Draft Calculations/ Topic Report draft

Unable to publish calculations due to inability of blogger to paste in equation format.
Topic Report drafting-

Introduction-

This extended experimental investigation is aimed at investigating the scientific theory of ‘like dissolves like.’ In order to thoroughly examine this topic and gain both qualitative and quantitative evidence, the solubility constant of the salt sodium chloride (NaCl) will be investigated in both polar (water and ethanol) and non-polar (oil) substances. A Solubility constant is referring to the equilibrium of a solid and its corresponding ions in a saturated solution. The higher the concentration of ions from the solid dissociated in the solution, the greater the solubility constant. The overall solubility of the solute is dependent upon the level of the K_sp (solubility constant). Additionally, the polarity of the solvent is also an integral factor in determining the solubility constant of the solute, in this case Sodium Chloride. Additionally, the solution must be saturated in order to gain an accurate measure of the Ksp. Temperature is also an experimental factor investigated, though its ability to either negate or increase the effects of polarity will be the main focus. Both the polarity of the solvent and the solubility product of the Sodium Chloride are relatable to the theory ‘like dissolves like,’ which states that the solubility or miscibility of a product depends on the degree to which polar substances dissolve polar substances and non-polar substances dissolve non-polar substances.

Discussion-

Ever since scientists first developed the central idea of ‘like dissolves like,’ researchers world-wide have based and implemented their experiments generally according to this guideline. However, a metaphorical ‘grey area’ which still exists in comparison to other theories is the extent to which like dissolves like. As mentioned, there is a key component being investigated which could determine the level of dissolution of similar substances- solubility product, in terms of polarity of the solution and the temperature.  The substances being used to dissolve the sodium chloride are water, ethanol and oil.

Water makes up around 70% of the surface of the earth, either in the form of salt or fresh water (Chemistry in Use 1, 2006). Water is unusual in terms of its properties, even though each molecule is made up of the relatively simple structure of two hydrogen atoms and one oxygen atom. Compared to other similar molecules, water has an abnormal boiling and freezing point, as well as high surface tension, cohesion and heat of vaporisation. However, the polarity of both the intra- molecular and inter-molecular forces within water can be considered the cause for these anomalies. Each hydrogen atom within a water molecule is covalently bonded to the oxygen molecule. Oxygen, having space for two more valency electrons in its outer shell, and each hydrogen having one space available in its outer shell, allows this bonding to occur. Oxygen has the highest electronegativity (ability to ‘hold’ electrons) out of the two atoms in the bond, therefore it has a greater attachment to the shared electrons. This attraction creates a permanent dipole, where the oxygen atom is always slightly negatively charged and the hydrogen is slightly more positive. The molecule consequently becomes arranged as a bent structure, as the positive hydrogen atoms repel each other. The polarity intra-molecularly increases the inter-molecular forces and gives water a greater cohesion. In water, this attraction between the highly electronegative oxygen atom and the corresponding hydrogen atoms, as well as the inter-molecular forces between oxygen and hydrogen atoms from different molecules is known as hydrogen bonding (see figure below).

 

 

 

 

This polarity and high cohesion of water helps determine the solubility of different substances. According to the theory ‘like dissolves like,’ water should be able to act as a solvent for polar molecules (i.e. substances with a permanent dipole and to a lesser extent an instantaneous dipole). The greater the polarity of the solute, the greater the ability of the molecules within that solute to ‘break’ the hydrogen bonding between water molecules.

High temperature also increases the ability of a solute to dissolve in water.  When water is heated, the molecular energy of the water molecules is increased, meaning that movement is faster and collisions between molecules contain more force. If the heat energy is great enough, then the hydrogen bonds between the water molecules will break and the molecules will move closer to a gaseous state. Through this increase in molecular motion, solutes with lower polarity in relation to water are able to dissolve more easily due there being more space between the water molecules (Does temperature affect dissolving, n/d). However, water (regarded as the universal solvent) does not contain the same solubility properties as other substances.

