DOCSLIB.ORG

Lab.9. Calorimetry

Total Page：16

File Type：pdf, Size：1020Kb

Download full-text PDF Read full-text
Lab.9. Calorimetry Lab.6. Thermodynamics, Calorimetry Key words: Heat, energy, exothermic & endothermic reaction, calorimeter, calorimetry, enthalpy of reaction, specific heat, chemical & physical change, enthalpy of neutralization, law of conservation of energy, final temperature, initial temperature, lattice energy, hydratation energy, enthalpy of solution Literature: J. A. Beran; Laboratory Manual for Principles of General Chemistry, pp. 245-256. J.E. Brady, F. Senese: Chemistry – Matter and its Changes, 4th ed. Wiley 2003, Chapters 7, 20, . M. Hein and S. Arena: Introduction to Chemistry, 13th ed. Wiley 2011; pp. 157 - 161 J. Crowe. T. Bradshaw, P. Monk, Chemistry for the Biosciences. The essential concepts., Oxford University Press, 2006; pp. 416 - 450. J. Brady, N. Jespersen, A. Hysop, Chemistry, International Student Version, 7thed. Wiley, 2015, Chapter 18, 855 – 890. Theoretical background Accompanying all chemical and physical changes is a transfer of heat (energy); heat may be either evolved (exothermic) or absorbed (endothermic). A calorimeter (Fig. 1) is the laboratory apparatus that is used to measure the quantity and direction of heat flow accompanying a chemical or physical change. The calorimeter is well-insulated so that, ideally, no heat enters or leaves the calorimeter from the surroundings. For this reason, any heat liberated by the reaction or process being studied must be picked up by the calorimeter and other substances in the calorimeter. The heat change in chemical reactions is quantitatively expressed as the enthalpy (or heat) of reaction, H, at constant pressure. H values are negative for exothermic reactions and positive for endothermic reactions. Lab.6. Thermodynamics, Calorimetry Fig. 1. A set nested coffee cups is a good constant pressure calorimeter. Two quantitative measurements of heat are detailed in his experiment: measurements of the heat accompanying an acid-base reaction, and the heat associated with dissolution of a salt in water. Specific Heat The energy (heat, expressed in joules, J) required to change the temperature of one gram of a substance by 1⁰C is the specific heat of that substance: J energy(J ) specific heat = (1) g • C mass (g) T (C) or, rearranging for energy, J energy (J ) = specific heat mass (g) T (C) (2) g • C T is the temperature change of the substance. Although the specific heat of a substance changes slightly with temperature, for our purposes, we assume it is constant over the temperature changes of this experiment. Enthalpy (Heat) of neutralization of an Acid-Base Reaction The reaction of a strong acid with a strong base is an exothermic reaction that produces water and heat as products. Lab.6. Thermodynamics, Calorimetry Definition: Enthalpy of neutralization: energy released per mole of water formed in an acid- base reaction-an exothermic quantity. + − H 3O (aq) + OH (aq) → 2 H 2O (l) + heat The enthalpy (heat) of neutralization, Hn , is determined by (1) assuming the density and specific heats of the acid and base solutions are equal to that of water and (2) measuring the temperature change, T , when the two are mixed. H = − specific heat combined masses T (3) n H 2O acid +base Hn is generally expressed in units of kJ/mol of acid (or base) reacted. The mass (grams) of the solution equals the combined masses of the acid and base solutions. Enthalpy (Heat) of Solution for the Dissolution of a Salt Lattice energy: energy required to vaporize one mole of salt into its gaseous ions-an endothermic quantity. Hydratation energy: energy released when one mole of a gaseous ion is attracted to and surrounded by water molecules forming one mole of hydrated ion in aqueous solution-an exothermic quantity. When a salt dissolves in water, energy is either absorbed or evolved depending upon the magnitude of the salt’s lattice energy and the hydration energy of its ions. For the dissolution of KI: H2O + − KI (s) ⎯ ⎯⎯ → K (aq) + I (aq) Hs =+13 kJ/mol The lattice energy (an endothermic quantity) of a salt, HLE, and the hydration energy (an exothermic quantity), HHyd, of its composite ions account for the amount of heat evolved or absorbed when one mole of the salt dissolves in water. The enthalpy (heat) of solution, Hs, is the sum of two terms (Fig. 2): H s = H