Chemistry Notes - Energy Transfer and Enthalpy Change PDF

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Energy Transfer and Enthalpy Change...

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Chemistry Notes: Energy ENERGY TRANSFER Introduction You have a cup of hot coffee. That coffee’s particles are moving fast. This kinetic energy gives the coffee a high temperature. On the other hand, you have your hand whose particles are moving at a slower pace. This kinetic energy gives your hand a lower temperature. Once your hand comes into contact with the hot cup, the kinetic energy/thermal energy from the cup transfers to your hand trying to balance it out. This transfer is what we call HEAT. Thermal Energy  Thermal energy is the energy that comes from heat. This heat is generated by the movement of tiny particles within an object. The faster these particles move, the more heat is generated. Stoves and matches are examples of objects that conduct thermal energy. Temperature  The measurement of the agitation of the atoms and particles in a system. Heat  Is the transfer of thermal energy that can occur when two systems with different temperatures come into contact with each other. How does Heat transfer?  Well in physics and chemistry reactions are happening all the time without us even knowing. These reactions cause an increase in particle disorder, which is called entropy. Therefore when your hand came into contact with the hot cup, random moving particles transferred to your hand, making it feel warm. The reason for this is because thermal particles travel from warm environments to cooler environments. Random particles will continue to transfer until both objects are at the same exact temperature (want BALANCE). Heat Transfer Systems  Heat cannot be created or be destroyed, it can only be transferred or transformed. Therefore, you must be aware of its relation with the surrounding environment. There are three types of systems: Open system, Closed system and Isolated system.  An Open system is a system that is in contact with its surroundings (the air around them). Therefore it allows matter and energy to be exchanged between the two. For example, the hot cup of coffee is an open system as it is exposed to its surroundings. Therefore, even though the hand absorbed the heat, it was also released into the air.

 A Closed system does not allow exchanges in matter, however it does allow exchanges in energy. For example, you have a reaction going on in a closed beaker, since it is close matter can exit nor enter the beaker, however energy is able to be exchange (energy gets absorbed by the beaker which is in contact with air and from there are exchanges).  An Isolated system that allows no matter exchange and no energy exchange meaning it has no way to get into contact with its surrounding. For example, an isolated thermos filled with hot chocolate. Number one this thermos is closed therefore does not allow exchange in matter. Number two because it is an isolated thermos it means that what it is made of does not allow energy to be absorbed/escaped and therefore has no contact with the surrounding = no exchange in energy. Calorimeter  this is an isolated system that is used to measure heat transfer  You have a rigid reaction vessel placed in a container filled with water. There is a closed system between the water and the reaction vessel. However the whole thing is an isolated system.  How to measure heat? Well first you measure the temperature of the water before the reaction takes place. Then you have your sample in the reaction vessel. Once the reaction happens, energy will be exchanged with the water. Lastly you would measure the temperature of the water after the reaction. If temp is higher = the reaction released energy, if temp is lower = the reaction absorbed energy.

How to calculate Thermal Energy  Q= mcΔT  Q= thermal energy (in Joules) o If Q ends up being a negative number, it means that energy was released from the system to the environment o If Q is a positive number, it means that energy was absorbed by the system to the environment  m= mass of substance (in grams)  c= specific heat capacity (J/g °C) o It is the amount of energy need to make 1 g of a substance reach 1 °C o Note that different substances have different heat capacities o The specific heat capacity for water is 4.18 J/g °C  ΔT = t final – t initial (°C) Example: o Calculate the quantity of the thermal energy absorbed by a 5.00kg block of concrete (c= 2.105 J/g °C) to raise its temp from 17.1°C to 35.5 °C o Q= (5000g)(2.105 J/g °C)(35.5 °C – 17.1 °C) Q= 193 660 J = 1.94 x 10^5

How to calculate Thermal Energy with two systems  this can only be done by assuming that the heat released from one system is equal to the heat absorbed by the second system and by assuming that the reaction is happening in an isolated system  Formula: -Q1 = +Q2 or – (mcΔT)1 = + (mcΔT)2  Example: o Calculate the mass of cold water at 10°C needed to cool to 30°C and a 10g piece of glass at 95 °C needed to cool to 30°C (C of glass= 0.84) o – (10 g)(0.84 J/g °C)(30°C - 95 °C) = + (m)(4.18 J/g °C)(30°C - 10°C) o m= 6.5 g How to calculate the final temperature when not given So if you’re missing the final temperature you can find it using two systems. The reason for this is because two systems will keep on exchanging heat until they have reached the exact same temperature therefore their final temperature will be the same.  Formula: T final = (mcT initial)1 + (mcΔT initial)2 (mc)1 + (mc)2  Example: o a 500g package of frozen raspberries (c= 3.50 J/g °C) is at -4.0 °C. It is then thawed in an insulated container filled with 2kg of water initially at a temp of 40.0°C. What is the final temp? o T final = (500)(3.50)(-4.0) + (2000)(4.18)(40.0) = 32 °C (500)(3.50) + (2000)(4.18) ENERGY (ENTHALPY) CHANGE Introduction: Ethanol (CH3CH2OH) is an alcohol derived from the fermentation of plants and is increasingly being used as a biofuel. This means that ethanol is burned to produce carbon dioxide, water and energy. It’s combustion reaction is therefore: (CH3CH2OH) + O2  2CO2 + 3H2O + Energy. For every one mole of ethanol that is burned 1267 kJ of energy are released (at 25 °C). Each fuel releases a different amount of energy. For example, for every one mole of propane (C 3H8) that is burned 2220 kJ of energy are released (at 25 °C) or for every one mole of gasoline that is burned 5013 kJ of energy is released. All these different values are what we call ENTHALPY CHANGE (ΔH), more specifically molar enthalpy change. What is Enthalpy Change?

