A solution is an Homogeneous mixture (a mixture that has uniform properties through out the material).
Water is the most common solvent that we use.
Are solutions with water as the solvent
A solution is formed when one substances mixes uniformly throughout another. The process of solvation occurs when the attractive forces between the solute and solvent are approximately the same or greater then the solute-solute and solvent-solvent attractive forces. When this happens, the solvent particles form a cage around the solute particles.
Hydration is the process of attaching water onto a substance. In the process of solvation, the hydrogen (positive) end of water associates with negative ions and the oxygen end (negative) associates with positive cations.
In polar solvents, the solvent tends to form a solvent cage around the solute. In this process, the negative and positive ends of the polar molecules line up with each other.
Water and ethanol: hydrogen of water aligns with oxygen of ethanol, hydrogen of ethanol aligns with oxygen of water.
For ionic compounds, the positive dipole of the solvent aligns with the anions, and the negative dipole of the solvent aligns with the cation. The cations and anions totally separate in solution.
The solvation of an ionic compound in water into cations and anions can be written as:
Ca(NO3)2(aq) à Ca2+(aq) + 2NO3-(aq)
NaCl(aq) à Na+(aq) + Cl-(aq)
KF(aq) à K+(aq) + F-(aq)
The enthalpy of solvation essentially consists of three components. We have to break the intermolecular attraction between solvent molecules, we have to break the intermolecular attraction between solute particles (both of these are endothermic), then we make the intermolecular attraction between solute and solvent particles (exothermic).
The sum of these components may be endothermic or exothermic.
However, the spontaneous direction of flow in this universe is towards lower energy (i.e. exothermic processes.
This would mean that only exothermic processes could happen; however, there is another component that influences spontaneous processes: Entropy. Entropy is described as randomness or disorder. It is a measure of the number of possible arrangements of the system.
Example: Solid salt in water; the salt ions are packed into a crystal structure that does not allow rearrangement. The water molecules can move around and have a certain amount of entropy (possible arrangements), and there is a limited amount of arrangements the solid salt can take in the bottom of the water. After the salt dissolves, the salt ions can be arranged throughout the water, and the water has more arrangement possibilities with cations, anion, and water molecules to interact with, so the arrangement possibilities has increased significantly, the entropy has increased.
The natural order of flow (spontaneous) in this universe is towards greater entropy.
So spontaneous processes are a combination of increasing entropy and decreasing energy (exothermic). Therefore, endothermic reactions occur when the energy gained is offset by a corresponding increase in entropy.
Solution processes all involve an increase in entropy.
Soluble: a substance that dissolves to a significant extent in a solvent. [can be taken to be less than 0.1 g/100 g water]
Insoluble: a substance that does not dissolves to a significant extent in a solvent.
Two liquids that are infinitely soluble in each other are said to be miscible. Water and ethyl alcohol [drinking], water and vinegar, water and isopropyl alcohol [rubbing].
Two liquids that separate into separate layers (insoluble in each other) are immiscible. Water and oil, water and gasoline.
Solubility is the maximum amount of solute that can dissolve in solution at a given temperature; usually given in grams per 100 mL of solvent.[at times g/100 g]
Saturated: a solution holding the maximum amount of solute (evidenced by solute that will not dissolve)
Unsaturated: a solution holding less than the maximum amount of solute
Supersaturated: a solution holding more than the stable maximum amount of solute. This is a metastable solution. It can remain liquid until something triggers crystallization (a seed crystal), and then crystallizes rapidly.
In general, the stronger the attraction between solute and solvent particles, the greater the solubility.
A liquid made up of polar molecules is a polar solvent. Water, H2O; Ethyl alcohol, CH3CH2OH (C2H5OH); Acetic acid, CH3COOH
A liquid made up of nonpolar molecules is a nonpolar solvent. Carbon Tetrachloride, CCl4; Hexane, C6H14, Ethyl Ether, C4H10O
Most molecules containing oxygen are polar, some exceptions occur. Ethers (C-O-C) are only mildly polar and behave mostly nonpolar.
