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20,133,001 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 4.3607e-04 M | A strong monoprotic acid is fully dissociated. At concentration 4.3607e-04 mol/L, [H⁺] = 4.3607e-04 M and pH = −log₁₀[H⁺] = 3.36. 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. |
20,133,002 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 1.5911e-04 M | A strong monoprotic acid is fully dissociated. At concentration 1.5911e-04 mol/L, [H⁺] = 1.5911e-04 M and pH = −log₁₀[H⁺] = 3.798. 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. |
20,133,003 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.03429 M | A strong monoprotic acid is fully dissociated. At concentration 0.03429 mol/L, [H⁺] = 0.03429 M and pH = −log₁₀[H⁺] = 1.465. 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. |
20,133,004 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.01993 M | A strong monoprotic acid is fully dissociated. At concentration 0.01993 mol/L, [H⁺] = 0.01993 M and pH = −log₁₀[H⁺] = 1.7. 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. |
20,133,005 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 5.5603e-04 M | A strong monoprotic acid is fully dissociated. At concentration 5.5603e-04 mol/L, [H⁺] = 5.5603e-04 M and pH = −log₁₀[H⁺] = 3.255. 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. |
20,133,006 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.01639 M | A strong monoprotic acid is fully dissociated. At concentration 0.01639 mol/L, [H⁺] = 0.01639 M and pH = −log₁₀[H⁺] = 1.785. 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. |
20,133,007 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.001407 M | A strong monoprotic acid is fully dissociated. At concentration 0.001407 mol/L, [H⁺] = 0.001407 M and pH = −log₁₀[H⁺] = 2.852. 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. |
20,133,008 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 7.4714e-04 M | A strong monoprotic acid is fully dissociated. At concentration 7.4714e-04 mol/L, [H⁺] = 7.4714e-04 M and pH = −log₁₀[H⁺] = 3.127. 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. |
20,133,009 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.01887 M | A strong monoprotic acid is fully dissociated. At concentration 0.01887 mol/L, [H⁺] = 0.01887 M and pH = −log₁₀[H⁺] = 1.724. 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. |
20,133,010 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.003212 M | A strong monoprotic acid is fully dissociated. At concentration 0.003212 mol/L, [H⁺] = 0.003212 M and pH = −log₁₀[H⁺] = 2.493. 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. |
20,133,011 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.004185 M | A strong monoprotic acid is fully dissociated. At concentration 0.004185 mol/L, [H⁺] = 0.004185 M and pH = −log₁₀[H⁺] = 2.378. 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. |
20,133,012 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.01435 M | A strong monoprotic acid is fully dissociated. At concentration 0.01435 mol/L, [H⁺] = 0.01435 M and pH = −log₁₀[H⁺] = 1.843. 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. |
20,133,013 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.03322 M | A strong monoprotic acid is fully dissociated. At concentration 0.03322 mol/L, [H⁺] = 0.03322 M and pH = −log₁₀[H⁺] = 1.479. 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. |
20,133,014 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.03286 M | A strong monoprotic acid is fully dissociated. At concentration 0.03286 mol/L, [H⁺] = 0.03286 M and pH = −log₁₀[H⁺] = 1.483. 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. |
20,133,015 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.003989 M | A strong monoprotic acid is fully dissociated. At concentration 0.003989 mol/L, [H⁺] = 0.003989 M and pH = −log₁₀[H⁺] = 2.399. 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. |
20,133,016 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 1.9526e-04 M | A strong monoprotic acid is fully dissociated. At concentration 1.9526e-04 mol/L, [H⁺] = 1.9526e-04 M and pH = −log₁₀[H⁺] = 3.709. 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. |
20,133,017 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 8.9721e-04 M | A strong monoprotic acid is fully dissociated. At concentration 8.9721e-04 mol/L, [H⁺] = 8.9721e-04 M and pH = −log₁₀[H⁺] = 3.047. 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. |
20,133,018 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.001108 M | A strong monoprotic acid is fully dissociated. At concentration 0.001108 mol/L, [H⁺] = 0.001108 M and pH = −log₁₀[H⁺] = 2.956. 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. |
20,133,019 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 4.9840e-04 M | A strong monoprotic acid is fully dissociated. At concentration 4.9840e-04 mol/L, [H⁺] = 4.9840e-04 M and pH = −log₁₀[H⁺] = 3.302. 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. |
20,133,020 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.00218 M | A strong monoprotic acid is fully dissociated. At concentration 0.00218 mol/L, [H⁺] = 0.00218 M and pH = −log₁₀[H⁺] = 2.661. 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. |
20,133,021 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.001284 M | A strong monoprotic acid is fully dissociated. At concentration 0.001284 mol/L, [H⁺] = 0.001284 M and pH = −log₁₀[H⁺] = 2.892. 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. |
20,133,022 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.003801 M | A strong monoprotic acid is fully dissociated. At concentration 0.003801 mol/L, [H⁺] = 0.003801 M and pH = −log₁₀[H⁺] = 2.42. 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. |
20,133,023 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.01304 M | A strong monoprotic acid is fully dissociated. At concentration 0.01304 mol/L, [H⁺] = 0.01304 M and pH = −log₁₀[H⁺] = 1.885. 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. |
20,133,024 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.001673 M | A strong monoprotic acid is fully dissociated. At concentration 0.001673 mol/L, [H⁺] = 0.001673 M and pH = −log₁₀[H⁺] = 2.777. 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. |
20,133,025 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 2.7450e-04 M | A strong monoprotic acid is fully dissociated. At concentration 2.7450e-04 mol/L, [H⁺] = 2.7450e-04 M and pH = −log₁₀[H⁺] = 3.561. 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. |
20,133,026 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.0662 M | A strong monoprotic acid is fully dissociated. At concentration 0.0662 mol/L, [H⁺] = 0.0662 M and pH = −log₁₀[H⁺] = 1.179. 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. |
20,133,027 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.03029 M | A strong monoprotic acid is fully dissociated. At concentration 0.03029 mol/L, [H⁺] = 0.03029 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. |
20,133,028 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.02842 M | A strong monoprotic acid is fully dissociated. At concentration 0.02842 mol/L, [H⁺] = 0.02842 M and pH = −log₁₀[H⁺] = 1.546. 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. |
20,133,029 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.001837 M | A strong monoprotic acid is fully dissociated. At concentration 0.001837 mol/L, [H⁺] = 0.001837 M and pH = −log₁₀[H⁺] = 2.736. 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. |
20,133,030 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 1.8403e-04 M | A strong monoprotic acid is fully dissociated. At concentration 1.8403e-04 mol/L, [H⁺] = 1.8403e-04 M and pH = −log₁₀[H⁺] = 3.735. 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. |
20,133,031 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.02622 M | A strong monoprotic acid is fully dissociated. At concentration 0.02622 mol/L, [H⁺] = 0.02622 M and pH = −log₁₀[H⁺] = 1.581. 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. |