 SO enthalpy (H) is the total amount of energy present in a system. However this is very difficult to measure as it is the sum of all the different types of energy present. Therefore, we calculate enthalpy change instead (ΔH).  Enthalpy change is the amount of energy exchanged between a system and its environment during a chemical reaction or physical change (at a constant pressure).  It could also be called as heat of reaction  To find the enthalpy change you must find the energy present in the reactants (before the reaction) and then subtract it from the energy of the products (after the reaction) in order to find the change in energy. Formula: ΔH = Hp – Hr Molar Enthalpy Change (ΔH°)  This is the same thing as enthalpy change, however instead of calculating the total of amount of energy before and after, you are calculating the change of energy in one mole.  Molar enthalpy change that happens at 25 °C is known as standard molar enthalpy change and is represented by ΔH°  For example, in the introduction we know that each mole of ethanol releases 1267 kJ of energy when burned. Since it is releasing energy, we know that it is an exothermic reaction, meaning the system is losing energy (-). Therefore, it can be said that ethanol as a standard molar enthalpy change of –1267 kJ/mol, as it is losing 1267 kJ per mole burned (ΔH°= -1267 kJ/mol). Exothermic and Endothermic reactions in ΔH  Exothermic reaction: fuel + oxygen= carbon dioxide + water + Energy  This means that the enthalpy of the products will be less than the enthalpy of the reactants  Diagram: Example: Ethanol

(negative)

 Endothermic Reaction: reactant + Energy  products  This means that energy is being absorbed and therefore the enthalpy of the products will be greater than the enthalpy of the reactants.  Diagram: Example: Calcium carbonate

(positive)

Enthalpy change for physical changes  Enthalpy change can also be measured for physical change of matter  For example, if you have an ice cube (reactant) and it melts to from water (product). These are two systems that are made up of the same thing, which is H2O, however they are in two different states of matter.  This fusion of water can be represented by: H2O(solid)  H2O(liquid)  It take 6.03 kJ of energy to melt one mole of ice. Therefore it is an endothermic reaction and its diagram will look this: Note that for the reverse reaction will have the same value just different sign.

Endothermic and Exothermic reaction in physical changes During and endothermic reaction, energy is absorbed and increased the kinetic energy of the particles, causing the forces of attraction between the particles to weaken and to break (solid  liquid). Therefore, the physical reaction will result in greater separation of its particles.  During an exothermic reaction, energy is being released and the kinetic energy of the particles start to decrease, causing the forces of attraction between the particles to strengthen (liquid  solid). Therefore, Therefore, the physical reaction will result in a greater attraction of its particles. Temperature as a function of heat added

Energy Balance  Since it is impossible to know the initial and final enthalpy of a system in which a chemical reaction occurs (unless they give it to us in a test), another method

must be used to determine the enthalpy change of a reaction= ENERGY BALANCE  Energy balance is a method that compares the bonds present in the reactants and the products. Each chemical bond posses a characteristic bond energy that corresponds to the energy needed to break a bond or to the energy released to form a bond. Meaning these bonds can tell us the enthalpy change.  Formula: ΔH = ΔH bonds broken + (-) ΔH bonds formed  ΔH= enthalpy change of the reaction in kJ/mol ΔH bonds broken= enthalpy change of the breaking of reactant bonds in kJ/mol ΔH bonds formed = enthalpy change of the forming of product bonds in kJ/mol  Example: Calculate the enthalpy change of the following reaction by performing an energy balance. Indicate if the reaction is endothermic or exothermic. H2 + Cl2  2HCl o Step 1: draw the bonds: H-H + Cl-Cl  H-Cl H-Cl o Step 2:look at table for values and plug them into equation ΔH = ΔH bonds broken + (-) ΔH bonds formed ΔH = (EH-H + ECl-Cl) + (-)(2 x EH-cl) ΔH = (436 + 243) + (-)( 2 x 432) ΔH= -185 kJ for two moles (divide by 2)  -92.5 kJ/mol o Step three: see if it is exo or endo: this is an exothermic reaction (because it is negative)...