A thermodynamic system undergoes a process in which it's internal energy decreases by 500J?
817267
I think such questions bug you because you equate Internal energy with heat. Work and heat both are energies in transit for any process. Think like this, if 220 J work is done on the system its energy would have increased by 220 J.  But instead its internal energy has decreased by 500 J. So only way this is possible is if [220 + 500] J = 720 J amount of energy leaves the system as heat.  If you like to work with formula then\n\nU = Q - W\n\nhere U = -500; Q unknown: W = -220\n\n-500 = Q - (-220)\n-500 = Q + 220\n\nor Q = -500-220 =  -720 J\n\nSo the system has lost 720 J as heat
920392
it is exothermic, meaning the combustion reaction releases energy in the form of heat
604334
since vapourisation is a change in state, it can also be called the heat of vapourisation, q.\n\nΔS = q/T\n\nΔS = (30x1000)/(273.15+76.8)\n\nΔS = 85.73 J/mol/K\n\nThat answer is only correct to a certain degree, since the temprature changes during the process, but it does.  otherwise you have to do a lot of maths.
846365
Thermal energy IS heat.
764428
For chemical changes, the energy scale is related to Boltzmann's constant, which is pretty small, and temperature, which is typically 1-1000 degrees.  For nuclear changes the energy scale is related to Einstein's famous equation, E=mC^2.  the mass is typically very small, say the mass of an electron, but the speed of light is huge, and gets squared, making nuclear changes much, much more energetic than chemical changes.  That's why the sun is billions of degrees, and a campfire is only thousands of degrees.
390030
heat mainly - the energy is wasted in the form of heat instead of transferable energy
406343
it's heat absorbed or radiated during a change of a phase at constant temperature/pressure.
741934
This question is not very difficult, but you have to be careful with the unit.\n\nFirst, to solve this question, you have to make the assumption that there is no energy lost to the surrounding. So, this reduces to a two system question: one system releases energy while the other one absorbs energy. In this particular case, the energy released by one system equals to the energy absorbs by another system.\n\nTo picture this, consider the following, one system goes through the following steps:\n\n[100 degC H2O (g)] --- (step 1) ---> [100 degC H2O (l)] --- (step 2) ----> [79 deg C H2O (l)]\n\nThe other system goes through:\n\n[260g, 19 degC H2O (l)] --- (step 3) ----> [260g, 79 degC H2O (l)]\n\nIt is obvious that both the energy involved in step 1 and 2 is negative (release energy) while the energy involved in step 3 is positive (absorb energy). We will now solve this question backwards. Starts from finding out the energy required by step 3, then find out the mass of steam that is needed to do the job.\n\nEnergy absorbed during step 3:\nE3 = m * c * (delta T) = 260 * 4.18 * (79-19) = 65208 J\n\nEnergy released in step 1:\n(you have to first convert heat of vaporization to unit of J/g: 40.7*1000/18 = 2261.111 J/g)\nE1 = m' * l = m' * 2261.111 = 2261.111m'\n(since we don't know how much steam is needed, let's just leave it as a variable)\n\nEnergy released in step 2:\n(the temperature of the same mass as above will be lowered to 79 degC)\nE2 = m' * 4.18 * (100-79) = 87.78m'\n\nTotal energy released = (2261.111 + 87.78)m' = 2348.89m'\n\nEnergy released = energy absorbed\n2348.89m' = 65208\nm' = 27.76 g
435509
An Endo-Thermic reaction takes in heat and the Exo-thermic one gives off heat.\n\nExemples:\n\nExothermic processes         \nmaking ice cubes              \nrusting iron                  \n\nEndothermic processes \nmelting ice cubes \ncooking an egg
866663
B\n\nStrong bonds requires more energy to break and therefore releases more energy when formed.\n\nSo when enthalpy(heat) is lowered you need something that eats more energy than it creates.
751958
heat change\ncolor change\nemmission of gas
486588
In chemistry, an exothermic reaction is one that releases heat. It is the opposite of an endothermic reaction. Expressed in a chemical equation:\n   reactants -->product + energy\n\nWhen using a calorimeter, the change in heat of the calorimeter is equal to the opposite of the change in heat of the system. This means that when the solution in which the reaction is taking place gains heat, the reaction is exothermic.\n\nIn an exothermic reaction the total energy absorbed in bond breaking is less than the total energy released in bond making.\n\nThe absolute amount of energy in a chemical system is extremely difficult to measure or calculate. The enthalpy change, ΔH, of a chemical reaction is much easier to measure and calculate. A bomb calorimeter is very suitable for measuring the energy change, ΔH, of a combustion reaction. Measured and calculated ΔH values are related to bond energies by:\n\nΔH = energy used in bond breaking reactions - energy released in bond making products.\n\nFor an exothermic reaction, this gives a negative value for ΔH as a larger value is subtracted from a smaller value. For example, when hydrogen burns:\n\n\nIn Chemistry an endothermic reaction is one in which the products have more energy than the reactants, and thus a net input of energy, usually in the form of heat, is required. Endothermic reactions are often described as reactions that "feel cold", and contrast with exothermic reactions, in which heat is released.\n\nAlthough the process of bond breaking amongst reactants in a chemical process requires an initial input of energy (the activation energy), in the case of an endothermic reaction, the energy released when bonds are formed to create reactants is less than that required to break the bonds in the products; bonding electrons in the products are therefore at a higher energy than the reactants. Heat energy from the material surrounding the reactants is usually what breaks their bonds, so as heat energy is transferred from the surroundings to the reactants, the surroundings get colder. This is often summarized in a chemical equation as follows:\n\nReactants + Energy → Products
174163
thermodynamics?  heat transfer?  physics?
930423
All energy eventually ends up as heat.  It's a consequence of the second law of thermodynamics
705010
The specific heat of water is 4.184 J/g C.\n\nSince water's density is 1 g/ mL, 500 mL of water will have a mass of 500 g.\n\nThe equation to calculate the joules of energy required is:\nheat = mass x specific heat x change in temperature\n\nenergy = 500 (4.184)(95-25)\nenergy = 146440 J or 146.44 kJ\n\nHope that helped.
801823
Quantity of energy lost by the water:\n\nQ = (m)(deltaT)(c)\n    = (50.0g)(2.00 C)(4.18 JK-1g-1)\n    = 418 J\n\nSince the water got colder, the reaction was endothermic, and the enthalpy change for the reaction is +418 J.
967087
a.) It decreases.\n\nThis is because only the most energetic particles evaporate.  They need to achieve the evaporation temperature, and then also absorb the energy of vaporization.  This can happen to individual particles even when the liquid as a whole is not near this temperature.  Because only the most energetic particles evaporate, the remaining particles will have a lower internal energy, and thus a lower temperature.