For an exothermic reaction, increasing the temperature shifts the equilibrium to the left, favoring the reactants, as the system tries to absorb the added heat.

For an exothermic reaction, increasing the temperature shifts the equilibrium to the left, favoring the reactants, as the system tries to absorb the added heat.

For an exothermic reaction, increasing the temperature shifts the equilibrium to the left, favoring the reactants, as the system tries to absorb the added heat.

For an exothermic reaction, increasing the temperature shifts the equilibrium to the left, favoring the reactants, as the system tries to absorb the added heat.

According to Le Chatelier's Principle, increasing the temperature of an exothermic reaction shifts the equilibrium to the left, favoring the reactants.

In a back titration, the first step is to add an excess of titrant to the analyte, followed by titrating the unreacted titrant.

Using q = mcΔT, ΔT = q/(mc) = 4200 J / (100 g * 4.18 J/g°C) = 10 °C.

q = mcΔT = 50 g × 4.18 J/g°C × (75-25)°C = 1045 J.

A negative q value indicates that heat is released by the system, signifying an exothermic process.

A negative ΔH indicates that the reaction releases heat to the surroundings, classifying it as exothermic.

A positive q value indicates that the system absorbs heat from the surroundings.

A large equilibrium constant (K) indicates that at equilibrium, the concentration of products is much greater than that of reactants, favoring products.

A negative ΔH indicates that the reaction releases heat to the surroundings, classifying it as exothermic.

A negative ΔH indicates that the reaction releases heat, classifying it as exothermic.

According to Gay-Lussac's Law, pressure is directly proportional to temperature at constant volume (P/T = k).

According to Gay-Lussac's Law, at constant volume, the pressure of a gas is directly proportional to its temperature in Kelvin.

ΔU = Q - W = 100 J - 40 J = 60 J.

If the internal energy of a closed system increases, it indicates that work is done on the system.

If the internal energy increases in a closed system, the enthalpy also increases, as ΔH = ΔU + PΔV.

According to Boyle's Law, if the volume is halved, the pressure will double (P1V1 = P2V2).

According to Boyle's Law, if the volume is halved, the pressure will double (P1V1 = P2V2).

According to Boyle's Law, decreasing the volume of a gas at constant temperature increases the pressure.

E = E° - (0.059/n)log([Ag⁺]1/[Ag⁺]2); E = 0 - (0.059/1)log(10) = 0.059 V.

E = (0.059 V/n) * log([Cathode]/[Anode]) = (0.059 V/2) * log(10) = 0.059 V.

E = (0.059 V/n) * log([Cathode]/[Anode]) = (0.059 V/2) * log(1/0.1) = 0.059 V * 1 = 0.059 V.

In a concentration cell, the cell potential is dependent on the concentration differences of the reactants.

In a concentration cell, the cell potential is generated due to differences in ion concentrations.

For a constant pressure process, the work done by the system is related to the change in enthalpy by the equation ΔH = Q + PΔV, where Q is the heat added.

The work done in a constant pressure process is calculated using W = PΔV, where P is pressure and ΔV is the change in volume.

The work done by the system at constant pressure is given by W = PΔV, where P is pressure and ΔV is the change in volume.