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Chemistry 100(02) Fall 2013

Chemistry 100(02) Fall 2013. Instructor: Dr. Upali Siriwardane e-mail : upali@coes.latech.edu Office : CTH 311 Phone 257-4941 Office Hours : M,W, 8:00-9:30 & 11:30-12:30 a.m Tu,Th,F 8 :00 - 10:00 a.m.   Or by appointment Test Dates :.

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Chemistry 100(02) Fall 2013

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  1. Chemistry 100(02) Fall 2013 Instructor: Dr. UpaliSiriwardane e-mail: upali@coes.latech.edu Office: CTH 311 Phone257-4941 Office Hours: M,W, 8:00-9:30 & 11:30-12:30 a.m Tu,Th,F8:00 - 10:00 a.m.   Or by appointment Test Dates: September 30, 2013 (Test 1): Chapter 1 & 2 October 23, 2013 (Test 2): Chapter 3 & 4 November 13, 2013 (Test 3) Chapter 5 & 6 November 14, 2013 (Make-up test) comprehensive: Chapters 1-6 9:30-10:45:15 AM, CTH 328

  2. REQUIRED: Textbook:Principles of Chemistry: A Molecular Approach, 2nd Edition-Nivaldo J. Tro - Pearson Prentice Hall and also purchase the Mastering Chemistry Group Homework, Slides and Exam review guides and sample exam questions are available online: http://moodle.latech.edu/ and follow the course information links. OPTIONAL: Study Guide: Chemistry: A Molecular Approach, 2nd Edition-Nivaldo J. Tro 2nd Edition Student Solutions Manual: Chemistry: A Molecular Approach, 2nd Edition-Nivaldo J. Tro2nd Text Book & Resources

  3. Chapter 5. Gases 6.1 Chemical Hand Warmers…………………………………………………………………..231 6.2 The Nature of Energy: Key Definitions…………………………………………………...232 6.3 The First Law of Thermodynamics: There Is No Free Lunch…………………………….234 6.4 Quantifying Heat and Work……………………………………………………………….240 6.5 Measuring for Chemical Reactions: Constant-Volume Calorimetry……………………...246 6.6 Enthalpy: The Heat Evolved in a Chemical Reaction at Constant Pressure ……………..249 6.7 Constant-Pressure Calorimetry: Measuring……………………………………………….253 6.8 Relationships Involving……………………………………………………………………255 6.9 Determining Enthalpies of Reaction from Standard Enthalpies of Formation……………257 6.1 0 Energy Use and the Environment……………………………………………………….263

  4. Chapter 6. KEY CONCEPTS: Thermochemistry

  5. Energy and Thermodynamics • Energy is a dynamic quality. • To a physicist: • Energy is the ability to effect change. • To a biologist: • Energy is responsible for growth and development of an organism. • To a chemist: • Energy is the ability to do work or transfer heat. • Energy in the form of heat is associated with molecular motions. • Energy is transferred as heat until thermal equilibriumis established. • A change in temperature (∆T) measures energy transferred. • Thermodynamics is the science of energy transfer as heat (q). • All of thermodynamics depends on the law of CONSERVATION OF ENERGY.

  6. Forms of Energy Energy - the ability to do work. Work - when a force is applied to an object. There are several types of energy: Thermal - heat Electrical Radiant - including light Chemical Mechanical - like sound Nuclear

  7. Energy units • Kinetic energy was defined as: • kinetic energy = ½ mv2 • m = mass and v = velocity. • Joule (J)- the energy required to move a 2 kg mass at a speed of 1 m/s. It is a derived SI unit. • J = kinetic energy= (2 kg) (1 m/s)2= 1 kg m2s-2 • Volume expansion work; P DV • 24.5 L atm x 101. 3 J = 2482 J • 1 L atm 1 2

  8. Energy units and unit conversion 1J = 1 kg m2/sec2 1 cal = 4.184 J 1kcal = 1 Cal thus 1 Cal = 1 kcal = 1000 cal = 4.184 kJ = 4184 J

  9. 1) What is energy?

