* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Download Chemistry and the material world
Heat exchanger wikipedia , lookup
Equipartition theorem wikipedia , lookup
Thermal radiation wikipedia , lookup
Temperature wikipedia , lookup
Copper in heat exchangers wikipedia , lookup
Thermoregulation wikipedia , lookup
Calorimetry wikipedia , lookup
R-value (insulation) wikipedia , lookup
Countercurrent exchange wikipedia , lookup
Heat capacity wikipedia , lookup
Conservation of energy wikipedia , lookup
Heat equation wikipedia , lookup
Heat transfer wikipedia , lookup
Internal energy wikipedia , lookup
Thermal conduction wikipedia , lookup
First law of thermodynamics wikipedia , lookup
Heat transfer physics wikipedia , lookup
Second law of thermodynamics wikipedia , lookup
Chemical thermodynamics wikipedia , lookup
Thermodynamic system wikipedia , lookup
Chemistry and the material world 123.102 Unit 4, Lecture 2 Matthias Lein Thermodynamics Definitions of Work, Heat and Energy adiabatic and non-adiabatic paths state functions internal energy the first law of thermodynamics heat capacity enthalpy – energy at constant pressure Standard enthalpy changes Hess' law and its applications Thermodynamic builds on Work, Heat and Energy: E.g.: Work is done, when a process can be used to change the height of a weight somewhere in the surrounding. To measure the amount of work we use the definition: work = force × way If a builder pulls a bucket of bricks with a pulley from the ground floor to the roof he has to exert a certain force. The higher he pulls the bucket, the more work is done. Energy is the ability to perform work If we do work on an otherwise isolated system, the system's ability to perform work itself is increased i.e. its energy is increased. If a system does work, its energy is decreased because the system's ability to perform work is decreased. Experiments have shown that work is not the only way to change a system's energy. If there is a temperature difference between the system and its surrounding, the system's energy will change when the temperatures start to equilibrate. If the energy of a system changes because of a temperature difference we speak of heat-flow. Systems which allow heat-flow are called diathermic, systems which do not allow heat-flow are called adiabatic. Thermos flasks are good approximations to adiabatic systems if the heat-flow is very slow compared to the time-scale of the experiment. Heat-flow is usually measured by observing temperature changes in the surrounding. Often in a calorimeter. The first law of thermodynamics: If a system is changed from one state to another in an adiabatic but otherwise arbitrary way, the work is always the same irrespective of the method used. If we assume a system goes from an initial state with the internal energy Ui to a final state with the internal energy Uf, it could be also assumed that the work wad for this process depends on the way the system takes from the initial state to the final state and which intermediate states are visited. But: The first law of thermodynamics states that the work is always: wad = Uf – Ui Important: The “height”-differences are independent of the way you take, they only depend on the “height” of the initial state and the “height” of the final state. This “height” is called the internal energy of the system U Any function (like U) that only depends on the difference between an initial and a final state and not on the path taken is called a state-function. Let us look at a another process: If we assume a system going from the same initial state to the same final state, but now along a non-adiabatic way. That means that the system is in thermal contact with its surrounding and heat is allowed to flow. The energy difference ΔU will be the same (because it depends only on the state and not the way) but the work done doesn't necessarily have to be the same. The difference between the energy difference ΔU and the work w is then called the heat q: q = ΔU - w This way the first law of thermodynamics gives us a fundamental definition of heat because we can measure ΔU for a process that takes s system from an initial to a final state by measuring wad for the adiabatic path and w for the non-adiabatic path. q = wad – w Finally, from the first law of thermodynamics also follows that the internal energy of an isolated system cannot change. Because for an isolated system there is w = 0 and q = 0 and with ΔU = q + w it follows that ΔU = 0. The state of an isolated system can undergo change but the energy remains the same. The energy of an isolated system is constant Sign convention: For the internal energy: ΔU = q + w we will use the sign for q and w in the following way If the internal energy is decreased by work being done by the system or heat flowing out of the system w (and or q respectively) will be negative. If the internal energy is increased by work being done to the system or heat flowing into the system w (and or q respectively) will be positive. More conventions: – SI unit of energy ● Joule ● 1 J = 1 kg m2 s-2 ● 1 kJ = 103 J – Thermodynamic equations require the temperature in kelvin ● Temperature difference (ΔT) of 1 K is numerically equal to ΔT of 1 °C Compression and expansion work: The work done if a piston moves against a