LE + H Hyd (4) Whereas HLE and HHyd are difficult to measure in the laboratory, Hs is easily measured. Lab.6. Thermodynamics, Calorimetry Fig.2. Energy changes in the dissolving of solid KI in water. The enthalpy of solution for the dissolution of a salt, Hs, is determined experimentally by adding the heat changes of the salt and water when the two are mixed. Hs is expressed in units of kilojoules per mole of salt. H = (− heat change )+ (heat change ) (5) s H 2O salt specific heat mass T specific heat mass T H = − H2O H2O H2O + − salt salt salt (6) S mole salt molesalt A temperature rise for the dissolution of a salt, indicating an exothermic process, means that the HHyd is greater than the HLE for the salt; conversely, a temperature decrease in the dissolution of the salt indicates that HLE is greater than HHyd and Hs is positive. Lab.6. Thermodynamics, Calorimetry Table 1. Specific Heat of Some Salts Salt Formula Specific Heat (J/g•⁰C) Ammonium chloride NH4Cl 1.57 Ammonium nitrate NH4NO3 1.74 Ammonium sulfate (NH4)2SO4 1.409 Calcium chloride CaCl2 0.657 Sodium carbonate Na2CO3 1.06 Sodium hydroxide NaOH 1.49 Sodium sulfate Na2SO4 0.903 Sodium thiosulfate pentahydrate NaS2O3•5H2O 1.45 Potassium bromide KBr 0.439 Potassium nitrate KNO3 0.95 Plot the Data. Plot the temperature ( y axis) versus time ( x axis) on the top half of a sheet of linear graph paper. The maximum temperature is the intersection point of two lines: (1) the best line drawn through the data points on the cooling portion of the curve and (2) a line drawn perpendicular to the time axis at the mixing time [when the acid is added to the base or when a salt is added to the water]. (Fig.3). Lab.6. Thermodynamics, Calorimetry Fig. 3. Extrapolation of temperature vs. time data (not scale) for an exothermic process. EXPERIMENTAL PART PART 1 Enthalpy (Heat) of Neutralization for an Acid-Base Reaction Calculate the enthalpy of neutralization reaction: q = m∙ c ∙ΔT = m∙ c∙ ( Tf - Ti ) ∆H = q /n, where c - specific heat of mixture, n – moles of water created Step 1 Measure 50.0 mL of 1.1 M HCl in a clean, graduated cylinder. Measure temperature with temperature sensor and record it in Table 2. Step 2 Using another clean, graduated cylinder transfer 50.0 mL of a standard 1.0 M NaOH solution to the dry calorimeter. Record the temperature and exact molar concentration of the NaOH solution in Table 2. Lab.6. Thermodynamics, Calorimetry Step 3 Carefully, but quickly add the acid to the base, replace the calorimeter lid, and swirl gently. Read and record the temperature and time every 5 seconds for 1 minute and thereafter every 15 seconds for other 2 minutes. Step 4 Calculate the thermodynamic data and fill the proper places in Table 3. Step 5 Repeat steps 1-4, for 1.1 M HNO3. Table 2. Data for acid–base neutralization HCl + NaOH HNO + NaOH 3 Trial 1 Trial 2 Trial 1 Trial 2 1. Volume of acid (mL) 2. Temperature of acid ⁰C 3. Volume of NaOH (mL) 4. Temperature of NaOH (⁰C) 5. Exact molar concentration of NaOH (mol/L) 6. Maximum temperature from graph( ⁰C) Table 3. Calculation for Enthalpy (Heat) of Neutralization for an Acid-Base Reaction HCl + NaOH HNO3 + NaOH Trial Trial Trial 1 Trial 2 1 2 Average initial temperature of acid and base ( ⁰C ) Temperature change, T ( ⁰C ) Volume of final mixture (mL) Mass of final mixture (g) (Assume the density of the solution is 1.0 g/mL) Specific heat of mixture 4.18J/g∙◦C 4.18J/g∙◦C Heat evolved (J) Amount of OH- reacted, the limiting reactant (mol) Lab.6. Thermodynamics, Calorimetry Amount of H2O formed (mol) Heat evolved per mole of H2O, ∆Hn (kJ/mol H2O) Average ∆Hn (kJ/mol H2O) Answer to the question: is the enthalpy different for both neutralization reactions, whether or is the same (in frame of experimental error)? Comment on your answer. Disposal: Rinse the calorimeter twice with deionized water beforewater before next step PART 2. Enthalpy (Heat) of Solution for the Dissolution of a Salt Calculation for Enthalpy: specific heat mass T specific heat mass T StepH 6=. − H2O H2O H2O + − salt salt salt S mole salt molesalt Prepare the salt. On weighing paper, measure about 5.0 g of the assigned salt. Record the name of the salt in Table 4. Step 7 Using your clean, graduated cylinder, add 20 g (20.0 mL) of deionized water to the calorimeter and record its temperature in Table 4. Secure the thermal sensor below the water