Ethanol (ethyl alcohol-CH₃CH₂OH) is one such substance very different to water in terms of its ability to dissolve polar solutes (such as NaCl). The hydroxyl group (OH group) present in ethanol signifies that the molecule has a certain degree of polarity. To what degree is yet to be determined in the experiment, however due to the fact that ethanol is a hydrocarbon and the length of the ‘R’ group is large in comparison to the OH group, this would decrease polarity and consequently solubility of polar substances. The oxygen atom present in the hydroxyl group has a high degree of electronegativity, therefore enabling other polar substances to dissolve. There is only one oxygen atom per ethanol molecule, however; therefore, the overall polarity is reduced from if ethanol was a smaller molecule. Due to the nature of the ethanol molecule, it is also capable of dissolving other hydrocarbons. This is also indicative of the theory ‘like dissolves like,’ as the non-polar component of ethanol (CH₃CH₂) has the ability to dissolve other non-polar substances.

Temperature increases the rate of solubility of either a polar or non-polar substance within ethanol. The higher the temperature, the greater the motion of the molecules, which weakens the dispersion forces inter-molecularly. This allows the solute to disperse more easily and to a greater extent throughout the solution. Although the polarity and consequent solubility of ethanol is much less than that of water, it is a more versatile solvent and still dissolves polar substances better than completely non-polar substance like oil.

Olive oil is a tri-glyceride made up of a variety of different fatty acids and is predominantly considered to be not polar. Tri-glyceride is a specific type of lipid, which is also known as fats. Oil contains molecules which are bent and irregular, therefore weakening the already poor dispersion forces between molecules. When a polar solute is added to the liquid, the oil molecules tend to clump together and rise to the top. The molecules clump together because the dispersion forces of oil are too weak to force apart the strong dispersion forces or dipoles which hold together polar substances. Additionally, the polar molecules are not attracted to the oil molecules because they do not hold a significant charge, let alone any highly electronegative atom. However, in the case of an ionic substance being used as the solute, a very small dipole is created when the charged particles are added to the oil, bonding the two molecules together briefly (Intermolecular and Interatomic particles, n/d). Temperature has no effect on the solubility of polar substances in oil, as no matter how much energy is injected into molecule motion and which intermolecular bonds are broken, the oil will still end up clumped together (because it has nothing to keep it attracted to polar molecules).

Sodium Chloride (NaCl) was the ionic (and consequently very polar) substance used throughout the following investigations to determine its solubility product when added to the aforementioned solvents. NaCl is essentially two charged ions. A positively charged sodium ion and a negatively charged chlorine ion combined together to cancel each other out and create a neutral compound. When sodium chloride is added to water, the sodium and chloride ions (through the strength of the hydrogen bonding in water over the ionic bonding in NaCl) dissociate and become attracted to their oppositely charged particles in water. The chlorine ion attracts two hydrogen atoms and the sodium attracts the oxygen in a process known as ion-dipole interaction (Intermolecular and Interatomic particles, n/d). At 20 degrees, a saturated solution is created when there is 35.7g of salt per 100mL. When sodium chloride dissociates in ethanol the procedure works much the same way, except there is only one OH group so limited ion-dipole interactions can occur. Therefore, the consequent solubility product for a solution containing ethanol and salt should be less than that of water and salt. Sodium Chloride in oil should have a solubility product equal to zero, even though the sodium chloride will create a very small dipole upon contact with the oil molecules. The formula for calculating the solubility product of NaCl is:

 Solubility product itself is independent of temperature. The rate of the dissolution of a substance in a solvent, however, is altered depending on the amount of kinetic energy available to molecules. If time constraints are added, then the solubility product could differ depending on the amount of time the sodium chloride ions take to dissociate in solution. If there was a known amount of NaCl within a solution, then the amount of ions in the solution would not change unless a greater proportion of the precipitate which was in equilibrium dissolved. However, the solution would then face the possibility of becoming super-saturated.  If the Ksp is equal to one, an equilibrium is present between the ions and the precipitate. The Q value, however, can be above or below the equilibrium constant if the substance is not at equilibrium. A higher Ksp value indicates that more ions have dissociated in comparison to amount of precipitate formed, while a low Ksp value means there is a higher relative concentration of the undissolved compound.

