Chapter 7: Solution chemistry (C2075801)

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1 Ions in solution

Solution is a homogenous/complete mixture of two or more compounds, into only one phase. Homogeneity indicates that the proportion of the mixtures throughout the population, is consistent, with any sample. Although solution is only thought of between liquids, brass is a solution of copper and zinc.

“My favorite solution is the ice coffee,” Blaire said.

“A nice mix of coffee bean, sugar, and milk,” Mandy said, “but it’s not a perfect example of solution, because it’s likely just suspension and sedimentation.”

The solute is the substance dissolved into another substance, the solvent. The solvent is generally the substance that is in greater amount, but if two liquids are made solution in another, they can be both called solvent. The name matters very little.

“That’s sort of like how Jamie wants to call himself ‘the resident Christian bad boy’,” Mandy commented.

“He can call himself that, everybody can know himself by that, but the truth is he aren’t that bad…” Mandy giggled, “Looking!”

Ideal solution are solutions with properties analogous to ideal gases, with the exception that the intermolecular forces cannot be neglected [per ideal gases], but rather:

  • The bonds between the solute and solvent, are analogous to the bonds within the solute, and the bonds within the solvent
  • The size of the solute and solvent molecules, are analogous in size

For example, 1-butanol and 2-butanol create an almost ideal solution, because they are similar in size and bonding properties.

Ideally dilute solution is a solution that is so excessively dilute, the solute molecules are unable to interact with one another.

Non-ideal solutions are solutions that are neither ideal, nor ideally dilute.

Colloids [or colloidal suspensions] may have the appearance of solutions, but are distinct because the solute particles conglomerate to some degree. Gravity can cause an [unstable] colloid to sediment after time. Colloids can cause the Tyndall effect, which is the scattering of light by particles in the colloid, causing certain frequencies such as blue light to be more scattered than red light. Colloidal solutes can be extracted by pores of an ultrafiltration membrane, or alternatively, heating or chemical reaction thereby causing coagulation. It can also be separated in dialysis, which is the process of separating molecules by difference in their rates of diffusion through a semipermeable membrane.

2 Solubility

When a solute is mixed into solvent, it is solvated. Polar solvents solvate polar solutes, and non-polar solvents solvate non-polar solutes, which can be memorized with the mnemonic “like dissolves like”. This is because the intermolecular forces are electrostatic attraction, which are stronger in polar than non-polar. Therefore, polar and non-polar cannot mix, because the electrostatic force between polar molecules, is too great for non-polar to break into.

“That’s like how Mandy can never join the cool club,” Megan remarked, “because she can never break into that!”

“Ouch,” Jamie replied.

The process where the solvent surrounds the solute is known as a solvation shell. Where the solvent is water, it is referred to as a hydration shell. Aqueous solution is a solute with a solvation shell. The solvation number is the number of solvent molecules required to surround the solute molecule [or ion], depending on the charge and size of the molecule [or ion]. Where this occurs with water, it is known as the hydration number, which is commonly 4 or 6.

Common ions include:

Compounds which form ions in solution, are known as electrolytes. Electrolytes in aqueous solution, conduct electricity.

Concentration of solution can be measured in:

  • Molarity, which is [mathjax]c=\dfrac{n_{solute}}{V_{solvent}}[/mathjax], where [mathjax]n[/mathjax] is the moles of solute, and [mathjax]V[/mathjax] is the volume of solvent. The problem is that when temperature changes, the volume of solvent changes, due to thermal expansion. Therefore, a better measure may be–
  • Molality, which is [mathjax]b=\dfrac{n_{solute}}{m_{solvent}}[/mathjax], where [mathjax]m[/mathjax] is the mass of solvent. Evidently, molality doesn’t fluctuate with temperature. Molality is often used with colligative properties, which are properties that depend on number, and not kind of chemical species present ( discussed)
  • Mole fraction, which is [mathjax]x=\dfrac{n_{solute}}{n_{total}}[/mathjax], which also doesn’t fluctuate with temperature
  • Mass percent, which is the mass fraction as a percentage (i.e. multiplied by 100%). The mass fraction is [mathjax]w=\dfrac{m_{solute}}{m_{total}}[/mathjax]
  • Parts per million (PPM), which is mass fraction multiplied by a million. Note therefore, that PPM is based on mass, and not number of molecules
  • Normality, which is the equivalence per liter of solution. Because the definition of equivalence may not be unequivocal, IUPAC discourages the use of normality. Equivalence may be:
    • In acid-base chemistry, the mass of acid or base, able to donate or accept 1 mol of protons. For example, [mathjax]\ce{H2SO4}[/mathjax] is a 2N solution
    • In redox reactions, the quantity of oxidizing or reducing agent, that can accept or donate 1 mol of electrons

 Dissolution (process by which a solute forms a solution in a solvent) occurs in there steps:

  • The breaking of bonds, which requires energy, and therefore must be endothermic (positive [mathjax]\Delta H[/mathjax])
    • Solvent-solvent bonds are broken
    • Solute-solute bonds are broken
  • The forming of bonds, which releases energy, and therefore must be exothermic (negative [mathjax]\Delta H[/mathjax])
    • Solute-solvent bonds are formed

The enthalpy [or commonly, heat] of solution, is the sum of these there enthalpies ([mathjax]\Delta H[/mathjax]). If the enthalpy of solution is positive, the breaking of bonds requires more energy than the forming of bonds, meaning the new bonds are weaker than the old bonds. In contrast, if the enthalpy of solution is negative, the converse is true, meaning the new bonds are stronger than the old bonds. In accordance with the 2ndlaw of thermodynamics, the entropy of solution is always positive, because mixing two systems creates disorder.

