Solutions and Acids & Bases Definitions

Solubility and Dissociation in Chemistry

• Solvent molecules are attracted to surface ions, and each ion is surrounded by solvent molecules.
• Enthalpy changes (∆H) occur with each interaction broken or formed during the process of dissolving an ionic solid in water.
• The ions are solvated when surrounded by solvent, and if the solvent is water, the ions are hydrated, with ion-dipole intermolecular force.
• The enthalpy of solution, DHsoln, can be positive or negative and is only one side of the story (entropy).
• The degree of saturation determines whether a solution is unsaturated or saturated, and solubility is an equilibrium term between solid and saturated solution.
• For dissolution to happen, chemists use the axiom “like dissolves like,” where polar substances tend to dissolve in polar solvents and nonpolar substances tend to dissolve in nonpolar solvents.
• The solubility constant, Ks, is used to determine the equilibrium in a solution, and generally, the solubility of solid solutes in liquid solvents increases with increasing temperature.
• The common ion effect, based on Le Chatelier's principle, states that adding ions will affect the equilibrium constant.
• Water dissociates into ions, and the autoprotolysis constant of water is Kw = 10^-14.
• The pH scale measures the acidity or basicity of a solution, and a buffer solution can withstand small additions of acid or base without much change in pH.
• The dissociation of acids and bases is governed by equilibrium constants, and the Henderson-Hasselbalch equation relates pH, pKa, and ionization.
• The ratio of [A-] to [HA] in a solution can be determined using the Henderson-Hasselbalch equation, with practical applications in pharmaceuticals.

Solubility and Dissociation in Chemistry

• Solvent molecules are attracted to surface ions, and each ion is surrounded by solvent molecules.
• Enthalpy changes (∆H) occur with each interaction broken or formed during the process of dissolving an ionic solid in water.
• The ions are solvated when surrounded by solvent, and if the solvent is water, the ions are hydrated, with ion-dipole intermolecular force.
• The enthalpy of solution, DHsoln, can be positive or negative and is only one side of the story (entropy).
• The degree of saturation determines whether a solution is unsaturated or saturated, and solubility is an equilibrium term between solid and saturated solution.
• For dissolution to happen, chemists use the axiom “like dissolves like,” where polar substances tend to dissolve in polar solvents and nonpolar substances tend to dissolve in nonpolar solvents.
• The solubility constant, Ks, is used to determine the equilibrium in a solution, and generally, the solubility of solid solutes in liquid solvents increases with increasing temperature.
• The common ion effect, based on Le Chatelier's principle, states that adding ions will affect the equilibrium constant.
• Water dissociates into ions, and the autoprotolysis constant of water is Kw = 10^-14.
• The pH scale measures the acidity or basicity of a solution, and a buffer solution can withstand small additions of acid or base without much change in pH.
• The dissociation of acids and bases is governed by equilibrium constants, and the Henderson-Hasselbalch equation relates pH, pKa, and ionization.
• The ratio of [A-] to [HA] in a solution can be determined using the Henderson-Hasselbalch equation, with practical applications in pharmaceuticals.

Solubility and Dissociation in Chemistry

• Solvent molecules are attracted to surface ions, and each ion is surrounded by solvent molecules.
• Enthalpy changes (∆H) occur with each interaction broken or formed during the process of dissolving an ionic solid in water.
• The ions are solvated when surrounded by solvent, and if the solvent is water, the ions are hydrated, with ion-dipole intermolecular force.
• The enthalpy of solution, DHsoln, can be positive or negative and is only one side of the story (entropy).
• The degree of saturation determines whether a solution is unsaturated or saturated, and solubility is an equilibrium term between solid and saturated solution.
• For dissolution to happen, chemists use the axiom “like dissolves like,” where polar substances tend to dissolve in polar solvents and nonpolar substances tend to dissolve in nonpolar solvents.
• The solubility constant, Ks, is used to determine the equilibrium in a solution, and generally, the solubility of solid solutes in liquid solvents increases with increasing temperature.
• The common ion effect, based on Le Chatelier's principle, states that adding ions will affect the equilibrium constant.
• Water dissociates into ions, and the autoprotolysis constant of water is Kw = 10^-14.
• The pH scale measures the acidity or basicity of a solution, and a buffer solution can withstand small additions of acid or base without much change in pH.
• The dissociation of acids and bases is governed by equilibrium constants, and the Henderson-Hasselbalch equation relates pH, pKa, and ionization.
• The ratio of [A-] to [HA] in a solution can be determined using the Henderson-Hasselbalch equation, with practical applications in pharmaceuticals.

This quiz covers definitions related to solutions, solubility, solubility constant, common ion effect, pKw, Ka, Kb, pH, pOH, Henderson-Hasselbalch equation, ionization of weak acids and bases, and buffer solutions.