives from “calorie,” the pre-SI unit for heat energy, now replaced by the Joule

  Joulometry probably didn’t catch on too well

• Since many chemical process occur in water (and because water has such a high SHC), calorimetry often uses water as the working fluid to absorb the energy

• The basic idea behind Calorimetry is to have a chemical reaction inside a well-insulated container

• If you measure the masses of all the substances involved, and the temperature changes which occur, then you can use $Q-mc\Delta T$ to calculate now much energy as involved in the reaction

• While this is generally initially measure per gram (or per kg), its often converted to Moles (Yes, those moles)

1.3. Energy Changes in Chemical Reactions

1.3.1. Exothermic Reactions

• Exothermic reactions are reactions which release more energy than they require
• The amount of energy involved is referred to as the $\Delta H$ (“Delta-H”) of the reaction, and is measured per mole.
• Exothermic reactions have a negative $\Delta H$
• When chemicals lose energy, the temperature in a calorimeter rises, because the energy released heats up the water in the calorimeter
  • When this happens, the energy quantity is considered negative, because the chemicals in the reaction have lost energy

  $\text{Formula: }\Delta H=E_{Products}-E_{Reactants}$

1.3.2. Endothermic Reactions

• Endothermic Reactions are the reactions which use up more energy than they produce.
• Endothermic reactions have a positive $\Delta H$
• The reaction absorbs energy from the surrounding environment, so the temperature drops

1.3.3. Heat of Solution $(\Delta H_{sol})$

• The heat of solution is the energy change that occurs when 1mol of a substance is dissolved in water.
• While the dissolving of an ionic compound in water is usually considered a physical cahgne, dissolving involves the breaking of ionic bonds within a crystal lattice.
  • Because bonds are being broken and reformed, Dissolving can be considered a chemical reaction, which can be measured by calorimetry.

1.4. Calculation Plan

Review 1: Questions

1. $824kJ$ of energy is released when one mole of $Fe(II)$ metal is converted to $Fe(III)$ Oxide. Write the chemical equation for the reaction, including the $\Delta H$ value.
2. Calculate the Molar Heat of solutions when $2.4g$ of $NaOH$ is dissolved in $100mL$ of water $(\ce{H2O})$ to increase the temperature from $22°C$ to $28.4°C$
3. What is the temperature increase for $1L$ of water $(\ce{H2O})$ when it absorbs all the energy released from the combustion of $12g$ of Acetylene gas $(\ce{C2H2})$ and the molar heat of combustion for Acetylene is $1290kJ/mol ?$

Review 1: Solutions

1. $\ce{4Fe +3O2 ->2Fe2O3},\Delta H =-3296kJ$

2. $\Delta H_{(sol)}=m\cdot c \cdot \Delta T=102.4\cdot4.18\cdot6.4$

   $n_{(NaOH)}=\frac{m}{MM}=\frac{2.4}{39.998}=0.06mol$

   $\Delta H_{(sol)}=\frac{2.7394}{0.06}=45.7kJ/mol$

3. $\text{Fuel = Acetylene: }n_{(\ce{C2H2})}=\frac{m}{MM}=\frac{1}{24.02\cdot2.016}=0.0384 mol$

   $\Delta H_c =1290 kJ/mol \rightarrow \Delta H_c =1290\cdot 0.0384=49.536kJ$

   $\Delta H_c =49536kJ\rightarrow \Delta T =\frac{\Delta H_c}{m\cdot c}=\frac{49536}{1000\times4.18}=11.85°C$

1.5. Activation Energy $(E_a)$

• Activation energy is the energy required to start a chemical reaction
• For example, to burn a piece of paper, an exteranl source of heat, such as a match, is required to start the paper burning
• However, once the reaction has started, the source of the activation energy may not be needed anymore

1.5.1. Combustion Reactions

• Combustion is the most common source of activation energy
• Combustion is observed as burning
• It always involves a “fuel” compound combining with oxygen to form “oxide” compounds of whatever elements are present in the fuel
• Combustion reactions are ALWAYS exothermic, releasing chemical energy as heat, light, and sometimes sound.
• Hydrocarbons are the most common types of fuels
  • For example, Natural Gas and LPG (Liquified Petroleum Gas) contain methane and propane
  • The methane reacts as $\ce{CH4 +O2 ->CO2 +2H2O}$
  • A similar example is petrol, which uses Octane: $\ce{C8H18 +\frac{25}{2}O2 ->CO2 +2H2O ->8Co2 +9H2O}$

  And yes, you can do equations with fractions.