When there is temp rise in a chemical reaction, is this exothermic or endothermic? Is the sign (+) or (-)?
812725
When heat is given off, the reaction is exothermic. The sign is negative because energy is given off by the system.
980614
Most reactions are exothermic.\nExamples are\n\n1) Combustion of methane (natural gas).\n\n2) Neutralisation of dilute hydrochloric acid and dilute sodium hydroxide.\n\n3) Adding concentrated sulphuric acid to water is highly exothermic\n(gives out a lot of heat).\n\n4) Adding water to anhydrous copper(II) sulphate is exothermic.\nAnhydrous copper(II) sulphate is white.\nAnhydrous is pronounced "an-high-drus"\nAnhydrous means "without water".\n\nAnhydrous copper(II) sulphate is copper(II) sulphate which is completely dry.\nWhen water is added to it, it turns into the familiar hydrated blue crystals.\nHydrated is pronounced "high-dray-tid", it means 'with water'.\nThis is used as a test for water.\n\nIf blue (hydrated) copper(II) sulphate crystals are heated,\nan endothermic reaction occurs, they lose their water,\nturn white and become anhydrous copper(II) sulphate crystals again.\nThe reaction is reversible.\nBlue copper(II) sulphate crystals will also turn white in the presence of  concentrated sulphuric acid.\n\n \n\nFew reactions are endothermic.\nExamples are\n\n1) Photosynthesis.\nChlorophyll is a very clever catalyst,\nwhich allows plants to make sugar from carbon dioxide in the air.\nThe energy needed for the reaction comes from sunlight.\n\n2) The reaction of ethanoic acid with sodium carbonate.\n\n3) Dissolving some salts in water is an endothermic process.\nPotassium chloride and ammonium nitrate\nboth take in energy when they dissolve.\nIf you put a thermometer in the solution as they are dissolving,\nyou will see that the temperature drops.\n\n4) Melting, boiling and evaporation\nare all endothermic processes (not reactions).
1172999
1) exothermic reaction gives off heat\n2) a catalyst speeds up a rxn\n3) chemical reaction\n4) 1  + 4 --> 1 + 1 + 2\n\ncoefficiant of HCl is one
1117740
since the heat evolced is dissipated in 100 ml in the first case (by 2 g of NaOH) and in 200 ml of water in the second case (by 4 g of NaOH) the net change in the temperature of water remains same. hence 'a' is the correct answer\n answer : 'a'
676662
What did you use to measure the activation energy?\nYou probably used a thermometer which was probably marked to 0.1 degrees. That means that your measurements are probably accurate to 0.05 degrees. Reactant temp may be measured as 25.1 (+/- 0.05) so this means that the temp could be 25.05 or 25.15.\nDo the same for the transition state.\nDraw small vertical bars through each point, e.g for reactants the bar would go from 25.05 to 25.15, with the cross in the middle.\n\nI can't remember the formula, but errors are a lot more pronounced on small measurements. It may be something like the difference between the highest and lowest possible values divided by the recorded value, all multiplied by 100.
758038
omg i just had this test today too!!!!!\nu use hess's law to find the change--->write the equation, balance it \nand then use the moles (coefficients) to multiply the change is s for each element.\nhess's law says that the change in s (or h, or G, its usable for any of the thermo concepts) is equal to the sum of the entropy of the products minus the sum of the entropy of the reactants. its pretty easy after that--just plug in the values and there you go! \ngood luck! this was a hard test for me..so confusing.
621304
If you are dealing with a reaction involving ammonium carbonate in solution, and are then measuring the temperature rise, you need the heat capacity of water, not ammonium carbonate!\n\nPlease post the whole problem if that is not enough to get you started.
799142
Hi Kitty,\nHere is what, I think, you need.\n\nExothermic, dH < 0, Entropy increases (dS > 0), Reaction is spontaneous and dG < 0\n\nExothermic, dH < 0, Entropy decreases (dS < 0), Reaction is spontaneous only at low temps, if |T dS| < |dH|\n\nEndothermic, dH > 0, Entropy increases (dS > 0), Reaction is spontaneous only  at high temps, if T dS > dH\n\nEndothermic, dH > 0, Entropy decreases (dS < 0), Reaction is not spontaneous, dG > 0
462388
A reaction which gives out heat is known as 'exothermic'.\n\nhttp://en.wikipedia.org/wiki/Exothermic\nhttp://en.wikipedia.org/wiki/Exothermic_reaction\n\nA reaction which seems to absorb heat is known as 'Endothermic'\n\nhttp://en.wikipedia.org/wiki/Endothermic\nhttp://en.wikipedia.org/wiki/Endothermic_reaction
766931
exothermic=right\nendothermic=left
746528
The correct answer is C, standard potential should be positive.\nIf you can't remember the correct sign easily, think that the general rule is that for spontaneous reactions of any type ΔG<0 and since for redox reactions ΔG= -nFE, E has to be positive.