In order to determine the amount of chloride ions in a solution of sodium chloride in either water or ethanol, a titration must occur. Normally, silver nitrate (AgNO₃) is titrated when investigating salinity, and a potassium dichromate solution is used as the indicator (Salt concentration by titration, n/d). The substances react in the titration according to the following equations (the second equation being with respect to the indicator):

When the brick red solid precipitate is formed, the reaction is finished. In order to determine the solubility product of the salt, the chlorine concentration must first be determined.

Final Results

Table 1- Titration of NaCl against 1M Silver Nitrate in water

Temperature (°C)	Amount of NaCl (g)	Amount of water used to dissolve salt (mL)	% NaCl solution used in titration	Amount of AgNO3 used in titration (mL)	Observations
1) ≈ 25 (room temperature)	35.7	135	7.4	23.68 ≈ 23.7
2) 50	35.7	135	7.4	29.98 ≈ 30.0	Solution left for two days before used.

 

Table 2- Titration of NaCl against 1M Silver Nitrate in ethanol

Temperature (°C)	Amount of NaCl (g)	Amount of water used to dissolve salt (mL)	% NaCl solution used in titration	Amount of AgNO3 used in titration (mL)	Observations
3) ≈ 25 (room temperature)	5.2	35	100	4.54	Solution used 0.998M instead of 1M
4) 50	5.2	35	100	3.93

 

Table 3- Filtration of oil and NaCl and corresponding weights in order to find if any NaCl had dissolved

Temperature (°C)	Amount of NaCl (g)	Amount of Oil (mL)	Mass of beaker with 50mL water (g)	Mass of beaker with 50mL water + additional NaCl (g)	Difference between both masses (g)	Observations
5) ≈ 25 (room temperature)	5	10	47.8	52.8	5
6) 50	5	10	48.1	53.2	5.1	Not possible, as only 5g of NaCl was applied.

More research

http://en.wikipedia.org/wiki/Solubility

Factors affecting solubility [edit]

Solubility is defined for specific phases. For example, the solubility of aragonite and calcite in water are expected to differ, even though they are both polymorphs of calcium carbonate and have the same chemical formula.
The solubility of one substance in another is determined by the balance of intermolecular forces between the solvent and solute, and the entropy change that accompanies the solvation. Factors such as temperature and pressure will alter this balance, thus changing the solubility.
Solubility may also strongly depend on the presence of other species dissolved in the solvent, for example, complex-forming anions (ligands) in liquids. Solubility will also depend on the excess or deficiency of a common ion in the solution, a phenomenon known as the common-ion effect. To a lesser extent, solubility will depend on the ionic strength of solutions. The last two effects can be quantified using the equation for solubility equilibrium.
For a solid that dissolves in a redox reaction, solubility is expected to depend on the potential (within the range of potentials under which the solid remains the thermodynamically stable phase). For example, solubility of gold in high-temperature water is observed to be almost an order of magnitude higher when the redox potential is controlled using a highly oxidizing Fe₃O₄-Fe₂O₃ redox buffer than with a moderately oxidizing Ni-NiO buffer.^[4]

Solubility (metastable) also depends on the physical size of the crystal or droplet of solute (or, strictly speaking, on the specific surface area or molar surface area of the solute). For quantification, see the equation in the article on solubility equilibrium. For highly defective crystals, solubility may increase with the increasing degree of disorder. Both of these effects occur because of the dependence of solubility constant on the Gibbs energy of the crystal. The last two effects, although often difficult to measure, are of practical importance.^{[citation needed]} For example, they provide the driving force for precipitate aging (the crystal size spontaneously increasing with time).

