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LOGO International University of Sarajevo Lecture 5: Energy changes in Chemical reactions Course lecturer : Jasmin Šutković 1th April 2015.

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Presentation on theme: "LOGO International University of Sarajevo Lecture 5: Energy changes in Chemical reactions Course lecturer : Jasmin Šutković 1th April 2015."— Presentation transcript:

1 LOGO International University of Sarajevo Lecture 5: Energy changes in Chemical reactions Course lecturer : Jasmin Šutković 1th April 2015

2 Contents International University of Sarajevo 1. Thermochemistry 2.Energy and work 3.Enthalpy 4.Calorimetry 5.Enegy sources and the environment 6.Unit conversations ( tutorials )

3 1. Thermochemistry Abranch of chemistry that describes the energy changes that occur during chemical reactions. In some situations, the energy produced by chemical reactions is actually of greater interest to chemists than the material products of the reaction.

4 2. Energy and work Energy – Takes many forms – Some can be seen or felt – Defined by its effect on matter

5 2. Energy and work Forms of energy  Thermal – form of atomic and molecular motion  Radiant – energy of light,microwaves and radio waves  Electrical – flow of electrical charged particles  Chemical – atomic energy in molecular compounds  Nuclear- energy stored in the nucleus of an atom

6 Forms of energy

7 Electrical energy, nuclear energy, and chemical energy are different forms of potential energy ( PE ) Energy stored in an object because of its relative position or orientation (energy that is waiting to be released)  Kinetic energy ( KE ) – Energy due to the motion of an object (released energy ) Potential and kinetic energy

8  Energy can be converted from one form to another or transferred from one object to another.  Law of conservation of energy — total amount of energy in the universe remains constant.  Energy cannot be created or destroyed.

9 Heat, Energy and Work Work – One definition of energy is the capacity to do work. – Easiest form of work to visualize is mechanical work, which is the energy required to move an object a distance d when opposed by a force F: work = force  distance or w = F x d – Because F, the force that opposes the action, is equal to the mass of the object times its acceleration, the preceding equation becomes work = mass  acceleration  distance or w = m x a x d

10 Heat, Energy and Work  Heat (q) – Heat is thermal energy that can be transferred from an object at one temperature to an object at another temperature – Net transfer of thermal energy stops when the two objects reach the same temperature.

11 Heat, Energy and Work Energy – Energy of an object can be changed by only the transfer of energy to or from another object in the form of heat, work performed on or by the object, or some combination of heat and work. – Sum of heat produced and the work performed equals the change in energy, ΔE: energy change = heat + work or ΔE = q + w – Energy is an extensive property of matter.

12  Kinetic energy, KE, of an object is related to its mass, m, and velocity v: KE = ½ m x v 2  Potential energy is equivalent to work: PE = F x d = m x a x d = m x g x h m = mass, g = gravity (9.81 m/s 2 ) and h = height Kinetic and Potential Energy

13 Units of energy – Most common units of energy 1. S  unit of energy is the joule (J), defined as 1 (kilogrammeter 2 )/second 2, energy is also expressed in kilojoules (1 kJ = 10 3 J). 2. Non-S  unit of energy is the calorie where 1 calorie (cal) is the amount of energy needed to the raise the temperature of 1 g of water by 1°C. – One cal = J or 1J = cal. – Units of energy are the same, regardless of the form of energy. Kinetic and Potential Energy

14 To study the flow of energy during a chemical reaction, one must distinguish between the system and the surroundings. – System — the small, well-defined part of the universe in which we are interested in studying (such as a chemical reaction) – Surroundings — the rest of the universe (including the container in which the reaction is carried out) 3.Enthalpy

15  Three kinds of systems 1. Open system — can exchange both matter and energy with its surroundings 2. Closed system — can exchange energy but not matter with its surroundings 3. Isolated system — exchanges neither energy nor matter with the surroundings; total energy of the system plus the surroundings is constant Enthalpy cont....

16 State of a system — complete description of the system at a given time, including its temperature and pressure, the amount of matter it contains, its chemical composition, and the physical state of the matter State function — current property of a system (present situation) – Temperature, pressure, volume, and potential energy are all state functions – Heat and work are not state functions because they are path-dependent Enthalpy cont....

17 Direction of Heat Flow  Thermochemical equations are chemical equations in which heat is shown as either a reactant or a product. – Exothermic reaction — process in which heat (q) is transferred from the system to the surroundings: q < 0 – Endothermic reaction — process in which heat is transferred to the system from the surroundings: q > 0

18 First law of thermodynamics... First Law: Energy of the Universe is Constant E = q + w q = heat. Transferred between two bodies w = work. Force acting over a distance (F x d)

19 Enthalpy of Reaction  Change in energy, Δ E, is equal to the sum of the heat (q) produced and the work performed.  The symbol E represents the change of energy states (internal energy of a system), which is the sum of the kinetic and potential energy of all of its components.  To measure the energy changes that occur in chemical reactions, a quantity called enthalpy (H) is used.  Dont mix up Heat (q) with enthalpy (H) !!!!

