Chemistry. which statement about delta hf is true? a.) it is zero for any compound in its standard state. b.) it is positive when the bonds of the product store more energy than those of the reactants. c.) it is negative when a compound forms from elements in their standard states. d.) it is zero for any element that is in the liquid state.

The answer is B on edge

Explanation:

Here are my notes on this section for anyone that needs them

Enthalpy and State Function

Bonds contain potential energy. Breaking and forming bonds involves energy. Reactants and products contain energy. Enthalpy (H) is a measure of heat and internal energy in a system.

A state function is a quantity whose change in magnitude during a process depends only on the beginning and end points the process, not the path taken between them. Enthalpy change during reaction depends only on the identity of reactants and products and their initial and finial condition

Enthalpy of Formation

enthalpy of formation (Hf) is the energy absorbed or released when a pure substance forms from elements in their standard states

Units: kJ/mol, kcal/mol

Standard state is the natural state of an element at 1 atm (atmosphere of pressure) and 25 degrees celsius. Hf for a pure element in its standard state is 0 kJ/mol.

H (hydrogen):    H2(g)

N (nitrogen):N2(g)

O (oxygen):O2(g)

F (fluorine):F2(g)

Cl (chlorine):Cl2(g)

Br (bromine):Br2(l)

Hg (mercury):Hg(l)

Enthalpy of Reaction

Enthalpy of reaction (Hrxn) is energy absorbed or released during a chemical reaction

Hrxn negative: exothermic reaction

Hrxn positive: endothermic reaction

Hess's Law: Hrxn =  Σ(ΔHƒ, products) − Σ(ΔHƒ, reactants)

thermochemical equation: the chemical equation that shows the state of each substance involved and the energy change involved in a reaction

Find the kJ/mol of the product and then subtract the kJ/mol of the reactants.

B: It is positive when the bonds of the product store more energy than those of the reactants.

Explanation:

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answer: i dont know the answer but this might

explanation:

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