Temperature [edit]

The solubility of a given solute in a given solvent typically depends on temperature. For many solids dissolved in liquid water, the solubility increases with temperature up to 100 °C.^[5] In liquid water at high temperatures, (e.g., that approaching the critical temperature), the solubility of ionic solutes tends to decrease due to the change of properties and structure of liquid water; the lower dielectric constant results in a less polar solvent.
Gaseous solutes exhibit more complex behavior with temperature. As the temperature is raised, gases usually become less soluble in water (to minimum, which is below 120 °C for most permanent gases^[6]), but more soluble in organic solvents.^[5]
The chart shows solubility curves for some typical solid inorganic salts (temperature is in degrees Celsius).^[7] Many salts behave like barium nitrate and disodium hydrogen arsenate, and show a large increase in solubility with temperature. Some solutes (e.g., sodium chloride in water) exhibit solubility that is fairly independent of temperature. A few, such as cerium(III) sulfate, become less soluble in water as temperature increases. This temperature dependence is sometimes referred to as "retrograde" or "inverse" solubility. Occasionally, a more complex pattern is observed, as with sodium sulfate, where the less soluble decahydrate crystal loses water of crystallization at 32 °C to form a more soluble anhydrous phase.^{[citation needed]}

The solubility of organic compounds nearly always increases with temperature. The technique of recrystallization, used for purification of solids, depends on a solute's different solubilities in hot and cold solvent. A few exceptions exist, such as certain cyclodextrins.^[8]

Solubility constants are used to describe saturated solutions of ionic compounds of relatively low solubility (see solubility equilibrium). The solubility constant is a special case of an equilibrium constant. It describes the balance between dissolved ions from the salt and undissolved salt. The solubility constant is also "applicable" (i.e., useful) to precipitation, the reverse of the dissolving reaction. As with other equilibrium constants, temperature can affect the numerical value of solubility constant. The solubility constant is not as simple as solubility, however the value of this constant is generally independent of the presence of other species in the solvent.
The Flory-Huggins solution theory is a theoretical model describing the solubility of polymers. The Hansen Solubility Parameters and the Hildebrand solubility parameters are empirical methods for the prediction of solubility. It is also possible to predict solubility from other physical constants such as the enthalpy of fusion.
The partition coefficient (Log P) is a measure of differential solubility of a compound in a hydrophobic solvent (octanol) and a hydrophilic solvent (water). The logarithm of these two values enables compounds to be ranked in terms of hydrophilicity (or hydrophobicity).
The energy change associated with dissolving is usually given per mole of solute as the enthalpy of solution.

Research for Hypothesis/Introduction finally added

http://www.chem.tamu.edu/class/majors/tutorialnotefiles/ksp.htm

Solubility equilibrium is base on the assumption that solids dissolve in water to give the basic particles from which they are formed.

·  Molecular solids dissolve to give individual aqueous molecules.

·  Ionic solids dissociate to give their respective positive and negative ions:

The ions in formed from the dissociation of ionic solids can carry an electrical current. Salt solutions, therefore, are good conductors of electricity. Molecular solids, however, do not dissociate in water to give ions, so no electrical current can be carried.

Solubility

1.      The ratio of the maximum amount of solute to the volume of solvent in which this solute can dissolve.

1.      Generally expressed in two ways:

1.      grams of solute per 100 g of water

2.      moles of solute per Liter of solution

Soluble: Dissolve - Do NOT form a solid precipitate. **alkali metal ions and ammonium ion: Li+, Na+, K+, NH4+ acetate ion: C2H3O21- nitrate ion: NO31- halide ions (X): Cl-, Br-, I- (Exceptions: AgX, HgX, and PbX2 are insoluble) sulfate ion: SO42- (Exceptions: SrSO4, BaSO4, and PbSO4 are insoluble; AgSO4, CaSO4, and Hg2SO4 are slightly soluble)

 

Insoluble: Do NOT Dissolve - Do form a solid precipitate. carbonate ion: CO32- chromate ion: CrO42- phosphate ion: PO43- sulfide ion: S2- (Exceptions: CaS, SrS, and BaS are soluble) hydroxide ion: OH- (Exceptions: Sr(OH)2 and Ba(OH)2 are soluble; Ca(OH)2 is slightly soluble)

• A salt is considered soluble if it dissolves in water to give a solution with a concentration of at least 0.1 M at room temperature.
• A salt is considered insoluble if the concentration of an aqueous solution is less than 0.0001 M at room temperature.
• Salts with solubilities between 0.0001 M and 0.1 M are considered to be slightly soluble.

