Energetics: Exothermic and Endothermic Reactions

Okay, here is the converted markdown format of the document: #Enthalpy, Energetics and Reactions ### 19 Energetics Some chemical reactions produce heat. Others need to be heated constantly to make them occur at all. This chapter explores some examples of both kinds of reaction and examines how ene...

Okay, here is the converted markdown format of the document: #Enthalpy, Energetics and Reactions ### 19 Energetics Some chemical reactions produce heat. Others need to be heated constantly to make them occur at all. This chapter explores some examples of both kinds of reaction and examines how energy changes during reactions can be measured by experiments or calculated using bond energies. ### Learning Objectives * Know that chemical reactions in which heat energy is given out are described as **exothermic**, and those in which heat energy is taken in are described as **endothermic**. * Describe simple calorimetry experiments for reactions such as combustion, displacement, dissolving and neutralisation. * Calculate the heat energy change from a measured temperature change using the expression $\text{Q = mcΔT}$ * Calculate the molar enthalpy change $(\Delta H)$ from the heat energy change, $Q$ **Practical:** Investigate the temperature changes accompanying some of the following types of change:Salts dissolving in water, neutralisation reactions, displacement reactions and combustion reactions. **Chemistry Only:** Draw and explain energy level diagrams to represent exothermic and endothermic reactions. Know that bond breaking is an endothermic process and that bond making is an exothermic process. Use bond energies to calculate the enthalpy change during a chemical reaction. **Did you know?** Calcium oxide is known as **quicklime**. Adding water to it is described as **slaking** it. The calcium hydroxide produced is known as **slaked lime**. **Exothermic Reactions** Some chemical reactions give out energy in the form of heat. A reaction that gives out heat to the surroundings is said to be **exothermic**. If you are holding a test-tube in which an exothermic reaction is occurring, you will notice that the test-tube gets warmer. An example of an exothermic reaction is adding water to calcium oxide. If you add water to solid calcium oxide, the heat produced is enough to boil the water and produce steam. Calcium hydroxide is produced. $\text{CaO(s) +H2O(l) → Ca(OH)2 (s)}$ In an exothermic reaction, the products of the reaction have less *(chemical)* energy than the reactants. In the reaction, chemical energy *(stored in the bonds of chemicals)* is converted to heat energy, which is released to the surroundings. The temperature of the reaction mixture and its surroundings goes up. **Safety Note:** It is less hazardous to use a lump of calcium oxide, rather than powered calcium oxide. Figure 19.3 is a diagram demonstrating an exothermic reaction converting chemical energy into heat energy and causing the temperature to rise. Figure 19.4 shows the burning of hydrogen in a use case of equipment being used underwater. Any reaction that produces a flame is exothermic. Burning things produces heat energy. **Combustion Reactions** For instance, hydrogen burns in oxygen, producing water and lots of heat: $2H_2(g) + O_2(g) \rightarrow 2H_2O(l)$ Apart from burning, other exothermic changes include: * the reactions of metals with acids * neutralisation reactions * displacement reaction When magnesium reacts with dilute sulfuric acid, for example, the mixture gets very warm: $Mg(s) + H_2SO_4(aq) \rightarrow MgSO_4 (aq) + H_2(g)$ **Reminder**: This reaction is described in detail in Chapter 14 (pages 152-153). **Neutralisation Reactions** About the only interesting thing that you can observe happening when sodium hydroxide solution reacts with dilute hydrochloric acid is that the temperature rises: $NaOH(aq) + HCl(aq) \rightarrow NaCl(aq) + H_2O(l)$ **Reminder**: You can read about this reaction in Chapter 16 (page 170). We will investigate how much heat energy is released in a typical neutralisation reaction later in this chapter (pages 217-219). **Displacement Reactions** The thermite reaction between powdered aluminium and iron(III) oxide is a displacement (competition) reaction. This reaction releases a large amount of heat, which can be used in railway welding: $2Al(s) + Fe_2O_3(s) \rightarrow 2Fe(I) + Al_2O_3(S)$ **Enthalpy Change of Reaction** **Key point**:$\Delta H$ is pronounced 'delta H'. The Greek letter $\Delta$ is used to mean 'change in'. $\Delta H$ means 'change in heat'. Note: it is not possible to measure how much enthalpy ($H$) something has - you can only measure the change in enthalpy ($\Delta H$) when it reacts. **Reminder**: The mole is a unit for the amount of a substance. You can read more about it in Chapter 5. **Key Point**: The units for $\Delta H$ can be written as $kJ/mol$ or $kJ mol^{-1}$. **Key Point**: The term stability is usually used to describe the relative energies of the reactants and the products in a chemical reaction. The more energy a chemical has, the less stable it is. **Reminder**: Remember that in a chemical reaction, the