20960
Many chemical reactions release energy in the form of heat, light, or sound. These are exothermic reactions. Exothermic reactions may occur spontaneously and result in higher randomness or entropy (ÄS > 0) of the system. They are denoted by a negative heat flow (heat is lost to the surroundings) and decrease in enthalpy (ÄH < 0). In the lab, exothermic reactions produce heat or may even be explosive. \nThere are other chemical reactions that must absorb energy in order to proceed. These are endothermic reactions. Endothermic reactions cannot occur spontaneously. Work must be done in order to get these reactions to occur. When endothermic reactions absorb energy, a temperature drop is measured during the reaction. Endothermic reactions are characterized by positive heat flow (into the reaction) and an increase in enthalpy (+ÄH). \n\nExamples of Endothermic and Exothermic Processes \n\nPhotosynthesis is an example of an endothermic chemical reaction. \nIn this process, plants use the energy from the sun to convert carbon dioxide and water into glucose and oxygen. This reaction requires 15MJ of energy (sunlight) for every kilogram of glucose that is produced: \nsunlight + 6CO2(g) + H2O(l) = C6H12O6(aq) + 6O2(g) \n\nAn example of an exothermic reaction is the mixture of sodium and chlorine to yield table salt. This reaction produces 411 kJ of energy for each mole of salt that is produced:\n\nNa(s) + 0.5Cl2(s) = NaCl(s) \n\nEndothermic reactions take in energy from their surroundings\n\nDissolution of salt is endothermic because it takes energy from its surroundings not much but some.\n\nAs far as all salt types that is hard to say and you would have to do an experiment to find that out.
1138895
They are similar, but not exactly the same as exothermic and endothermic as used in chemistry.\n\nExergonic reactions are defined by a negative delta G, that is they have less free energy at the completion of the reaction than at the beginning. An exergonic reaction may require some input of energy (called activation energy) to start the reaction, but will release more energy than is required to start the reaction. \n\nAn endergonic reaction will have a positive delta G, that is it will have more free energy after the reaction than before it starts. \n\nHere is a good website with graphs that explains this:\nhttp://academic.brooklyn.cuny.edu/biology/bio4fv/page/exergon.htm\n\nHere is another website that explains the similarity between exergonic, exothermic, and endergonic, endothermic. \nhttp://www.mansfield.ohio-state.edu/~sabedon/sta5202.htm
613069
1.  There is no process going on in beaker A.\n2.  The process in beaker B is endothermic, as it consumes heat from the surroundings.\n3.  He's studying the heat of solution of sodium nitrate.\n4.  It's negative, as it's an exothermic process (the heat of the system is lower than it was to start, as heat was given off to the surroundings).
816266
At constant temperature and pressure, a reaction will occur spontaneously if the Gibbs free energy change of the reaction is less than zero.\n\nThe Gibbs free energy change, delta-G, is given by:\n\ndelta-G = delta-H - T*delta-S\n\nwhere T is the absolute temperature, \n\ndelta-H is the enthalpy change for the reaction at the temperature and pressure of interest, and \n\ndelta-S is the entropy change for the reaction at the temperature and pressure of interest.\n\nWe are given the enthalpy and entropy changes for the reaction of interest in this question.  I will assume that the enthalpy change and entropy change for this reaction are independent of temperature. Note also that the values of both of these thermodynamic parameters are given in "per mole" units in this question.  I will assume these values are normalized to one mole of one of the reactants, and not one mole of the product.\n\nPlugging in the enthalpy and entropy values, and setting delta-G to be less than zero:\n\n0 > delta-G = 180500 J/mol - T*24.8 J/(mol*K)\n\nT  > (180500 J/mol)/(24.8 J/(mol*K))\n\nT > 7278 K\n\nThis is an extremely high temperature, and the assumption that we made above that the enthalpy and entropy changes were independent of temperature is probably not very good given the large temperature extrapolation being made here.
596796
Exothermic reactions produce heat, so the temperature of the reactants will become higher, i.e. hotter.\n\nEndothermic reactions absorb heat from the surroundings, so the temperature of the reactants will become lower, i.e. colder.\n\nFor a specific reaction involving Co(II), would have to know the reaction equation and then calculate the delta enthalpy.
919726
Sodium and Potassium are correct.  They are extremely reactive with water, forming Hydroxides and releasing Hydrogen gas, which is usually ignited from the exothermic heat of the reaction.\n\n2 Na + 2 H2O -> 2 NaOH + H2\n\n2 K + 2H2O -> 2 KOH + H2