Salts that have extremely low solubilities dissociate in water according to the principles of equilibrium. For example, the reaction for the dissociation of the salt AgCl is:

The reverse reaction for the dissolving of the salt would be the precipitation of the ions to form a solid:

The system has reached equilibrium when the rate at which AgCl dissolves is equal to the rate at which AgCl precipitates.

• Saturated solution - Contains the maximum concentration of ions that can exist in equilibrium with the solid salt at a given temperature.

The equilibrium reaction for the dissociation of AgCl is:

Solubility product equilibrium constant (K_sp) - The product of the equilibrium concentrations of the ions in a saturated solution of a salt. Each concentration is raised to the power of the respective coefficient of ion in the balanced equation.

• NOTE: There is no denominator in the solubility product equilibrium constant. The key word to remember is PRODUCT which can remind you that you should have a multiplication (or product) of the concentrations of the ions. The reason that the solid reactant is not written is because its concentration effectively remains constant.

For example, the solubility product equilibrium constant for the dissociation of AgCl is:

Let's try another example of a solubility product equilibrium constant. Consider the reaction for the dissociation of CaF₂ in water:

The solubility product equilibrium constant for this reaction would be the product of the concentration of Ca²⁺ ion and the concentration of the F^- ion raised to the second power (squared):

NOTE: Unlike K_a and K_b for acids and bases, the relative values of K_sp cannot be used to predict the relative solubilities of salts if the salts being compared produce a different number of ions.

The solubility product is literally the product of the solubilities of the ions in units of molarity (mol/L)

Sample Calculations

• K_sp can be calculated from the solubility of a salt. Conversely, the solubility of a salt can be calculated from K_sp.

Let's try an example calculation problem to demonstrate the relationship between the solubility and the solubility product of a salt.

1) Calculate the solubility of CaF₂ in g/L (K_sp = 4.0 x 10^-8)

First, write the BALANCED REACTION:

Next, set up the SOLUBILITY PRODUCT EQUILIBRIUM EXPRESSION:

In the above equation, however, we have two unknowns, [Ca²⁺] and [F^-]². So, we have to write one in terms of the other using mole ratios. According to the balanced equation, for every one mole of Ca²⁺ formed, 2 moles of F^- are formed. To simplify things a little, let's assign the the variable X for the solubility of the Ca²⁺:

If we SUBSTITUTE these values into the equilibrium expression, we now only have one variable to worry about, X:

We can now SOLVE for X:

We assigned X as the solubility of the Ca²⁺ which is equal to the solubility of the salt, CaF₂. However, our units right now are in molarity (mol/L), so we have to convert to grams:

Now, let's try to do the opposite, i.e., calculate the K_sp from the solubility of a salt.

2) The solubility of AgCl in pure water is 1.3 x 10^-5 M. Calculate the value of K_sp.

First, write the BALANCED REACTION:

Next, set up the SOLUBILITY PRODUCT EQUILIBRIUM EXPRESSION:

It is given in the problem that the solubility of AgCl is 1.3 x 10^-5. Since the mole ratio of AgCl to both Ag⁺ and Cl^- is 1:1, the solubility of each of the ions is equal to the solubility of AgCl. SUBSTITUTE the solubility given into the equilibrium expression to get K_sp:

http://pharmlabs.unc.edu/labs/solubility/structure.htm

Solute and Solvent Structure/Polarity

Solute molecules are held together by certain intermolecular forces (dipole-dipole, induced dipole-induced dipole, ion-ion, etc.), as are molecules of solvent. In order for dissolution to occur, these cohesive forces of like molecules must be broken and adhesive forces between solute and solvent must be formed.

The solubility of a drug in a given solvent is largely a function of the polarity of the solvent. Solvents may be considered polar, semi-polar or non-polar. Polar solvents will dissolve ionic and other polar solutes (i.e. those with an asymmetric charge distribution [like dissolves like]), whereas, non-polar solvents will dissolve non-polar molecules. Semi-polar solvents (eg. alcohols and ketones) may induce a certain degree of polarity in non-polar molecules and may thus act to improve the miscibility of polar and non-polar liquids. The relationship between polarity and solubility may be used in practice to alter the solubility of a drug in a pharmaceutical solution.