Vapor pressure is the tendency of particles to escape from a liquid. Although particles escape, some will fall back down, the vapor pressure is the pressure of the vapor, whilst in equilibrium between these two phenomena. Water evaporates well below boiling temperature, for example, at room temperature, because although temperature relates to average kinetic energy [as stated ], there is an enormous range of kinetic energies, and some are moving fast enough to escape from the liquid. Note it makes sense that as temperature increases, average kinetic increases, and therefore vapor pressure.

Substances with high vapor pressures at normal temperatures are referred to as volatile. Assuming a volatile solvent, and a non-volatile solute, some of the non-volatile solute would block the volatile solvent from escaping, such that only a fraction can now escape. Raoult’s law states that [mathjax]p=p*x[/mathjax], where [mathjax]p[/mathjax] is the vapor pressure of the solution [with a non-volatile solute], [mathjax]p*[/mathjax] is the vapor pressure of the pure solvent, and [mathjax]x[/mathjax] is the mole fraction of the solvent. With a volatile solute, despite the volatile solute still blocks the volatile solvent from escaping, the volatile solute can itself escape. In this circumstance, Raoult’s law can be modified by just adding the partial effect of the solute, namely, [mathjax]p=p_{solvent}+p_{solute}[/mathjax]. Raoult’s law needs to be adapted to non-ideal solutions, by accounting for intermolecular forces. If the solute-solvent bonds are stronger, the vapor pressure would be lower; and if the solute-solvent bonds are weaker, the vapor pressure would be higher. Remember from  that enthalpy of solution provides information on strength comparison of the before- and after- bonds, such that if enthalpy of solution is negative, new bonds are stronger [than old bonds], causing a negative deviation from Raoult’s law, meaning that actual vapor pressure should be lower [than calculated]. In contrast, if enthalpy of solution is positive, actual vapor pressure should be higher.

Solubility is the tendency of a solute to solvate [in a solvent]. Where there is excess solute that doesn’t have an associated hydration shell, it will begin to precipitate. Saturation is the point at which precipitation begins to occur. The extent of solubility of a substance is measured as the concentration at saturation. Solubility product constant ([mathjax]K_{sp}[/mathjax]) is a constant specific to each solute and solvent, at a specific temperature. It is therefore like any other chemical equilibrium constant, but specifically refers to solute/solvent solution reactions. For example, [mathjax]Ca(OH)_{2} \rightleftarrows Ca^{2+} + 2OH^-[/mathjax]. The [mathjax]K_{sp}=[Ca^{2+}].[OH^-]^2[/mathjax], noting that [mathjax]Ca(OH)_2[/mathjax] doesn’t appear in the equation because it is a pure substance. Using a textbook value of [mathjax]K_{sp}[/mathjax], it is then possible to figure out the [mathjax][Ca^{2+}][/mathjax], the maximum concentration of calcium ions that can be dissolved, and thus [mathjax][Ca(OH)_{2}][/mathjax]. If substances are added as reactants, but don’t participate in the reaction, they are known as spectator ions. For example, if [mathjax]NaOH[/mathjax] is added to a fully saturated solution of [mathjax]Ca(OH)_{2}[/mathjax], the [mathjax]Na^+[/mathjax] ions are spectator ions, and the [mathjax]OH^-[/mathjax] ions would, by Chatelier’s principle, the equilibrium position will shift to counteract the effects of the disturbance, by shifting to the left, thereby causing precipitation of [mathjax]Ca(OH)_{2}[/mathjax]. This is the common ion effect, which is where the solubility of one salt is decreased, when another salt, which has an ion common with it, is added. A solubility chart indicates how well ions mix with other ions, and can be summarized as follows:

Factors that affect solubility, include pressure and temperature. In practice, pressure has negligible effect on solubility in condensed phases (solids or liquids). However, pressure does affect the solubility of a gas. Henry’s law states that in sufficiently dilute solutions, [mathjax]c=kp[/mathjax], where [mathjax]c[/mathjax] is solubility of gas, [mathjax]k[/mathjax] is Henry’s law constant, and [mathjax]p[/mathjax] is the [partial] pressure of gas, meaning solubility is directly proportional with pressure. This occurs in a soft drink can, where under pressure, the carbon dioxide is soluble in the soft drink. However, when the seal is broken, the carbon dioxide is insoluble, and released into the air. Solubility of solids increases with temperature, as enthalpy of solution requires quiet some energy to break the bonds of the reactants. However, solubility of gases decreases with temperature. This makes sense because gases expand with higher temperature, thereby leaving its solution.

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Pre-med science (MED5118352)


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Chapter 7: Solution chemistry - General chemistry - Pre-med science - MR. SHUM'S CLASSROOM