The Like Dissolves Like Rule states that a solute dissolved in a solvent because their molecules are similar in polarity. This works for liquids or solids dissolving in a liquid.
Ionic compounds tend to dissolve in polar solvents and not nonpolar solvents. Although many ionic compounds don't dissolve readily in polar solvents.
Two liquids that are infinitely soluble in each other are said to be miscible. Water and ethyl alcohol [drinking], water and vinegar, water and isopropyl alcohol [rubbing].
Two liquids that separate into separate layers are immiscible. Water and oil, water and gasoline
Item of Interest: Methanol, miscible with both water and gasoline
Like dissolves like
Ionic compounds are similar to polar compounds.
Nonpolar solute soluble in nonpolar solvent: grease, turpentine, oil, gasoline, ethyl ether
Polar and ionic compounds soluble in polar solvents: Water, table sugar (C12H22O11), table salt (NaCl)
A solution is composed of a solute dissolved in a solvent. Normally, the solute is the lesser quantity and the solvent is the greater quantity.
Solubility is the maximum amount of solute that can dissolve in solution at a given temperature; usually given in grams per 100 mL of solvent.
The solubility of most solids and liquids increases with increasing temperature. This increase in solubility can be linear or exponential.
The solubility of a gas in liquid decreases rapidly with increasing temperature.
The solubility of a gas in liquid increases rapidly with increasing partial pressure of the gas over the liquid.
Henry's Law states that the solubility of a gas in liquid is proportional to the partial pressure of the gas over the liquid.
Sg = kPg
Where, Sg is the solubility of the gas, usually in molarity; Pg is the partial pressure of the gas over the solution and k is the proportionality constant known as Henry's Law Constant.
One of the ways of telling how much solute is dissolved in a solvent is by using the mass (weight/weight) percent concentration; %(w/w).
%(w/w) = (mass solute/mass solution)(100%)
or mass of solute in 100 g of solution.
As a Ratio:
mass solute / mass solution = %(w/w) / 100%
Also: mass solute per sum of mass solute and solvent (x 100%)
10 grams of NaCl in 90 grams water is (10/(10 + 90)) x 100% = 10%
10 grams of NaCl in 100 grams water is (10/(10 + 100)) x 100% = 9.1%
The concept of Mass percent concentration can be used to relate mass solute with mass solution, mass solute with mass solvent, and mass solvent with mass solution.
The decimal equivalent of the mass percent concentration [ mass solute/mass solution] relates the mass of the solute with the mass of the solution.
Example: 25% mass concentration gives a unit factor of 0.25 (0.25/1 or 25/100) mass solute/mass solution
This can be modified to relate mass solute with mass
solvent
[mass solute/mass solvent = mass solute/(mass solution - mass solute)]
Example: 25% mass concentration gives a solute/solvent unit factor of (0.25/(1-.025) = (0.25/0.75) = 0.33) mass solute/mass solvent
Another modification can relate mass solvent to mass
solution
[mass solvent/ mass solution = (mass solution - mass solute)/mass solution]
Example: 25% mass concentration gives a solvent/solution unit factor of ((1-.025)/1 = (0.75/1) = 0.75) mass solute/mass solvent
A Saline Nasal Spray contains 0.65% sodium chloride.
What mass of sodium chloride is in 45 g of solution? [solution to solute]
0.65/100 g NaCl/g solution x 45 g solution = 0.29 g NaCl
Table vinegar contains 5% acetic acid. How much table vinegar contains 15 g of acetic acid? [solute to solution]
15 g acetic acid / (5 g acetic acid/100 g vinegar) = 15 g acetic acid x (100 g vinegar/5 g acetic acid) = 300 g vinegar
How much water is required to make 300 g of table vinegar? [solution to solvent]
300 g vinegar x ((100 - 5) g water/ 100 g vinegar) = 285 g water
How much acetic acid needs to be added to 500 g water to make 5% vinegar solution? [solvent to solute]
500 g water x (5 g acetic acid)/((100 - 5) g water) = (500 x 5)/95 = 26.3 g acetic acid
A common concentration unit for dilute solutions is parts per million (ppm). This can be described two ways.
ppm solute = (mass solute/mass solution) x 106
or
ppm solute = mg solute/kg solution
For even more dilute solutions parts per billion (ppb) can be used.
ppb solute = (mass solute/mass solution) x 109
or
ppm solute = mg solute/kg solution
The ratio of the number of moles of one component of a mixture to the total moles of all components of the mixture.
c1 = n1/nTotal
An expression of concentration
Molar Concentration, also Molarity, symbol (M), is defined as
Molarity = moles of solute/ liters of solution
M = n/V
Molarity is often more useful since chemicals react on a mole stoichiometry basis and not a weight basis.