  10. 2) What forms of energy are available in the Universe? a) e) b) f) c) g) d) h)

  11. Kinetic energy and potential energy Externalor Macroscopic Energy • Potential Energy: Energy of an object as a result of its position • Kinetic Energy: Energy of an object as a result of its motion. Internalor submicroscopic (nano-scale) Energy) • Potential Energy: Energy of an atoms or molecules as a result of its position at nano-scale • Kinetic Energy: Energy of an object as a result of motion of its atoms and molecules at nano-scale. Temperature is directly proportional to kinetic energy (thermal energy) of atoms and molecules • Total Energy = Kinetic + Potential

  12. Internal Energy • The sum of the individual energies of all nano-scale particles (atoms, ions, or molecules) in that sample • Chemical Energy: Potential energy as stored in bonds • Nuclear energy: E = 1/2mc2 • Thermal Energy: Depends on the temperature • Total Internal Energy: Depends on the type of particles, and how many of them there are in the sample

  13. 3) Explain the differences between following categories of energy. a) Kinetic Energy: b) Potential Energy: c) Macroscopic Energy: e) Microscopic or Internal energy:

  14. Law of Conservation of Energy “Energy cannot be created or destroyed in a chemical reaction.” During a reaction, energy can change from one form to another. Example. Combustion of natural gas. Chemical bonds can be viewed as potential energy. So during the reaction: 2CH4 (g) + 3O2 (g) 2CO2(g) + 2H2O (l) + thermal energy + light some potential energy is converted to thermal energy and light.

  15. First Law of Thermodynamics • the total amount of energy in the universe is a constant • the amount of heat transferred into a system plus the amount of work done on the system must result in a corresponding increase of internal energy in the system

  16. Thermodynamics Most product-favored reactions are exothermic. Often referred to as spontaneousreactions. “Spontaneous” does not imply anything about time for reaction to occur. Kinetic factors are more important for certain reactions. Will the rearrangement of a system decrease its energy? If yes, system is favored to react — a product-favored system.

  17. 4) What is thermodynamics?

  18. 5) What is first law of thermodynamics?

  19. 6) What is first law of thermodynamics?

  20. D Internal Energy (E) and the D Enthalpy (qp or H) relationship • PE + KE = Internal energy (E) • Internal energy (E) of a chemical system depends on • Number of particles • Type of particles • Temperature • The higher the T, the higher the internal energy of the chemical system. • Changes in T (∆T) are used to monitor changes in energy (∆U). • Internal energy (DE) = q (heat) + w (work or PDV) • Work (w) under standard conditions is constant. • Under constant pressure, DU = q for the w (PDV) component falls out. • Then qp = DH, where DH is the change in ENTHALPY.

  21. What is the internal energy change (DE orU) of a system? DU is associated with changes in atoms, molecules and subatomic particles Etotal= Eke + E pe + DU DE(U) = heat (q) + w (work) DE(U) = q + w DE(U) = q -P DV; w =- P DV

  22. Volume Expansion Workw = -p DV, Why is negative sign? • Work has a sign: performed (- , loss) or done on the system (+, gain) • volume expansion work: compression and expansion • compression DV = Vf -Vi ;DV is negative expansion DV = Vf -Vi ;DV is positive • compression: w = -p DV; DV = -, w is + expansion: w = -pDV; DV = +, w is -

  23. Volume Expansion Type Work w = PDV Expansion w is + compression w is - DV = Vinitial + Vfinal (increase) P Vincrease P Vinitial qp = +2kJ

  24. 6) Given DU= q + w; [DU = internal energy, q=heat change to the system and w= work involved by the system. w= -PDV (volume expansion work, DV= Vf-Vi)]. Show this equation follows first law of thermodynamics.

  25. Thermochemistry Heat changes during chemical reactions Thermochemical equation. eg. H2 (g) + O2 (g) ---> 2H2O(l) DH =- 256 kJ; DH is called the enthalpy of reaction. if DH is + reaction is called endothermic if DH is - reaction is called exothermic

  26. What exactly is DH? • Heat measured at constant pressure qp • Chemical reactions are exposed to atmosphere and are held at a constant pressure. • Volume of materials or gases produced can change. ie: work = -PDV • DE = qp + w • qp = DE + PDV; w = -PDV • DH = DE + PDV; qp = DH(enthalpy )

  27. Exothermic Reaction Reactants Energy Products Since excess energy is released, the products are more stable.

  28. Endothermic Reaction Products Energy Reactants Additional energy is required because the products are less stable.

  29. 7) Consider following chemical reactions. a) Taken place at room temperature, 25C and atmospheric pressure: 2H2 (g) + O2 (g) ---> 2H2O(l); DHrex(1) = - 572 Kj b) Taken place at room temperature, 200C and higher pressure: 2H2 (g) + O2 (g) ---> 2H2O(g); D Hrex(2) = ? kJ Compare the values of DHrex(1) to DHrex(2). Which one would be larger?