surface. Lab.6. Thermodynamics, Calorimetry Step 8 Carefully add (do not spill) the salt to the calorimeter, replace the lid, and swirl gently. Read and record the temperature and time at 5-second intervals for 1 min and thereafter every 15 seconds for 2 min. Record the data. Step 9 Calculate thermodynamic data for salt and fill the Table 4. Step 10 Repeat steps 6-9, for other salt assigned by TA. Disposal: Discard the salt solution into the “Waste Salts” container, followed by additional tap water. Table 3. Enthalpy (Heat) of solution for the Dissolution of a Salt Name of the salt .............................. ................................... Mass of salt (g) Moles of salt Mass of water (g) Initial temperature of water ( ⁰C ) Final temperature of mixture Table 4.
Load more

Recommended publications
  • Lesson 4 Temperature Change TEACHER GUIDE
    Lesson 4 Temperature change TEACHER GUIDE Lesson summary Students meet scientist Jason Williams, an industrial chemist who designs the materials and processes for making solar cells. He explains that during the summers, Antarctic days are very long, sometimes lasting a couple of weeks or more. That makes solar energy an abundant natural resource near the South Pole. Solar cells convert energy from the sun into electrical energy. Converting energy from one form to another is an important process. In the activity, students will conduct two chemical reactions that convert chemical energy into thermal energy. Key concepts + A change in temperature is a clue that a chemical reaction may have occurred. + When the temperature increases during a chemical reaction, it is called an exothermic reaction. + When the temperature decreases during a chemical reaction, it is called an endothermic reaction. + It takes energy to break chemical bonds in the reactants. + Energy is released when chemical bonds form in the products. Safety Be sure you and the students wear properly fitting goggles. Read and follow all safety warnings on the labels of the sodium bicarbonate, citric acid, calcium chloride, and universal indicator containers. Also follow the warnings on the packaging of the foot warmer. Have students wash their hands after the activity. Proper disposal At the end of the lesson, have students pour their used solutions in a waste container. The resulting solution from the final demo should be placed in this container, too. Then dispose of all of this liquid waste down the drain or according to local regulations. The left-over sodium bicarbonate, citric acid, and calcium chloride powders can be disposed of with the classroom trash.
    [Show full text]
  • 1 Introduction
    1 Field-control, phase-transitions, and life’s emergence 2 3 4 Gargi Mitra-Delmotte 1* and A.N. Mitra 2* 5 6 7 139 Cite de l’Ocean, Montgaillard, St.Denis 97400, REUNION. 8 e.mail : [email protected] 9 10 2Emeritius Professor, Department of Physics, Delhi University, INDIA; 244 Tagore Park, 11 Delhi 110009, INDIA; 12 e.mail : [email protected] 13 14 15 16 17 18 19 *Correspondence: 20 Gargi Mitra-Delmotte 1 21 e.mail : [email protected] 22 23 A.N. Mitra 2 24 e.mail : [email protected] 25 26 27 28 29 30 31 32 33 34 35 Number of words: ~11748 (without Abstract, references, tables, and figure legends) 36 7 figures (plus two figures in supplementary information files). 37 38 39 40 41 42 43 44 45 46 47 Abstract 48 49 Critical-like characteristics in open living systems at each organizational level (from bio- 50 molecules to ecosystems) indicate that non-equilibrium phase-transitions into absorbing 51 states lead to self-organized states comprising autonomous components. Also Langton’s 52 hypothesis of the spontaneous emergence of computation in the vicinity of a critical 53 phase-transition, points to the importance of conservative redistribution rules, threshold, 54 meta-stability, and so on. But extrapolating these features to the origins of life, brings up 55 a paradox: how could simple organics-- lacking the ‘soft matter’ response properties of 56 today’s complex bio-molecules--have dissipated energy from primordial reactions 57 (eventually reducing CO 2) in a controlled manner for their ‘ordering’? Nevertheless, a 58 causal link of life’s macroscopic irreversible dynamics to the microscopic reversible laws 59 of statistical mechanics is indicated via the ‘functional-takeover’ of a soft magnetic 60 scaffold by organics (c.f.