20,133,032 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.03552 M | A strong monoprotic acid is fully dissociated. At concentration 0.03552 mol/L, [H⁺] = 0.03552 M and pH = −log₁₀[H⁺] = 1.45. 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. |
20,133,033 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.01248 M | A strong monoprotic acid is fully dissociated. At concentration 0.01248 mol/L, [H⁺] = 0.01248 M and pH = −log₁₀[H⁺] = 1.904. 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. |
20,133,034 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.002455 M | A strong monoprotic acid is fully dissociated. At concentration 0.002455 mol/L, [H⁺] = 0.002455 M and pH = −log₁₀[H⁺] = 2.61. 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. |
20,133,035 | chemistry | acids_bases | strong_acid_ph | 4 | worked_example | pH of strong acid at concentration 0.04291 M | A strong monoprotic acid is fully dissociated. At concentration 0.04291 mol/L, [H⁺] = 0.04291 M and pH = −log₁₀[H⁺] = 1.367. 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. |
20,133,036 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 2.244 | 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 = 2.244. 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°. |
20,133,037 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.004942 | 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 = 0.004942. 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°. |
20,133,038 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.005175 | 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 = 0.005175. 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°. |
20,133,039 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 138.2 | 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 = 138.2. 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°. |
20,133,040 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.1534 | 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 = 0.1534. 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°. |
20,133,041 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.1161 | 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 = 0.1161. 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°. |
20,133,042 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.00468 | 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 = 0.00468. 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°. |
20,133,043 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.06427 | 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 = 0.06427. 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°. |
20,133,044 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.04676 | 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 = 0.04676. 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°. |
20,133,045 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.003225 | 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 = 0.003225. 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°. |
20,133,046 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.2828 | 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 = 0.2828. 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°. |
20,133,047 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 3.561 | 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 = 3.561. 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°. |
20,133,048 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.07372 | 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 = 0.07372. 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°. |
20,133,049 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.4199 | 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 = 0.4199. 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°. |
20,133,050 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 2.152 | 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 = 2.152. 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°. |
20,133,051 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.001803 | 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 = 0.001803. 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°. |
20,133,052 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 2.638 | 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 = 2.638. 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°. |
20,133,053 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.6218 | 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 = 0.6218. 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°. |
20,133,054 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.05675 | 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 = 0.05675. 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°. |
20,133,055 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.9057 | 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 = 0.9057. 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°. |
20,133,056 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 2.087 | 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 = 2.087. 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°. |
20,133,057 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 974.7 | 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 = 974.7. 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°. |
20,133,058 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 6.963 | 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 = 6.963. 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°. |
20,133,059 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 137.9 | 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 = 137.9. 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°. |
20,133,060 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 58 | 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 = 58. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reactio... | 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°. |
20,133,061 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.002748 | 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 = 0.002748. 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°. |
20,133,062 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 19.67 | 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 = 19.67. 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°. |
20,133,063 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 15.5 | 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 = 15.5. 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°. |
20,133,064 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 273.5 | 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 = 273.5. 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°. |
20,133,065 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 525.3 | 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 = 525.3. 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°. |
20,133,066 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.003984 | 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 = 0.003984. 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°. |
20,133,067 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 19.03 | 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 = 19.03. 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°. |
20,133,068 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.1849 | 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 = 0.1849. 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°. |
20,133,069 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.04131 | 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 = 0.04131. 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°. |
20,133,070 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.003101 | 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 = 0.003101. 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°. |