reactants are the chemicals you start with. You can measure the amount of heat energy taken in or released in a chemical reaction. It is called the **enthalpy change** of the reaction and is given the symbol $\Delta H$. The enthalpy change is the amount of heat energy taken in or given out in a chemical reaction. It is the difference between the energy of the products and the energy of the reactants. $\Delta H$ is given a minus or a plus sign to show whether heat is being given out or adsorbed by the reaction. You always look at it from the point of view of the reactants. For an exothermic reaction, $\Delta H$ is given a negative number because the reactants are losing energy as heat. That heat is transferred to the surroundings, which then get warmer. $\Delta H$ is measured in units of kJ/mol (kilojoules per mole). In an equation, the amount of heat given out or taken in can be shown as, for example: $Mg(s) + H_2SO_4(aq) \rightarrow MgSO_4(aq) + H_2(g) \quad \Delta H = -466.9 \text{kJ/mol}$ The $\Delta H$ written next to an equation represents the enthalpy change of the reaction, i.e. 466.9 kJ of heat is given out when one mole of magnesium reacts in this way. You know heat has been given out because $\Delta H$ has a negative sign. **Chemistry Only** **Showing an exothermic change on an energy level diagram** In an exothermic reaction, the reactants have more (chemical) energy than the products; we say that the products are more stable than the reactants. As the reaction happens, energy is given out in the form of heat. That energy warms up both the reaction itself and its surroundings. Figure 19.5 is a diagram demonstrating an exothermic reaction showing the progress of the reaction and how heat is evolved. Figure 19.6 is of a boy using a cold pack to relieve pain in his elbow. **Endothermic Reactions** A reaction that absorbs heat from the surroundings is said to be **endothermic**. If you hold a test-tube in which an endothermic reaction is occurring you will notice that it gets colder. In an endothermic reaction, the products have more (chemical) energy than the reactants. In order to supply the extra energy that is needed to convert the reactants (lower energy) into the product **Reminder**: Not all reactions that have to be heated are endothermic reactions. Sometimes we just heat something to make it go faster. Needs to be absorbed from the surroundings. This heat energy is converted to chemical energy *energy stored in the bonds of chemicals*. The transformation of the reaction mixture and surrounding temperature. Figure 19.7 In an endothermic reaction heat energy. Has to be absorbed from the surroundings. 19.8 depicts an endothermic change and plots the energy against the reaction. You have seen the thermal decomposition of metal carbonates before in Chapter 13. These are examples of endothermic reactions. You have to heat a carbonate constantly to make it decompose. For example, copper(II) carbonate (green) decomposes on heating to produce copper(II) oxide (black). $CuCO_3(s) \rightarrow CuO(s) + CO_2(g)$ Similarly, zinc carbonate decomposes to form zinc oxide when heated. $ZnCO_3(s) \rightarrow ZnO(s) + CO_2(g)$ **Chemistry Only** **Showing an endothermic change on an energy level diagram** In an endothermic change, the products have more energy than the reactants so we say that the products are less stable than the reactants. That extra energy has to come from somewhere, and it is taken from the surroundings. In the case of the thermal decomposition of carbonates in the laboratory, it comes from the Bunsen burner. Because the reactants are gaining energy, the enthalpy change of the reaction $\Delta H$ is given a positive sign. For example: $CaCO_3(s) \rightarrow CaO(s) + CO_2(g) \quad \Delta H = +178 \text{ kJ/mol}$ This means that 178 kJ of heat energy must be absorbed to convert 1 mole of calcium carbonate into calcium oxide and carbon dioxide. **Measuring Enthalpy Changes of Reactions** Here we discuss how we can measure how much heat is taken in or given out by a chemical reaction, in the other words the enthalpy change. The technique that we use to do this is called calorimetry. **Specific Heat Capacity** **Key Point**: The unit for specific heat capacity is J/g/°C or J g-1°C-1. **Key Point**: Again, we are using the delta $(\Delta)$ symbol to indicate a change, but this time a temperature change. The concept of specific heat capacity also applies to the cooling of a substance. For example, to cool 1 g of water by 1 °C we need to take out 4.18J of energy. When we heat something up, it gets hotter. The **specific heat capacity** tells us about how much energy has to be put in to increase the temperature of something. The specific heat capacity of a substance is defined as the amount of heat needed to raise the temperature of 1 gram of a substance by 1 °C. For water, the value is 4.18 J/g/°C (joules per gram per degree Celsius). This means that 4.18 J of heat energy is needed if we want to increase the temperature of 1 g of water by 1 °C. If you want the temperature of 1 g of water to go up by 2 °C, then 4.18 x 2 = 8.36 J of heat energy must be supplied. If now you have 2 g of