    [Show full text]
  • Comparison of the Results from Six Calorimeters in the Determination of the Thermokinetics of a Model Reaction Wassila Benaissa, Douglas Carson
    Comparison of the results from six calorimeters in the determination of the thermokinetics of a model reaction Wassila Benaissa, Douglas Carson To cite this version: Wassila Benaissa, Douglas Carson. Comparison of the results from six calorimeters in the determina- tion of the thermokinetics of a model reaction. AIChE Spring Meeting 2011 & 7. Global Congress on Process Safety (GCPS), Mar 2011, Chicago, United States. pp.NC. ineris-00976225 HAL Id: ineris-00976225 https://hal-ineris.archives-ouvertes.fr/ineris-00976225 Submitted on 9 Apr 2014 HAL is a multi-disciplinary open access L’archive ouverte pluridisciplinaire HAL, est archive for the deposit and dissemination of sci- destinée au dépôt et à la diffusion de documents entific research documents, whether they are pub- scientifiques de niveau recherche, publiés ou non, lished or not. The documents may come from émanant des établissements d’enseignement et de teaching and research institutions in France or recherche français ou étrangers, des laboratoires abroad, or from public or private research centers. publics ou privés. Comparison of the Results from Six Calorimeters in the Determination of the Thermokinetics of a Model Reaction Wassila BENAISSA INERIS Parc Technologique Alata BP 2, F-60550 Verneuil-en-Halatte [email protected] Douglas CARSON INERIS Keywords: calorimeter, thermal runway, kinetics Abstract This paper deals with the comparison of experimental results from several types of commercially available calorimeters: a screening calorimeter (DSC), a Calvet calorimeter (C80), a reaction calorimeter (RC1), and various pseudo-adiabatic calorimeters (VSP 2, ARSST, and Phi-Tec 1). One exothermic reaction was selected as a case study: the esterification of acetic anhydride by methanol, a system which has been well studied in the literature.
    [Show full text]
  • The Use of Accelerating Rate Calorimetry (ARC) for the Study of the Thermal Reactions of Li-Ion Battery Electrolyte Solutions J.S
    Journal of Power Sources 119–121 (2003) 794–798 The use of accelerating rate calorimetry (ARC) for the study of the thermal reactions of Li-ion battery electrolyte solutions J.S. Gnanaraja, E. Zinigrada, L. Asrafa, H.E. Gottlieba, M. Sprechera, D. Aurbacha,*, M. Schmidtb aDepartment of Chemistry, Bar-Ilan University, Ramat-Gan 52900, Israel bMerck KGaA, D-64293 Darmstadt, Germany Abstract The thermal stability of 1M LiPF6, LiClO4, LiN(SO2CF2CF3)2 (LiBETI) and LiPF3(CF2CF3)3 (LiFAP) solutions in mixtures of ethylene carbonate, diethyl carbonate and dimethyl carbonate in the temperature range 40–350 8C was studied by ARC and DSC. NMR was used to analyze the reaction products at different reaction stages. The least thermally stable are LiClO4 solutions. LiPF3(CF2CF3)3 solutions showed higher thermal stability than LiPF6 solutions. The highest thermal stability was found for LiN(SO2CF2CF3)2 solutions. Studies by DSC and pressure measurements during ARC experiments with LiPF6 and LiFAP solutions detected an endothermic reaction, which occurs before a number of exothermic reactions as the temperature increases. Fluoride ions are formed and react with the alkyl carbonate molecules both as bases and as nucleophiles. # 2003 Elsevier Science B.V. All rights reserved. Keywords: Accelerating rate calorimetry (ARC); Differential scanning calorimetry (DSC); Thermal stability; Alkyl carbonate solutions 1. Introduction 2. Experimental Accelerating rate calorimetry (ARC) is an important One molar LiPF6, LiClO4, LiN(SO2CF2CF3)2 and method for studying the thermal behavior of materials LiPF3(CF2CF3)3 solutions in mixture of ethylene carbonate [1,2]. LiPF6 solutions in alkyl carbonates are widely used (EC), diethyl carbonate (DEC), dimethyl carbonate (DMC) in commercial Li-ion batteries in spite of their relatively low (2:1:2 v/v/v) were obtained from Merck KGaA (highly pure, thermal stability.