20 Enthalpy of Reaction cont...  Enthalpy of a system is defined as the sum of its internal energy, E, and the product of its pressure P and volume: H = E + ( P x V )  At constant pressure, the heat transferred from the surroundings to the system (or vice versa) is identical to the change in enthalpy, ΔH, of the system: q = H final –H initial = ΔH

21 Law which states that the energy change( Δ H) in an overall chemical reaction is equal to the sum of the energy changes( Δ H )in the individual reactions comprising it.  H for a process involving the transformation of reactants into products is not dependent on pathway. Therefore, we can pick any pathway to calculate  H for a reaction. Hess’s Law

22 Graphical representation....

23  For example, suppose you are given the following data: Hess’s Law example Could you use these data to obtain the enthalpy change for the following reaction?

24  If we multiply the first equation by 2 and reverse the second equation, they will sum together to become the third. Hess’s Law

25 More Examples  getics/sums.html

26 4.Calorimetry  Thermal energy cannot be measured easily.  But temperature change,caused by the flow of thermal energy between objects or substances, can be measured.  Calorimetry is the set of techniques employed to measure enthalpy changes in chemical processes using calorimeters.

27 Heat Capacity  The level of the temperature change of an object,when it absorbs or loses thermal energy, depends on both the amount of thermal energy transferred (q) and the heat capacity of the object.  Heat capacity (C) — amount of energy needed to raise the temperature of the object exactly 1°C, and the units of C are joules per degree Celsius (J/°C)  Change in temperature (ΔT) is: ΔT = q / C, ΔT = T final  T initial (in units of °C).  If ΔT and q are positive, heat flows from the surroundings into the object, and if ΔT and q are negative, heat flows from the object into the surroundings.

28 Heat Capacity types  Heat capacity of an object depends also on mass and its composition.  Molar heat capacity (C p ) — amount of energy needed to increase the temperature of 1 mol of a substance by 1°C; units of C p are J/(mol°C)  Specific heat (C s ) — amount of energy needed to increase the temperature of 1 g of a substance by 1°C, units are J/(g°C) q = n x C p x ΔT where n = number of moles of substance q = m x C s x ΔT where m = mass of substance in grams

29 5.Energy Sources and the Environment  Society requires the constant expenditure of huge amounts of energy.  The United States consumes 30% of the world’s total energy usage.

30  Energy can never actually be consumed but can be changed from one form to another. It is more efficient to use primary sources of energy directly than to transform them to a secondary source. More than 80% all energy is provided by the combustion of fossil fuels: oil, coal, and natural gas. Natural gas and petroleum are the preferred fuels due to their availability, ease of transport, and facile conversion to convenient fuels. Fuels A fuel is any substance that is burned to provide heat or other forms of energy.

31 Coal – A complex solid material derived primarily from plants that died and were buried long ago and were subsequently subjected to high temperatures and pressures – Contains a large number of oxygen atoms that link parts of the structure together, in addition to the basic framework of carbon-carbon bonds with the presence of many different kinds of ring structures Fuels

32  Coal – Four distinct classes of coal 1. Lignite, high oxygen content, has the lowest enthalpy of combustion 2. Anthracite, has the highest enthalpy of combustion, and is the highest grade of coal 3. Bituminous — most abundant form, has a high sulfur content 4. Subbituminous — han H/C ratio of 0.9 and as has a high oxygen content Fuels

33 Converting coal to gaseous and liquid fuels – Oil and natural gas resources are limited – Coal is abundant but is difficult to transport and mine Current research focuses on developing methods to convert coal to gaseous fuels or liquid fuels One method is to react coal with steam to produce a mixture of CO and H 2, known as synthesis gas or syngas Syngas is also used as a reactant to produce methane and methanol Fuels

34  Burning of carbon-based fuels is done on a vast scale and produces large amounts of CO 2.  The amounts of CO 2 are overwhelming the natural ability of the planet to remove CO 2 from the atmosphere, and the elevated levels of CO 2 are affecting the temperature of the planet through the “greenhouse effect.”  Therfore the Carbon dioxide is a greenhouse gas, and the increased concentration of carbon dioxide in the atmosphere influence earth's radiation balance! The Carbon Cycle and the Greenhouse Effect

35 Greenhouse gases  Greenhouse gases are those that can absorb and emit infrared radiation, but not radiation in or near the visible spectrum. In order, the most abundant greenhouse gases in Earth's atmosphere are:  water vapour (H 2 O)  carbon dioxide (CO 2 )  methane (CH 4 )  nitrous oxide (N 2 O)  ozone (O 3 )

36 – The distribution and flow of carbon throughout the planet – Atmospheric CO 2 either dissolves in the oceans and eventually precipitates as carbonate rocks or is taken up by plants – Combustion of fossil fuels releases large amounts of CO 2 that upset the balance of the carbon cycle and result in a steady increase in atmospheric CO 2 levels The global carbon cycle

37 – CO 2 is a greenhouse gas that absorbs heat before it can be radiated from the earth into space – CO 2 in the atmosphere can result in increased surface temperatures (the greenhouse effect) – Temperature increases caused by increased CO 2 levels lead to nature disbalance and weather change The atmospheric greenhouse effect

38  Lecture 5 : Book pages : – follow this power point lecture while reading the book for details !!! Book pages info

39 Periodic system of elements

40 Homework 1  Strong acid and bases are good electrolytes, why ?  Products of a neutralization reaction have none of the properties of an acid or a base. Explain why?  What is the pH of the tape water? Explain why?  What is ΔE and what is ΔH, explain their correlation? Please do hand writing and use your own language as much as possible! SUBMIT BY 8 th APRIL – NO LATE HOMEWORKS ARE ACCEPTED

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