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2,901 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.003881 M | A strong monoprotic acid is fully dissociated. At concentration 0.003881 mol/L, [H⁺] = 0.003881 M and pH = −log₁₀[H⁺] = 2.411. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,902 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.023 M | A strong monoprotic acid is fully dissociated. At concentration 0.023 mol/L, [H⁺] = 0.023 M and pH = −log₁₀[H⁺] = 1.638. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,903 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 2.5941e-04 M | A strong monoprotic acid is fully dissociated. At concentration 2.5941e-04 mol/L, [H⁺] = 2.5941e-04 M and pH = −log₁₀[H⁺] = 3.586. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,904 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.001934 M | A strong monoprotic acid is fully dissociated. At concentration 0.001934 mol/L, [H⁺] = 0.001934 M and pH = −log₁₀[H⁺] = 2.714. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,905 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.05676 M | A strong monoprotic acid is fully dissociated. At concentration 0.05676 mol/L, [H⁺] = 0.05676 M and pH = −log₁₀[H⁺] = 1.246. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,906 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.04945 M | A strong monoprotic acid is fully dissociated. At concentration 0.04945 mol/L, [H⁺] = 0.04945 M and pH = −log₁₀[H⁺] = 1.306. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,907 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.0717 M | A strong monoprotic acid is fully dissociated. At concentration 0.0717 mol/L, [H⁺] = 0.0717 M and pH = −log₁₀[H⁺] = 1.144. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,908 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.003536 M | A strong monoprotic acid is fully dissociated. At concentration 0.003536 mol/L, [H⁺] = 0.003536 M and pH = −log₁₀[H⁺] = 2.452. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,909 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.01113 M | A strong monoprotic acid is fully dissociated. At concentration 0.01113 mol/L, [H⁺] = 0.01113 M and pH = −log₁₀[H⁺] = 1.953. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,910 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.0678 M | A strong monoprotic acid is fully dissociated. At concentration 0.0678 mol/L, [H⁺] = 0.0678 M and pH = −log₁₀[H⁺] = 1.169. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,911 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.002127 M | A strong monoprotic acid is fully dissociated. At concentration 0.002127 mol/L, [H⁺] = 0.002127 M and pH = −log₁₀[H⁺] = 2.672. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,912 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.02044 M | A strong monoprotic acid is fully dissociated. At concentration 0.02044 mol/L, [H⁺] = 0.02044 M and pH = −log₁₀[H⁺] = 1.69. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,913 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.01422 M | A strong monoprotic acid is fully dissociated. At concentration 0.01422 mol/L, [H⁺] = 0.01422 M and pH = −log₁₀[H⁺] = 1.847. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,914 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.003565 M | A strong monoprotic acid is fully dissociated. At concentration 0.003565 mol/L, [H⁺] = 0.003565 M and pH = −log₁₀[H⁺] = 2.448. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,915 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 1.6171e-04 M | A strong monoprotic acid is fully dissociated. At concentration 1.6171e-04 mol/L, [H⁺] = 1.6171e-04 M and pH = −log₁₀[H⁺] = 3.791. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,916 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.001035 M | A strong monoprotic acid is fully dissociated. At concentration 0.001035 mol/L, [H⁺] = 0.001035 M and pH = −log₁₀[H⁺] = 2.985. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,917 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.001603 M | A strong monoprotic acid is fully dissociated. At concentration 0.001603 mol/L, [H⁺] = 0.001603 M and pH = −log₁₀[H⁺] = 2.795. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,918 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.02075 M | A strong monoprotic acid is fully dissociated. At concentration 0.02075 mol/L, [H⁺] = 0.02075 M and pH = −log₁₀[H⁺] = 1.683. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,919 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.002005 M | A strong monoprotic acid is fully dissociated. At concentration 0.002005 mol/L, [H⁺] = 0.002005 M and pH = −log₁₀[H⁺] = 2.698. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,920 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.001513 M | A strong monoprotic acid is fully dissociated. At concentration 0.001513 mol/L, [H⁺] = 0.001513 M and pH = −log₁₀[H⁺] = 2.82. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,921 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 1.2921e-04 M | A strong monoprotic acid is fully dissociated. At concentration 1.2921e-04 mol/L, [H⁺] = 1.2921e-04 M and pH = −log₁₀[H⁺] = 3.889. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,922 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.05855 M | A strong monoprotic acid is fully dissociated. At concentration 0.05855 mol/L, [H⁺] = 0.05855 M and pH = −log₁₀[H⁺] = 1.232. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,923 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 2.8370e-04 M | A strong monoprotic acid is fully dissociated. At concentration 2.8370e-04 mol/L, [H⁺] = 2.8370e-04 M and pH = −log₁₀[H⁺] = 3.547. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,924 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.005293 M | A strong monoprotic acid is fully dissociated. At concentration 0.005293 mol/L, [H⁺] = 0.005293 M and pH = −log₁₀[H⁺] = 2.276. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,925 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.08574 M | A strong monoprotic acid is fully dissociated. At concentration 0.08574 mol/L, [H⁺] = 0.08574 M and pH = −log₁₀[H⁺] = 1.067. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,926 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.08712 M | A strong monoprotic acid is fully dissociated. At concentration 0.08712 mol/L, [H⁺] = 0.08712 M and pH = −log₁₀[H⁺] = 1.06. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,927 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.02496 M | A strong monoprotic acid is fully dissociated. At concentration 0.02496 mol/L, [H⁺] = 0.02496 M and pH = −log₁₀[H⁺] = 1.603. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,928 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.04241 M | A strong monoprotic acid is fully dissociated. At concentration 0.04241 mol/L, [H⁺] = 0.04241 M and pH = −log₁₀[H⁺] = 1.373. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,929 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.08279 M | A strong monoprotic acid is fully dissociated. At concentration 0.08279 mol/L, [H⁺] = 0.08279 M and pH = −log₁₀[H⁺] = 1.082. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,930 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.001811 M | A strong monoprotic acid is fully dissociated. At concentration 0.001811 mol/L, [H⁺] = 0.001811 M and pH = −log₁₀[H⁺] = 2.742. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,931 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.04543 M | A strong monoprotic acid is fully dissociated. At concentration 0.04543 mol/L, [H⁺] = 0.04543 M and pH = −log₁₀[H⁺] = 1.343. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,932 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.01266 M | A strong monoprotic acid is fully dissociated. At concentration 0.01266 mol/L, [H⁺] = 0.01266 M and pH = −log₁₀[H⁺] = 1.898. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,933 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.04175 M | A strong monoprotic acid is fully dissociated. At concentration 0.04175 mol/L, [H⁺] = 0.04175 M and pH = −log₁₀[H⁺] = 1.379. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,934 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.01461 M | A strong monoprotic acid is fully dissociated. At concentration 0.01461 mol/L, [H⁺] = 0.01461 M and pH = −log₁₀[H⁺] = 1.835. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,935 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 8.7314e-04 M | A strong monoprotic acid is fully dissociated. At concentration 8.7314e-04 mol/L, [H⁺] = 8.7314e-04 M and pH = −log₁₀[H⁺] = 3.059. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,936 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.001078 M | A strong monoprotic acid is fully dissociated. At concentration 0.001078 mol/L, [H⁺] = 0.001078 M and pH = −log₁₀[H⁺] = 2.968. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,937 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.03027 M | A strong monoprotic acid is fully dissociated. At concentration 0.03027 mol/L, [H⁺] = 0.03027 M and pH = −log₁₀[H⁺] = 1.519. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,938 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.05019 M | A strong monoprotic acid is fully dissociated. At concentration 0.05019 mol/L, [H⁺] = 0.05019 M and pH = −log₁₀[H⁺] = 1.299. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,939 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.00678 M | A strong monoprotic acid is fully dissociated. At concentration 0.00678 mol/L, [H⁺] = 0.00678 M and pH = −log₁₀[H⁺] = 2.169. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,940 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 2.7579e-04 M | A strong monoprotic acid is fully dissociated. At concentration 2.7579e-04 mol/L, [H⁺] = 2.7579e-04 M and pH = −log₁₀[H⁺] = 3.559. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,941 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.001629 M | A strong monoprotic acid is fully dissociated. At concentration 0.001629 mol/L, [H⁺] = 0.001629 M and pH = −log₁₀[H⁺] = 2.788. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,942 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.008544 M | A strong monoprotic acid is fully dissociated. At concentration 0.008544 mol/L, [H⁺] = 0.008544 M and pH = −log₁₀[H⁺] = 2.068. This relation follows directly from the definition of pH and the complete dissociation assumption. | pH = -log10 [H+] | mole_concept; logarithmic functions | Calculate the pH of a strong monoprotic acid solution. |
2,943 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 211.6 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 211.6. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,944 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 5.574 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 5.574. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,945 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 5.118 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 5.118. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,946 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.4496 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.4496. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse rea... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,947 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.1179 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.1179. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse rea... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,948 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 74.3 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 74.3. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse react... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,949 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 7.808 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 7.808. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,950 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 44.92 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 44.92. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,951 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 193.1 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 193.1. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,952 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.003204 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.003204. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse r... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,953 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 1.677 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 1.677. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,954 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.2188 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.2188. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse rea... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,955 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 20.57 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 20.57. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,956 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 190.2 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 190.2. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,957 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 1.327 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 1.327. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,958 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 1.757 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 1.757. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,959 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.01916 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.01916. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse re... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,960 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 1.175 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 1.175. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,961 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.07473 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.07473. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse re... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,962 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.1107 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.1107. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse rea... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,963 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 6.518 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 6.518. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,964 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.001382 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.001382. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse r... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,965 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 219.7 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 219.7. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,966 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 1.736 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 1.736. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,967 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.03566 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.03566. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse re... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,968 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 98.49 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 98.49. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,969 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.06663 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.06663. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse re... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,970 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.3439 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.3439. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse rea... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,971 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 1.076 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 1.076. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,972 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 24.66 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 24.66. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,973 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 112.1 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 112.1. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,974 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.03619 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.03619. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse re... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,975 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 1.109 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 1.109. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,976 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.005532 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.005532. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse r... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,977 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 34.62 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 34.62. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,978 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 444.7 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 444.7. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,979 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.4925 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.4925. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse rea... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,980 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 2.72 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 2.72. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse react... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,981 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.001075 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.001075. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse r... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,982 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.1081 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.1081. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse rea... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,983 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.0261 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.0261. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse rea... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,984 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 29.82 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 29.82. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,985 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 3.087 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 3.087. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,986 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 19.4 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 19.4. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse react... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,987 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 421.9 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 421.9. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,988 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 9.843 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 9.843. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,989 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 116.7 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 116.7. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,990 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.02529 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.02529. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse re... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,991 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.0624 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.0624. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse rea... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,992 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 10.04 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 10.04. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,993 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 77.71 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 77.71. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,994 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 2.733 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 2.733. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,995 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.006916 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.006916. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse r... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,996 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.1845 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.1845. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse rea... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,997 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 26.61 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 26.61. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,998 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.004279 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.004279. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse r... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
2,999 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 1.255 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 1.255. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reac... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
3,000 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.0143 | For a reversible reaction at a fixed temperature, the equilibrium constant K is a thermodynamic quantity determined solely by the standard Gibbs free-energy change: K = exp(−ΔG° / R T). In the present illustration K = 0.0143. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse rea... | K = exp(-ΔG° / R T); ΔG = ΔG° + R T ln Q | thermodynamics_first_law; mole_concept | Interpret the magnitude of an equilibrium constant and its relation to ΔG°. |
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