    [Show full text]
  • Chemical Reactions Involve Energy Changes
    Page 1 of 6 KEY CONCEPT Chemical reactions involve energy changes. BEFORE, you learned NOW, you will learn • Bonds are broken and made • About the energy in chemical during chemical reactions bonds between atoms • Mass is conserved in all • Why some chemical reactions chemical reactions release energy • Chemical reactions are • Why some chemical reactions represented by balanced absorb energy chemical equations VOCABULARY EXPLORE Energy Changes bond energy p. 86 How can you identify a transfer of energy? exothermic reaction p. 87 endothermic reaction p. 87 PROCEDURE MATERIALS photosynthesis p. 90 • graduated cylinder 1 Pour 50 ml of hot tap water into the cup • hot tap water and place the thermometer in the cup. • plastic cup 2 Wait 30 seconds, then record the • thermometer temperature of the water. • stopwatch • plastic spoon 3 Measure 5 tsp of Epsom salts. Add the Epsom salts to the beaker and immedi- • Epsom salts ately record the temperature while stirring the contents of the cup. 4 Continue to record the temperature every 30 seconds for 2 minutes. WHAT DO YOU THINK? • What happened to the temperature after you added the Epsom salts? • What do you think caused this change to occur? Chemical reactions release or absorb energy. COMBINATION NOTES Chemical reactions involve breaking bonds in reactants and forming Use combination notes new bonds in products. Breaking bonds requires energy, and forming to organize information on how chemical reactions bonds releases energy. The energy associated with bonds is called bond absorb or release energy. energy. What happens to this energy during a chemical reaction? Chemists have determined the bond energy for bonds between atoms.
    [Show full text]
  • 5.3 Controlling Chemical Reactions Vocabulary: Activation Energy
    5.3 Controlling Chemical Reactions Vocabulary: Activation energy – Concentration – Catalyst – Enzyme – Inhibitor - How do reactions get started? Chemical reactions won’t begin until the reactants have enough energy. The energy is used to break the chemical bonds of the reactants. Then the atoms form the new bonds of the products. Activation Energy is the minimum amount of energy needed to start a chemical reaction. All chemical reactions need a certain amount of activation energy to get started. Usually, once a few molecules react, the rest will quickly follow. The first few reactions provide the activation energy for more molecules to react. Hydrogen and oxygen can react to form water. However, if you just mix the two gases together, nothing happens. For the reaction to start, activation energy must be added. An electric spark or adding heat can provide that energy. A few of the hydrogen and oxygen molecules will react, producing energy for even more molecules to react. Graphing Changes in Energy Every chemical reaction needs activation energy to start. Whether or not a reaction still needs more energy from the environment to keep going depends on whether it is exothermic or endothermic. The peaks on the graphs show the activation energy. Notice that at the end of the exothermic reaction, the products have less energy than the reactants. This type of reaction results in a release of energy. The burning of fuels, such as wood, natural gas, or oil, is an example of an exothermic reaction. Endothermic reactions also need activation energy to get started. In addition, they need energy to continue.
