Define internal energy (E) .

The sum of all kinetic energy and potential energy of the components of a system. The change in internal energy, ΔE, is calculated as E final − E initial and indicates whether energy has been gained (positive ΔE) or lost (negative ΔE) by the system.

True or False? According to the first law of thermodynamics, any energy lost by the system must be gained by the surroundings.

True. The first law of thermodynamics states that energy can neither be created nor destroyed. Therefore, energy is conserved across the system and surroundings: any decrease in the system's energy is exactly matched by an increase in the surroundings' energy, and vice versa.

In the equation ΔE = q + w, why does work done by the system carry a negative sign?

When the system does work on the surroundings (e.g. expanding against external pressure ), it expends energy — energy leaves the system, so w is negative by convention. A negative w therefore reduces ΔE.

Define specific heat capacity (C_s) .

The quantity of heat required to raise the temperature of 1 g of a substance by 1°C (or 1 K). It is measured in J g -1 °C -1 and used in the equation q = m × C_s × Δ*T *.

True or False? The direction of heat flow between two objects in contact is determined by the difference in their specific heat capacities.

False. Heat flows from the object at higher temperature to the object at lower temperature, regardless of specific heat capacity or mass. Specific heat capacity affects how much the temperature of each object changes as heat is transferred, but not the direction of flow.

How do heating/cooling curves demonstrate that energy changes in a chemical system can involve both temperature change and phase transition?

On a heating curve, sloping segments show temperature rising as heat increases average kinetic energy , while flat segments (at melting and boiling points) show temperature held constant as heat overcomes intermolecular forces — increasing potential energy without changing kinetic energy. This demonstrates that energy input drives either temperature change or phase change, not both simultaneously.

The change in internal energy of a system is given by the equation ΔE = .......... + w, where q is the heat exchanged and w is the work done.

The change in internal energy of a system is given by the equation ΔE = q + w, where q is the heat exchanged and w is the work done.

An experimental technique used to measure enthalpy changes in chemical reactions by monitoring temperature changes in a known mass of liquid (typically water) within a calorimeter .

True or False? In a calorimetry experiment, the sign of ΔH for a combustion reaction is negative because heat is released by the reaction into the water.

True. Combustion is an exothermic process: the system (reaction) loses heat to the surroundings (water in the calorimeter), so q for the reaction is negative. Δ H = q ÷ n is therefore also negative, indicating heat is released.

Why do calorimetry experiments using a simple flame calorimeter typically give enthalpy of combustion values that are less exothermic than the accepted literature value?

Two main sources of error reduce the apparent heat released: (1) heat losses to the surroundings mean not all the energy from combustion reaches the water; and (2) incomplete combustion means the fuel is not fully oxidized, so less energy is released per mole than the theoretical maximum.

True or False? In calorimetry calculations using q = mcΔ T , the temperature change must be converted from °C to K before being substituted into the equation.

False. The temperature change ΔT is the same whether calculated in °C or K, because the interval size is identical on both scales. There is no need to convert units before calculating ΔT.

In a calorimetry calculation, the heat transferred is calculated using q = m × .......... × ΔT, where m is the mass of the liquid and ΔT is the temperature change.

In a calorimetry calculation, the heat transferred is calculated using q = m × C_s × ΔT, where m is the mass of the liquid and ΔT is the temperature change.

AP®ChemistryCollege BoardFlashcardsUnit 6: ThermochemistryCalorimetry

Calorimetry (College Board AP® Chemistry): Flashcards

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Define internal energy (E).

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Define internal energy (E).
The sum of all kinetic energy and potential energy of the components of a system. The change in internal energy, ΔE, is calculated as Efinal − Einitial and indicates whether energy has been gained (positive ΔE) or lost (negative ΔE) by the system.
True or False?
According to the first law of thermodynamics, any energy lost by the system must be gained by the surroundings.
True.
The first law of thermodynamics states that energy can neither be created nor destroyed. Therefore, energy is conserved across the system and surroundings: any decrease in the system's energy is exactly matched by an increase in the surroundings' energy, and vice versa.
In the equation ΔE = q + w, why does work done by the system carry a negative sign?
When the system does work on the surroundings (e.g. expanding against external pressure), it expends energy — energy leaves the system, so w is negative by convention. A negative w therefore reduces ΔE.
Define specific heat capacity (C_s).
The quantity of heat required to raise the temperature of 1 g of a substance by 1°C (or 1 K). It is measured in J g^-1 °C^-1 and used in the equation q = m × C_s × Δ*T*.
True or False?
The direction of heat flow between two objects in contact is determined by the difference in their specific heat capacities.
False.
Heat flows from the object at higher temperature to the object at lower temperature, regardless of specific heat capacity or mass. Specific heat capacity affects how much the temperature of each object changes as heat is transferred, but not the direction of flow.
How do heating/cooling curves demonstrate that energy changes in a chemical system can involve both temperature change and phase transition?
On a heating curve, sloping segments show temperature rising as heat increases average kinetic energy, while flat segments (at melting and boiling points) show temperature held constant as heat overcomes intermolecular forces — increasing potential energy without changing kinetic energy. This demonstrates that energy input drives either temperature change or phase change, not both simultaneously.
The change in internal energy of a system is given by the equation ΔE = .......... + w, where q is the heat exchanged and w is the work done.
The change in internal energy of a system is given by the equation ΔE = q + w, where q is the heat exchanged and w is the work done.
Define calorimetry.
An experimental technique used to measure enthalpy changes in chemical reactions by monitoring temperature changes in a known mass of liquid (typically water) within a calorimeter.
True or False?
In a calorimetry experiment, the sign of ΔH for a combustion reaction is negative because heat is released by the reaction into the water.
True.
Combustion is an exothermic process: the system (reaction) loses heat to the surroundings (water in the calorimeter), so q for the reaction is negative. ΔH = q ÷ n is therefore also negative, indicating heat is released.
Why do calorimetry experiments using a simple flame calorimeter typically give enthalpy of combustion values that are less exothermic than the accepted literature value?
Two main sources of error reduce the apparent heat released: (1) heat losses to the surroundings mean not all the energy from combustion reaches the water; and (2) incomplete combustion means the fuel is not fully oxidized, so less energy is released per mole than the theoretical maximum.
True or False?
In calorimetry calculations using q = mcΔT, the temperature change must be converted from °C to K before being substituted into the equation.
False.
The temperature change ΔT is the same whether calculated in °C or K, because the interval size is identical on both scales. There is no need to convert units before calculating ΔT.
In a calorimetry calculation, the heat transferred is calculated using q = m × .......... × ΔT, where m is the mass of the liquid and ΔT is the temperature change.
In a calorimetry calculation, the heat transferred is calculated using q = m × C_s × ΔT, where m is the mass of the liquid and ΔT is the temperature change.

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