constant outside pressure is given by: w = - pout ΔV ΔV is positive for an expansion of gas (Vfinal – Vinitial) work is negative for this process. Note that the work is only determined by the outside pressure even though it is the inside pressure that drives the piston. ● Heat Capacity – – – – Heat and temperature are not the same thing Heat is a transfer of energy due to a temperature difference q = CΔT q - heat (J) C - heat capacity (J K-1) ● Heat Capacity – – – Heat capacity depends on the size of the sample A property with a value that depends on the size of the sample is an extensive property A property with a value independent of the size of the sample is an intensive property ● Heat Capacity – Divide heat capacity (extensive property) by the mass of the sample to form specific heat capacity (intensive property) – c - specific heat capacity (J g-1 K-1) Divide by amount instead of mass to form molar heat capacity (J mol-1 K-1) – ● Heat Capacity q = cmΔT If a gold ring with a mass of 5.50 g changes in temperature from 25.0 to 28.0 °C, how much heat has it absorbed? m = 5.50 g c = 0.129 J g–1 K–1 ΔT = 3 K q = cmΔT = (0.129 J g-1 K-1) × (5.50 g) × (3 K) = 2.1 J • The determination of heat – Calorimeter • Apparatus designed to minimise heat loss between the system and surroundings – Bomb calorimeter • System remains at constant volume • ΔU = q + w • ΔU = qv ● Enthalpy: the heat of reaction at constant pressure ΔU = q + w ΔU = qp – pΔV – – Inconvenient as need to know ΔV Define a new thermodynamic property called enthalpy (H) ΔH = qp ● Enthalpy: the heat of reaction at constant pressure H = U + pV ΔH = ΔU + pΔV – Substituting ΔU = qp – pΔV gives ΔH = qp – pΔV + pΔV ΔH = qp ● Enthalpy: the heat of reaction at constant pressure – – – – – The heat of reaction at constant pressure is equal to ΔH The heat of reaction at constant volume is equal to ΔU ΔH > 0 reaction is endothermic ΔH < 0 reaction is exothermic The difference between ΔH and ΔU for a reaction is pΔV ● Standard enthalpy change N2(g) + 3H2(g) 2NH3(g) ΔHθ = –92.38 kJ – – – The above is a thermochemical equation Always gives the physical states of the reactants and products Its value of ΔHθ is only true when coefficients of reactants and products are numerically equal to the number of moles of the corresponding substances ● Hess’s law – – Method for combining known thermochemical equations in a way that allows us to calculate ΔHθ for another reaction One step C(s) + O2(g) CO2(g) ΔHθ = –393.5 kJ – Two step Step 1: C(s) + ½O2(g) CO(g) ΔHθ = – 110.5 kJ Step 2: CO(g) + ½O2(g) CO2(g)ΔHθ = – 283.0 kJ ● Hess’s law 2Fe(s) + 3CO2(g) Fe2O3(s) + 3CO(g) ΔHθ = +26.7 kJ 3CO(g) + O2(g) 3CO2(g) ΔHθ = -849.0 kJ 2Fe(s) + O2(g) Fe2O3(s) ΔHθ = -822.3 kJ – Rules for manipulating thermochemical equations: 1. When an equation is reversed the sign of ΔHθ must also be reversed. Fe2O3(s) + 3CO(g) 2Fe(s) + 3CO2(g) ΔHθ = -26.7 kJ 2Fe(s) + 3CO2(g) Fe2O3(s) + 3CO(g) ΔHθ = +26.7 kJ ● Hess’s law 2Fe(s) + 3CO2(g) Fe2O3(s) + 3CO(g) ΔHθ = +26.7 kJ 3CO(g) + O2(g) 3CO2(g) ΔHθ = -849.0 kJ 2Fe(s) + O2(g) Fe2O3(s) ΔHθ = -822.3 kJ – Rules for manipulating thermochemical equations: 2. Formulae can be cancelled from both sides of an equation only if the substance is an identical physical state. ● Hess’s law 2Fe(s) + 3CO2(g) Fe2O3(s) + 3CO(g) ΔHθ = +26.7 kJ 3CO(g) + O2(g) 3CO2(g) ΔHθ = -849.0 kJ 2Fe(s) + O2(g) Fe2O3(s) ΔHθ = -822.3 kJ – Rules for manipulating thermochemical equations: 3. If all the coefficients of an equation are multiplied or divided by the same factor, the value of ΔHθ must likewise be multiplied or divided by that factor. CO(g) + ½O2(g) CO2(g) ΔHθ = -283.0 kJ 3CO(g) + 3/2O2(g) 3CO2(g) ΔHθ = -849.0 kJ ● Standard enthalpy of combustion – ΔcHθ – Enthalpy change at temperature T when 1 mole of a substance is completely burned in pure oxygen gas – Combustion reactions are always exothermic – ΔcHθ always negative – kJ mol–1 ● Standard enthalpy of formation (ΔfHθ) Enthalpy change when 1 mole of substance is formed at 105 Pa and the specified temperature from its elements in their standard states – An element is in its standard state when it is in its most stable form and physical state at 105 Pa and the specified temperature – ΔfHθ for the elements in their standard states are 0 – ● Standard enthalpy of formation – – aA + bB cC + dD ΔH θ c Δf HCθ d Δf H Dθ a Δf H Aθ b Δf H Bθ Hess’s law equation Use either enthalpies of combustion or enthalpies of formation ● Bond enthalpies – – A bond enthalpy is the enthalpy change on breaking 1 mole of a particular chemical bond to give electrically neutral fragments Atomisation enthalpy (ΔatH) is the enthalpy change that occurs on rupturing all the chemical bonds in 1 mole of gaseous molecules ● Bond enthalpies and Hess’s law ● Lattice enthalpies and Hess’s law – the Born Haber cycle – Lattice enthalpies for ionic solids calculable using Hess’s law and thermodynamic data Today we covered Definitions of Work, Heat and Energy adiabatic and non-adiabatic paths state functions – that do not depend on the path internal energy – ΔU = q + w the first law of thermodynamics – conservation of energy heat capacity – q = CΔT to connect heat flow and temperature enthalpy – energy at constant pressure ΔH = ΔU + pΔV = qp Standard enthalpy changes Hess' law and its applications