    [Show full text]
  • 13 Energetics
    13 Energetics All chemical substances contain energy stored in their bonds. When a chemical reaction occurs, there is usually a change in energy between the reactants and products. This is normally in the form of heat energy, but may also be in the form of light, nuclear or electrical energy. Exothermic and endothermic reactions Based on energy changes occurring, reactions can be of two types: • An exothermic reaction produces heat which causes the reaction mixture and its surroundings to get hotter (it releases energy to the surroundings). Exothermic reactions include neutralisation reactions, burning fossil fuels and respiration in cells. • An endothermic reaction absorbs heat which causes the reaction mixture and its surroundings to get colder (it absorbs energy from the surroundings). Endothermic reactions include dissolving certain salts in water, thermal decomposition reactions and photosynthesis in plants. Breaking and forming bonds during reactions During any chemical reaction, existing bonds in the reactants are broken and new bonds are formed in the products: • Energy is absorbed when the existing bonds in the reactants are broken. • Energy is released when new bonds are formed in the products: reactants products existing bonds are broken new bonds are formed energy is absorbed energy is released • In an exothermic reaction: energy absorbed to break bonds < energy released when forming bonds The extra energy is released to the surroundings causing the temperature of the surroundings to increase. • In an endothermic reaction: energy absorbed to break bonds > energy released when forming bonds The extra energy is absorbed from the surroundings causing the temperature of the surroundings to decrease.
    [Show full text]
  • Energy and Enthalpy Thermodynamics
    Energy and Energy and Enthalpy Thermodynamics The internal energy (E) of a system consists of The energy change of a reaction the kinetic energy of all the particles (related to is measured at constant temperature) plus the potential energy of volume (in a bomb interaction between the particles and within the calorimeter). particles (eg bonding). We can only measure the change in energy of the system (units = J or Nm). More conveniently reactions are performed at constant Energy pressure which measures enthalpy change, ∆H. initial state final state ∆H ~ ∆E for most reactions we study. final state initial state ∆H < 0 exothermic reaction Energy "lost" to surroundings Energy "gained" from surroundings ∆H > 0 endothermic reaction < 0 > 0 2 o Enthalpy of formation, fH Hess’s Law o Hess's Law: The heat change in any reaction is the The standard enthalpy of formation, fH , is the change in enthalpy when one mole of a substance is formed from same whether the reaction takes place in one step or its elements under a standard pressure of 1 atm. several steps, i.e. the overall energy change of a reaction is independent of the route taken. The heat of formation of any element in its standard state is defined as zero. o The standard enthalpy of reaction, H , is the sum of the enthalpy of the products minus the sum of the enthalpy of the reactants. Start End o o o H = prod nfH - react nfH 3 4 Example Application – energy foods! Calculate Ho for CH (g) + 2O (g) CO (g) + 2H O(l) Do you get more energy from the metabolism of 1.0 g of sugar or
    [Show full text]
  • Chemical Reactions ©TCL
    Chemical Reactions ©TCL Keyword Definition Endothermic Reactions Oxidation Reactions In an endothermic reaction, thermal energy is taken in from the In an oxidation reaction, a substance gains oxygen. Metals and Endothermic Reactions that take in heat surroundings, therefore there is a temperature decrease. Thermal non-metals can take part in oxidation reactions. decomposition is an example. Metals react with oxygen in the air to produce metal oxides. For Exothermic Reactions that give out heat Exothermic Reactions example, copper reacts with oxygen to produce copper oxide In an exothermic reaction, thermal energy is given out to the when it is heated in the air. surroundings, therefore there is a temperature increase. Oxidation Reaction of other elements with oxygen Combustion, oxidation and neutralisation reactions are all examples. Copper + Oxygen → Copper Oxide 2Cu + O → 2CuO Temp decrease decrease Temp 2 → Combustion Burning fuel in oxygen Thermal Decomposition Thermal When a substance is broken down into 2 or Some compounds break down when heated, forming two or more Decomposition more products by heat products from one reactants. → Temp increase increase Temp Many metal carbonates can break down easily when it is heated: Reactivity series List of metals in order of reactivity Copper Carbonate → Copper Oxide + Carbon Dioxide Displacement A more reactive metal will displace a less Combustion Copper carbonate is green, copper oxide is black. We can test for reactive metal from its compound Combustion is another name for burning. It is an example of an carbon dioxide using limewater. Limewater is colourless, but turns exothermic reaction. There are two types of combustion – complete cloudy when carbon dioxide is bubbled through it.