water, then 2 x 8.36 J of energy would be needed to raise the temperature by 2 °С. The amount of heat energy required is directly proportional to the mass *(m)* and the temperature change ($\Delta$T) of the substance. The following equation can be used to calculate how much heat energy needs to be supplied to raise the temperature of mass m by $\Delta$T°C: heat energy change = mass x specific heat capacity x temperature change $Q = m \times c \times \Delta T$ **Calorimetry Experiments for Determining the Enthalpy Changes of Reactions** It is fairly uncomplicated to measure the amount of heat absorbed or given out in several kinds of chemical reactions and physical changes. The technique used to do this is called **calorimetry** and it is based on the idea that if we use the heat from a reaction to heat another substance, such as water, we can then use the equation introduced above ($Q = mc\Delta T$) to calculate the amount of heat released. Here the mass, the specific heat capacity and the temperature change are all referring to the substance heated. If we know how many moles of reactants are used in the reaction, we can then work out the **molar enthalpy change**, $\Delta H$, of the reaction in the unit kJ/mol. The following activity illustrates how we can use calorimetry to determine the molar enthalpy change of combustion of an alcohol, i.e. how much heat energy is released when 1 mole of alcohol burns. **Activity 1** **Practical: Measuring Enthalpy Changes in Combustion Reactions** **Safety Note:** Wear eye protection. Do not carry a lit spirit burner and do not fill or re-fill one when there is a naked flame nearby. **Key point**: A calorimeter is simply something that we do a calorimetry experiment in; here it is a copper can, in other experiments it can be a polystyrene cup. One of the most common calorimetry experiments at International GCSE is to measure the amount of heat given off when a number of small alcohols are burned. You could use methanol, ethanol, propan-1-ol and butan-1-ol (see Chapter 22 for an introduction to the naming of alcohols). The alcohols are burned in a small spirit burner, and the heat produced is used to heat some water in a copper can (the calorimeter). The following procedure could be used: * Measure 100 $cm^3$ of cold water using a measuring cylinder and transfer the water to a copper can. * Take the initial temperature of the water. * Weigh a spirit-burner containing ethanol with its lid on. The lid should be Kept on when the wick is not lit to prevent Arrange the apparatus as shown in Figure 9.9 so that the spirit-burner can be used to heat the water in the copper can. The apparatus is shielded as far as possible to prevent draughts. Light the wick to heat the water. Stop heating when you have a reasonable temperature rise of water (say, about 40). The flame can be extinguished by putting the lid back on the wick. Stir the water thoroughly and measure the maximum temperature of the water. Weigh the spirit burner again with is lid on. The experiment can be repeated with the same alcohol to check The alcohol has been tested for other alcohols. Figure 19.9 is a labelled diagram example experiment that uses a calorimeter and a temperature thermometer to measure change in enthalpy. **Sample Data** | | | | :-------------------------- | :------ | | volume of water/$cm^3$ | 100 | | Mass of burner before experiment/g| 137.330 | | Mass of burner after experiment/g | 136.58 | | Original temperature of water C | 21,5 | | Final temperature of water C | 62.8 | Table 19.1 is shows results from a chemical calorimetry experiment Combustion is exothermic so the water goes up. As total mass of burner and ethanol goes down we can use to determine how much heat is released **Calculations for Activity 1** We are going to use the equation $Q = mc\Delta T$, so we need to find out what each Quantity is. Temperture change of water = $\Delta T = 62.8 - 21.5 = 41.3$ C Mass of water being heated = $m = 100g$, the density of water is 1 cm³ c is the specific heat capacity of water, $C=4.18 J/g/C$ Heat Gained by the water = $Q = mc\Delta t = 100 * 4.18 * 41.3 = 1780J$ Divide Q by 1000 to get energy in $kJ= 17kJ$ **Key Point**: This means 17.8kJ of heat energy is released by the combustion of enthanol to calculate it if one mole if one mole of enthenol. **Hint**: We quoted the final answer to 3 significant figures best is to quote answers to 3 significant figures and and base our answers based on those figures. In some calculations perform the work and then show the results from calacator with end result. **enthenol formula $C_2H_5OH$: Number if moles $N = \frac {mass} {$ Relative molecular mass Molar enthalpy change of combustion What is it? = - 1020 The negative sign shows that this has been released and the combustion Is endothermic **Hint**: Always remember in the sign of if its exothermic and andthermic How accurate do you want to calculate this what are the factors affecting it like heat and it being affected by equipment Major for alcohol. We will talk more in chapter 23. **extension work** Can be only used that are gas. we can also work if we get an other measure the amount or released in