Energy: Applications & Uses Lowersixth Science Chemistry

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A quick true/false on enthalpy and energy changes.

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YAKILI LMS

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Energy: Applications & Uses Lowersixth Science Chemistry


Energy: Applications & Uses Lowersixth Science ChemistryOnline version A quick true/false on enthalpy and energy changes. by YAKILI LMS 1 Enthalpy changes can be used to compare efficiency of different reactions. Yes No 2 Absolute enthalpy values can be directly measured with a simple thermometer. Yes No 3 A system’s energy change is identical in sign to the surroundings' energy change. Yes No 4 Calorimetry experiments can be used to determine ΔH for a reaction. Yes No 5 The standard enthalpy of formation is always positive. Yes No 6 If ΔH is positive for a reaction, the system gains energy from the surroundings. Yes No 7 Combustion enthalpy is not used in calorimetry studies. Yes No 8 Enthalpy is a measure of the volume of a system at constant temperature. Yes No 9 In an exothermic reaction, ΔH is negative at constant pressure. Yes No 10 Standard enthalpy of formation is defined for products minus reactants to form a compound from elements in their standard states. Yes No 11 The first law of thermodynamics states that energy cannot be created or destroyed. Yes No 12 Standard states are defined only for gases at 290 K. Yes No 13 The standard enthalpy of reaction is the enthalpy change at any temperature. Yes No 14 Calorimeters are designed to minimize heat exchange with the surroundings. Yes No 15 Enthalpy change is denoted by ΔH. Yes No 16 Standard conditions for enthalpy change measurement are 0 K and 0 atm. Yes No 17 An isolated system cannot exchange energy with its surroundings. Yes No 18 Heat transfer during a reaction occurs until thermal equilibrium is reached. Yes No 19 The enthalpy change of fusion is an example of a phase change enthalpy. Yes No 20 Endothermic and exothermic refer to color changes in the reaction. Yes No 21 Reaction enthalpies are intrinsic properties independent of the path taken. Yes No 22 Reaction enthalpies are always dependent on the reaction pathway. Yes No 23 Calorimeters are designed to maximize heat exchange with the surroundings. Yes No 24 Endothermic reactions absorb heat from the surroundings. Yes No 25 The energy content of a system can be tracked by changes in internal energy and enthalpy. Yes No 26 Fusion enthalpy is not involved in phase changes. Yes No 27 Calorimetry cannot be used to determine enthalpy changes. Yes No 28 The term 'enthalpy' refers to entropy multiplied by pressure. Yes No 29 Standard conditions for enthalpy change measurements are commonly 298 K and 1 atm. Yes No 30 Hess's law states that enthalpy changes depend on the path of the reaction. Yes No 31 The enthalpy change of combustion is a common example studied in calorimetry. Yes No 32 Absolute zero is the only reference point for enthalpy calculations. Yes No 33 The standard enthalpy of formation uses elements in their standard states as zero. Yes No 34 The concept of enthalpy is widely used in chemical thermodynamics and physical chemistry. Yes No 35 The term 'enthalpy' combines internal energy and pressure-volume work. Yes No 36 The first law of thermodynamics allows creation of energy in a chemical reaction. Yes No 37 Exothermic reactions release heat to the surroundings. Yes No 38 Endothermic and exothermic describe energy transfer direction relative to the surroundings. Yes No 39 Enthalpy changes cannot be used to assess reaction feasibility. Yes No 40 If ΔHrxn is negative, the reaction is endothermic. Yes No 41 The first law implies that energy is created during chemical processes. Yes No 42 Measurement of enthalpy changes often uses calorimetry. Yes No 43 ΔH represents entropy change in a system. Yes No 44 The first law implies conservation of energy in chemical processes. Yes No 45 Endothermic reactions release heat to the surroundings. Yes No 46 Heat transfer stops before reaching thermal equilibrium in a reaction. Yes No 47 A system's energy change can be calculated from the surroundings' energy change with opposite sign. Yes No 48 Hess's law allows enthalpy changes to be added for overall reactions. Yes No 49 An open system cannot exchange energy with its surroundings. Yes No 50 The standard enthalpy of reaction is the enthalpy change when reactants in their standard states form products in their standard states. Yes No

Energy: Applications & Uses Lowersixth Science Chemistry

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Energy: Applications & Uses Lowersixth Science ChemistryOnline version

by YAKILI LMS

Enthalpy changes can be used to compare efficiency of different reactions.

Absolute enthalpy values can be directly measured with a simple thermometer.

A system’s energy change is identical in sign to the surroundings' energy change.

Calorimetry experiments can be used to determine ΔH for a reaction.

The standard enthalpy of formation is always positive.

If ΔH is positive for a reaction, the system gains energy from the surroundings.

Combustion enthalpy is not used in calorimetry studies.

Enthalpy is a measure of the volume of a system at constant temperature.

In an exothermic reaction, ΔH is negative at constant pressure.

Standard enthalpy of formation is defined for products minus reactants to form a compound from elements in their standard states.

The first law of thermodynamics states that energy cannot be created or destroyed.

Standard states are defined only for gases at 290 K.

The standard enthalpy of reaction is the enthalpy change at any temperature.

Calorimeters are designed to minimize heat exchange with the surroundings.

Enthalpy change is denoted by ΔH.

Standard conditions for enthalpy change measurement are 0 K and 0 atm.

An isolated system cannot exchange energy with its surroundings.

Heat transfer during a reaction occurs until thermal equilibrium is reached.

The enthalpy change of fusion is an example of a phase change enthalpy.

Endothermic and exothermic refer to color changes in the reaction.

Reaction enthalpies are intrinsic properties independent of the path taken.

Reaction enthalpies are always dependent on the reaction pathway.

Calorimeters are designed to maximize heat exchange with the surroundings.

Endothermic reactions absorb heat from the surroundings.

The energy content of a system can be tracked by changes in internal energy and enthalpy.

Fusion enthalpy is not involved in phase changes.

Calorimetry cannot be used to determine enthalpy changes.

The term 'enthalpy' refers to entropy multiplied by pressure.

Standard conditions for enthalpy change measurements are commonly 298 K and 1 atm.

Hess's law states that enthalpy changes depend on the path of the reaction.

The enthalpy change of combustion is a common example studied in calorimetry.

Absolute zero is the only reference point for enthalpy calculations.

The standard enthalpy of formation uses elements in their standard states as zero.

The concept of enthalpy is widely used in chemical thermodynamics and physical chemistry.

The term 'enthalpy' combines internal energy and pressure-volume work.

The first law of thermodynamics allows creation of energy in a chemical reaction.

Exothermic reactions release heat to the surroundings.

Endothermic and exothermic describe energy transfer direction relative to the surroundings.

Enthalpy changes cannot be used to assess reaction feasibility.

If ΔHrxn is negative, the reaction is endothermic.

The first law implies that energy is created during chemical processes.

Measurement of enthalpy changes often uses calorimetry.

ΔH represents entropy change in a system.

The first law implies conservation of energy in chemical processes.

Endothermic reactions release heat to the surroundings.

Heat transfer stops before reaching thermal equilibrium in a reaction.

A system's energy change can be calculated from the surroundings' energy change with opposite sign.

Hess's law allows enthalpy changes to be added for overall reactions.

An open system cannot exchange energy with its surroundings.

The standard enthalpy of reaction is the enthalpy change when reactants in their standard states form products in their standard states.