    [Show full text]
  • Topic 5.1 Exothermic and Endothermic Reactions Heat And
    Topic 5.1 Exothermic and Endothermic Reactions Read pages 189- 194 of your IB Chemistry text book and the resource material provided in this packet. Answer questions 1-14. Heat and Temperature Often the concepts of heat and temperature are thought to be the same, but they are not. Temperature is a number that is related to the average kinetic energy of the molecules of a substance. A regular thermometer uses the expansion of a fluid to measure temperature. When the liquid (mercury or alcohol) in a thermometer is heated the average kinetic energy of the liquid particles increases, causing the particles to take up more space expanding them up the tube. The absolute temperature or Kelvin scale is an artificial temperature scale. The Celsius scale is based on the behavior of water molecules, with 0 oC being the freezing point or the point where the motion of the water molecules ceases. The Celsius scale has limited use when describing the motion of many substances, especially gases whose motions can cease at much lower temperatures. The mathematical conversion between oC and Kelvin is: °C + 273 = K K -273 = °C If Temperature is measured in Kelvin, then it is directly proportional to the average kinetic energy of the particles. In other words if you double the Kelvin temperature of a substance, you double the average kinetic energy of its molecules. KE α Absolute temp (K) Heat is a measurement of the total energy in a substance. That total energy is the sum of the kinetic (motion) and potential (stored) energies of the molecules.
    [Show full text]
  • Modeling Exothermic and Endothermic Reactions (1) Qualitative Models Endothermic Reaction: (Endo: “In”); Energy Is Absorbed
    Exothermic and Endothermic Reactions Modeling Exothermic and Endothermic Reactions (1) Qualitative Models Endothermic Reaction: (Endo: “In”); Energy is absorbed Note: Energy of the Product is greater than the Energy of the Reactant Energy Exothermic Reaction : (Exo: “Out”); Energy is released Note: Energy of the Reactant is greater than the Energy of the Product Energy You Do Qualitative Models: Your Turn Choose a design of your choice to illustrate the Endothermic and Exothermic Reactions Endothermic Reaction: (Endo: “In”); Energy is absorbed Note: Energy of the Product is greater than the Energy of the Reactant Exothermic Reaction : (Exo: “Out”); Energy is released Note: Energy of the Reactant is greater than the Energy of the Product (2) Quantitative Models – 1: Guided Practice Let each arrow be a bond and the numeric value on it indicates its Bond Energy in kJ/mol) Endothermic Reaction: Energy is Absorbed 12 Energy 18 18 12 12 20 12 12 12 Exothermic Reaction: Energy is Released Energy 12 12 You Do (2) Quantitative Models -1: Your Turn Choose a design of your choice to illustrate the Endothermic and Exothermic Reactions Enthalpy SC2g Develop a model to illustrate the release or absorption of energy (endothermic or exothermic) from a chemical reaction system depends upon the changes in total bond energy. SC5a Plan and carry out an investigation to calculate the amount of heat absorbed or released by chemical or physical processes. SC5b Construct an explanation using a heating curve as evidence of the effects of energy and intermolecular forces on phase changes. 1. Enthalpy is the scientific term for Heat of Reactions.
    [Show full text]
  • There Are Many Types of Chemical Reactions
    Chapter 11 There are many types of chemical reactions. In this chapter we will begin with combustion reactions and end with oxidation/reduction reactions. Along the way we examine various aspects of reactions. Homework from the book: Exercises: 1, 4, 6, 8, 9, 11,12, 14-19, 23, 25, 26, 29 Questions: none Problems: none In the study guide: Multiple 4, 5, 6, 12-18. There are many types of chemical reactions. In this chapter we will begin with combustion reactions and end with oxidation/reduction reactions. Along the way we examine various aspects of reactions. Combustion reactions: All combustion reactions fall into the pattern: fuel + O2 → CO2 + H2O The coefficients of the balanced equation will change depending on the fuel. The fuel can be almost anything including methane (CH4), propane (C3H8), butane (C4H10), octane (C8H18) or sugar (C6H12O6). The balanced equation for methane is CH4 + 2 O2 → CO2 + 2 H2O The balanced equation for octane is 2 C8H18 + 25 O2 → 16 CO2 + 18 H2O The combustion of methane or octane is exothermic; it releases energy. CH4 + 2 O2 → CO2 + 2 H2O + energy The energies of the products are lower than the energiies of the reactants. The excess energy is released as heat and light. Matter tends to move to lower energy states. A good analogy for a reaction is a ball falling down a hill. Exothermic reactions are more likely to occur. Endothermic reactions absorb energy. In these reactions the products are higher in energy than the reactants. Endothermic reactions are less likely to occur. You seldom see a ball spontaneously going uphill.
    [Show full text]