  30. 8) What is thermodynamic standard state?

  31. universe system surroundings Why is it necessary to divide Universe into System and Surrounding Universe = System + Surrounding Boundary

  32. Types of Systems • Isolated system • no mass or energy exchange • Closed system • only energy exchange • Open system • both mass and energy exchange

  33. 9) A acid/base chemical reaction is carried out in 250 ml flask with a thermometer and stirrer by mixing 50 mL 0.1 M HCl (aq) and 50 mL of 0.1 M NaOH(aq) solutions and exotermic reaction observed. Draw a diagram for this set up: • What items constitute the system? • What items constitute the surrounding? • What is releasing heat? • What is absorbing heat?

  34. 10) Describe the following systems: a) Open System: b) Closed system: c) Isolated system:

  35. Sign Convention for Energy Exchange

  36. 11) What are the sings of following changes of the system? a) Heat REMOVED FROM a reaction mixture system TO the flask: System: Surrounding: Sign: b) Heat ABSORBED a reaction FROM the surroundings: System: Surrounding: Sign:

  37. 11) What are the sings of following changes of the system? c) Work is performed ON a gas trapped inside a piston by the surroundings: System: Surrounding: Sign: d) Work done by a heat generated inside a gas in piston ON the surroundings: System: Surrounding: Sign:

  38. Freezing and Melting

  39. Measuring Temperature

  40. Thermal Energy

  41. 12) What is heat energy?

  42. Measuring thermal energy changes Thermal energy cannot be directly measured. We can only measure differences in energy. To be able to observe energy changes, we must be able to isolate our system from the rest of the universe. Calorimeter - a device that is used to measure thermal energy changes and provide isolation of our system.

  43. Heat capacity vs Specific heat Every material will contain thermal energy. Identical masses of substances may contain different amounts of thermal energy even if at the same temperature. Heat capacity. The quantity of thermal energy required to raise the temperature of an object by one degree. Specific heat.The amount of thermal energy required to raise the temperature of one gram of a substance by one degree.

  44. Specific Heats at 25oC, 1 atm Substance DH Al(s) 0.90 Br2 (l) 0.47 C (diamond) 0.51 C (graphite) 0.71 CH3CH2OH (l) 2.42 CH3(CH2)6CH3 (l) 2.23 Substance DH Fe (s) 0.45 H2O (s) 2.09 H2O (l) 4.18 H2O (g) 1.86 N2 (g) 1.04 O2 (g) 0.92 DH = specific heat, J g-1oC-1

  45. Thermo Properties of Matter Specific Heat (Cs): The amount of energy (heat as q) required to raise ONE gram of matter’s temperature by 1 oC.

  46. Calorimetry: Technique to Measure Heat Flow • Calorimetry: • A quantitative method to determine heat flow (exchange) between two systems and their surroundings • Systems can be about physical (phase) changes or • Chemical changes (reactions) • Measures change in energy (DE) under constant pressure conditions

  47. Calorimetric Relationships q = m x Sp. ht. x DT • q is the heat “lost” or “gained” and is related to: • m: the sample mass in grams • DT: the change in temperature • where DT is Tinitial – Tfinal • Cs: the specific heat (Sp. ht.) of substance • Specific heat units cal/goC or J/goC

  48. Calorimetry • Some heat from the reaction warms water; therefore: qwater = (specific heat)(water mass)(DT) • Some heat from the reaction warms the “calorimeter bomb”; therefore: qbomb= (heat capacity, J/oC)(DT) • So, total heat evolved qtotal = qwater + qbomb

  49. Coffee-Cup Calorimeter Setup Example of simple calorimeter setup

  50. Calorimetry: Measuring Heats of Reactions PROBLEM: • 1.00 gram of octane is burned in a calorimeter containing 1200.0 grams of water. The temperature raises from 25.00 oC to 33.20 oC. 1. If the heat capacity of the calorimeter is 837 J/oC, calculate the energy of this combustion reaction. 2. Calculate the energy for 2 moles of octane. Unbalanced chemical reaction: _ C8H18 + _ O2_ CO2 + _ H2O

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