20,133,071 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.06497 | 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 = 0.06497. 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°. |
20,133,072 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.04736 | 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 = 0.04736. 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°. |
20,133,073 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.0125 | 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 = 0.0125. 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°. |
20,133,074 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.02565 | 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 = 0.02565. 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°. |
20,133,075 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.001009 | 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 = 0.001009. 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°. |
20,133,076 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 178.4 | 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 = 178.4. 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°. |
20,133,077 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.1411 | 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 = 0.1411. 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°. |
20,133,078 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.01057 | 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 = 0.01057. 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°. |
20,133,079 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 1.399 | 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 = 1.399. 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°. |
20,133,080 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.9973 | 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 = 0.9973. 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°. |
20,133,081 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 4.453 | 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 = 4.453. 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°. |
20,133,082 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.01101 | 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 = 0.01101. 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°. |
20,133,083 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 73 | 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 = 73. When Q (reaction quotient) < K the forward reaction is spontaneous; when Q > K the reverse reactio... | 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°. |
20,133,084 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 98.57 | 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 = 98.57. 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°. |
20,133,085 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.03458 | 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 = 0.03458. 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°. |
20,133,086 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 20.35 | 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 = 20.35. 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°. |
20,133,087 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.2151 | 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 = 0.2151. 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°. |
20,133,088 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 26.28 | 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 = 26.28. 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°. |
20,133,089 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.2177 | 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 = 0.2177. 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°. |
20,133,090 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.004417 | 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 = 0.004417. 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°. |
20,133,091 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 303.3 | 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 = 303.3. 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°. |
20,133,092 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.001938 | 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 = 0.001938. 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°. |
20,133,093 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.3568 | 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 = 0.3568. 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°. |
20,133,094 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 24.42 | 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 = 24.42. 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°. |
20,133,095 | chemistry | equilibrium | equilibrium_constant | 6 | explanation | Meaning of equilibrium constant K = 0.09013 | 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 = 0.09013. 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°. |
20,133,096 | chemistry | thermochemistry | hess_law | 5 | explanation | Hess's law and enthalpy as a state function | Enthalpy H is a state function: its change between two states is independent of path. Consequently, the enthalpy change of a reaction may be computed by summing the enthalpy changes of any convenient sequence of reactions that net to the same overall transformation (Hess's law). Standard enthalpies of formation provide... | ΔH°_rxn = Σ ΔH°_f(products) - Σ ΔH°_f(reactants) | first_law | Apply Hess's law to compute reaction enthalpies from tabulated formation data. |
20,133,097 | chemistry | thermochemistry | hess_law | 5 | explanation | Hess's law and enthalpy as a state function | Enthalpy H is a state function: its change between two states is independent of path. Consequently, the enthalpy change of a reaction may be computed by summing the enthalpy changes of any convenient sequence of reactions that net to the same overall transformation (Hess's law). Standard enthalpies of formation provide... | ΔH°_rxn = Σ ΔH°_f(products) - Σ ΔH°_f(reactants) | first_law | Apply Hess's law to compute reaction enthalpies from tabulated formation data. |
20,133,098 | chemistry | thermochemistry | hess_law | 5 | explanation | Hess's law and enthalpy as a state function | Enthalpy H is a state function: its change between two states is independent of path. Consequently, the enthalpy change of a reaction may be computed by summing the enthalpy changes of any convenient sequence of reactions that net to the same overall transformation (Hess's law). Standard enthalpies of formation provide... | ΔH°_rxn = Σ ΔH°_f(products) - Σ ΔH°_f(reactants) | first_law | Apply Hess's law to compute reaction enthalpies from tabulated formation data. |
20,133,099 | chemistry | thermochemistry | hess_law | 5 | explanation | Hess's law and enthalpy as a state function | Enthalpy H is a state function: its change between two states is independent of path. Consequently, the enthalpy change of a reaction may be computed by summing the enthalpy changes of any convenient sequence of reactions that net to the same overall transformation (Hess's law). Standard enthalpies of formation provide... | ΔH°_rxn = Σ ΔH°_f(products) - Σ ΔH°_f(reactants) | first_law | Apply Hess's law to compute reaction enthalpies from tabulated formation data. |
20,133,100 | chemistry | thermochemistry | hess_law | 5 | explanation | Hess's law and enthalpy as a state function | Enthalpy H is a state function: its change between two states is independent of path. Consequently, the enthalpy change of a reaction may be computed by summing the enthalpy changes of any convenient sequence of reactions that net to the same overall transformation (Hess's law). Standard enthalpies of formation provide... | ΔH°_rxn = Σ ΔH°_f(products) - Σ ΔH°_f(reactants) | first_law | Apply Hess's law to compute reaction enthalpies from tabulated formation data. |
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