Why did you use two coffee cups during the experiment instead of just one cup? in case a hole was punctured in the inside cup to add more weight in order to stabilize the calorimeter to hold all of the water in the experiment so that the water temperature would not change at all during the experiment to prevent heat …

The heat capacity and the specific heat are related by C=cm or c=C/m. The mass m, specific heat c, change in temperature ΔT, and heat added (or subtracted) Q are related by the equation: Q=mcΔT. Values of specific heat are dependent on the properties and phase of a given substance.

To calculate the amount of heat released in a chemical reaction, use the equation Q = mc ΔT, where Q is the heat energy transferred (in joules), m is the mass of the liquid being heated (in kilograms), c is the specific heat capacity of the liquid (joule per kilogram degrees Celsius), and ΔT is the change in …

This assumes no heat is lost to the surroundings from the calorimeter. Heat lost by hot water = heat gained by cold water + heat gained by calorimeter. The minus sign indicates that the hot water is losing heat, whereas the cold water and the calorimeter are gaining heat.

The specific heat capacity of water is 4.18 J/g/°C. We wish to determine the value of Q – the quantity of heat. To do so, we would use the equation Q = m•C•ΔT. The m and the C are known; the ΔT can be determined from the initial and final temperature.

Explanation: The “calorimeter constant” is just the specific heat of the calorimeter and its thermal conductivity. An “ideal” calorimeter would have a very low specific heat and zero thermal conductivity because the point is to conserve energy within the system.

Uses. The calorimeter constants are used in constant pressure calorimetry to calculate the amount of heat required to achieve a certain raise in the temperature of the calorimeter’s contents.

Heat Capacity – amount of energy necessary to raise the temperature of a gram of substance one degree celsius. Calorimeter Constant – amount of energy needed to raise the temperature of the calorimeter one degree celsius.

If the calorimeter had a low specific heat, it would absorb less heat, but its temperature would increase more. Therefore, a calorimeter with a high specific heat would be more effective because it would minimize heat transfer between the calorimeter and the surroundings.

The calorimeter constant can never be negative — if it is, you have made a mistake… Try performing multiple trials and averaging out the results of those trials to reduce your error. The uncertainty in your final average will be plus/minus 2x the standard deviation.

More reliable results can be obtained by repeating the experiment many times. The biggest source of error in calorimetry is usually unwanted heat loss to the surroundings. This can be reduced by insulating the sides of the calorimeter and adding a lid.

law of conservation of energy

The principle of calorimetry indicates the law of conservation energy, i.e. the total heat lost by the hot body is equal to the total heat gained by the cold body.

Hess’s Law describes the conservation of energy in chemical reactions, stating that the heat flow of a reaction is equal to the sum of the heat flow of its composite reactions. Calculating the Qcal, the heat of the calorimeter, allows you to adjust your readings to determine the total heat flow of a reaction.

Enthalpy changes Enthalpy change is the name given to the amount of heat evolved or absorbed in a reaction carried out at constant pressure. It is given the symbol ΔH, read as “delta H”.

Enthalpy (heat) of solution can be determined in the laboratory by measuring the temperature change of the solvent when solute is added. ΔH is negative if energy (heat) is released (exothermic). ΔH is positive if energy (heat) is absorbed (endothermic).

For example, when water changes from liquid to gas, delta H is positive; the water gains heat. When water changes from liquid to solid, delta H is negative; the water loses heat.

Q is the energy transfer due to thermal reactions such as heating water, cooking, etc. anywhere where there is a heat transfer. You can say that Q (Heat) is energy in transit. Enthalpy (Delta H), on the other hand, is the state of the system, the total heat content.

When heat is absorbed by the solution, q for the solution has a positive value. This means that the reaction produces heat for the solution to absorb and q for the reaction is negative. When heat is absorbed from the solution q for the solution has a negative value.

ΔH is the change in enthalpy, which is heat content. ΔE is the change in energy of a system.

If H (enthalpy change) is zero, it means that the spontaneity of the reaction only depends on the entropy change (S). If entropy change for this reaction is +ve, the reaction will always be spontaneous at all temperatures. If entropy change is -ve, the reaction will always be non-spontaneous at all temperatures.

When delta H is negative, it means the products in the reaction have lower energy compared to the reactants, so the reaction has lost energy and released it as heat, making it exothermic. You can think of it as positive delta H as gaining energy and negative delta H as releasing energy.

A negative delta S corresponds to a spontaneous process when the magnitude of T * delta S is less than delta H (which must be negative). delta G = delta H – (T * delta S). A negative delta S would mean that the products have a lower entropy than the reactants, which is not spontaneous by itself.

In adiabatic process, there is no exchange of heat between system and surroundings (q=0) The system is completely insulated from surroundings. For exothermic process, the temperature of system rises and for endothermic process, the temperature of system falls. Hence, ΔH=0.

According to the definition of an adiabatic process, ΔU=wad. Therefore, ΔU = -96.7 J. Calculate the final temperature, the work done, and the change in internal energy when 0.0400 moles of CO at 25.0oC undergoes a reversible adiabatic expansion from 200. L to 800.

The CHANGE in enthalpy is zero for isothermal processes consisting of ONLY ideal gases. For ideal gases, enthalpy is a function of only temperature. Thus, in any isothermal process involving only ideal gases, the change in enthalpy is zero.

An adiabatic process is a process in which no heat is exchanged. An adiabatic and reversible process has constant entropy s–it is isentropic. An isenthalpic process has constant enthalpy, and probably there is a myriad ways to realize such a process.

Enthalpy is the amount of internal energy contained in a compound whereas entropy is the amount of intrinsic disorder within the compound. Enthalpy is zero for elemental compounds such hydrogen gas and oxygen gas; therefore, enthalpy is nonzero for water (regardless of phase).

The assumption that a process is adiabatic is a frequently made simplifying assumption. For such an adiabatic process, the modulus of elasticity (Young’s modulus) can be expressed as E = γP, where γ is the ratio of specific heats at constant pressure and at constant volume (γ = CpCv ) and P is the pressure of the gas .

Isoenthalpic means constant enthalpy, and any material which passes through a system without a change of enthalpy has, by definition, passed through an isoenthalpic process. For instance, it is common to assume that a gas or vapor flashing through a valve is an isoenthalpic process.