One approach is to alter the polarity of the solute by shifting it between its molecular (undissociated) and ionic (dissociated) states. A shift toward the ionic form improves solubility of the solute in water and other polar solvents. A shift toward the molecular species improves solute solubility in non-polar solvents. Such shifts may be produced by altering the pH of the solution (or using the salt form of the compound).

Another approach is to mix solvents of different polarities to form a solvent system of optimum polarity to dissolve the solute. Such solvents must, obviously, be miscible. This method is referred to as solvent blending or cosolvency and uses the dielectric constant as a guide to developing the cosolvent system. Since many solvents may be toxic when ingested, most solvent blends are limited to mixtures containing water, ethanol, glycerin, propylene glycol, polyethylene glycol 400 or sorbitol solution. The list is somewhat expanded for solutions for external application.

The dielectric constant (δ) of a compound is an index of its polarity. A series of solvents of increasing polarity will show a similar increase in dielectric constant.

http://aceyourchemistry.blogspot.com.au/2013/01/part-1-like-dissolves-like-to-dissolve.html

Part 1 : "Like dissolves Like" - To dissolve or not to dissolve

To dissolve or not to dissolve, that is the question. ''Like dissolves like'' is an expression used by chemists to remember how solvents work. It refers to ''polar'' and ''non-polar'' solvents and solutes. For instance, water is polar while oil is non-polar. Therefore, water will not dissolve oil. In other words, they are immiscible. Meanwhile, ionic salt such as sodium chloride is ionic (which is considered to be extremely polar). Since like dissolves like, that means polar dissolves polar, thus the ionic salt is soluble in water.

 

However, how do I know whether a molecule is polar? There are two conditions that you need to consider before you determine if the molecule is polar.

 Condition 1: Is the bond polar? As long as the bonding atoms are different, they will have different electronegativity, which leads to a dipole moment in the bond i.e. the bond is polar. For example, H-Br is a polar bond whereas Br-Br is not.

 Condition 2: What is the molecular geometry?

 Consider the molecule CO2, which is carbon dioxide. The displayed structure looks like this i.e. O=C=O.  The C=O bond is polar because the bonding atoms are different and they have different electronegativity.  However, the molecule is linear in shape. As a result, the dipoles (treated as vectors) will cancel out each other. In the linear molecules, these dipole moments have the same magnitude, but since they are pointing in opposite directions, there is cancellation of dipoles that renders the molecule non-polar.

 Now, what about water? Isn't water a non-polar molecule? The answer is NO! If the water molecule is linear too, then it will not be polar. However if you were to draw the dot and cross diagram for water, you will find that the central oxygen atom has 2 lone pairs of electrons.

 As a result, the molecular geometry is actually bent. As a result, there is a collective effect of the individual polar O-H bond that give rise to a net dipole moment, and hence water is a polar molecule.

http://www.elmhurst.edu/~chm/vchembook/174temppres.html

Temperature and Pressure Effects on Solubility

Effect of Temperature on Solubility:

The solubility of solutes is dependent on temperature. When a solid dissolves in a liquid, a change in the physical state of the solid analogous to melting takes place. Heat is required to break the bonds holding the molecules in the solid together. At the same time, heat is given off during the formation of new solute -- solvent bonds.

CASE I: Decrease in solubility with temperature:

If the heat given off in the dissolving process is greater than the heat required to break apart the solid, the net dissolving reaction is exothermic (energy given off). The addition of more heat (increases temperature) inhibits the dissolving reaction since excess heat is already being produced by the reaction. This situation is not very common where an increase in temperature produces a decrease in solubility.

CASE II: Increase in solubility with temperature:

If the heat given off in the dissolving reaction is less than the heat required to break apart the solid, the net dissolving reaction is endothermic (energy required). The addition of more heat facilitates the dissolving reaction by providing energy to break bonds in the solid. This is the most common situation where an increase in temperature produces an increase in solubility for solids.

The use of first-aid instant cold packs is an application of this solubility principle. A salt such as ammonium nitrate is dissolved in water after a sharp blow breaks the containers for each. The dissolving reaction is endothermic - requires heat. Therefore the heat is drawn from the surroundings, the pack feels cold.