The unit factor from molarity is moles/liter or it can also be written moles/1,000 mL
The concept of molarity allows you to associate moles of solute with volume of solution. Using molar mass also allows a connection of mass of solute with volume of solution.
How many grams of NaCl is required to make 1.000 L of 1 M solution?
1.000 L x 1 mole/L x 58.44 g/mole = 58.44 grams NaCl
149.1 grams of KCl was added to enough water to make 1.000 L of solution. What is the Molarity?
(149.1 g KCl/74.55 g/mol)/ 1.000L = 2.000 moles/L = 2.000 M KCl
What volume of 8.00 M KCl solution is required to make 500 mL of 2.00 M KCl solution?
(2.00 M x 0.500 L) = (8.00 M x ? L)
(2.00 M x 0.500 L)/8.00 M = 0.125 L = 125 mL
Molality (m) = mol solute
kg solvent
Molality is used in calculations of colligative properties including freezing point depression and boiling point elevation
When an ionic compound dissolves, the Molarity represents the concentration of the compound in solution. It does not necessarily represent the concentration of ions in solution. the concentration of ions will depend on the number of ions produced by the compound.
Since concentration (molarity) times volume gives moles [MV = n], this can be used to calculate dilution volumes using
M1V1 = M2V2
Colligative Properties depend upon only the concentration of solute particles instead of their identity.
Psoln = XsolventPosolvent
Xsolvent is the mole fraction of the solvent
Posolvent = vapor pressure of pure solvent
PTotal = PA + PB = XAPoA + XBPoB
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Interaction between solute and solvent |
∆Hsoln |
∆T soln formation |
Deviation from Raoult's Law |
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solute/solvent same as solute/solute and solvent/solvent |
Zero |
Zero |
None (Ideal Solution) |
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solute/solvent greater than solute/solute and solvent/solvent |
Negative (exothermic) |
Positive |
Negative |
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solute/solvent less than solute/solute and solvent/solvent |
Positive (endothermic) |
Negative |
Positive |
∆T = Kbmsolute
Kb = molal boiling-point elevation constant
∆T = Kfmsolute
Kf = molal freezing-point depression constant
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Solvent |
BP (°C) |
Kb |
MP (°C) |
Kf |
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Water (H2O) |
100.0 |
0.51 |
0 |
1.86 |
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Benzene (C6H6) |
80.1 |
2.53 |
5.5 |
5.12 |
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The flow of solvent but not solute through a semipermeable membrane separating two solutions of different concentrations is called osmosis.
The pressure which when applied to one side that stops osmosis is called the osmotic pressure.
π = MRT
where
π is osmotic pressure in atm
M is molarity
R is gas constant
T is absolute temperature
Solutions that have identical osmotic pressures are isotonic solutions.
Reverse osmosis occurs when a pressure, greater than the osmotic pressure is applied to the more concentrated side of the semipermeable membrane. Under this condition the solvent flows from the concentrated side to the less concentrated side.
i = (moles of particles in solution)/(moles of solute in solution)
The expected value is the number of ions per formula unit.
Ion pairing can reduce the van't Hoff Factor from the expected value.
The van't Hoff Factor can be included in the appropriate colligative property equation.
DT = Kbimsolute
DT = Kfimsolute
p = iMRT
The colligative properties provide a useful way of experimentally determining molar mass. A measured mass of a solute can be used in affecting a colligative property. The measured change in the property is then used to calculate moles of the solute. From the mass and moles of the solute the Molar mass can be calculated.
If the sizes of the solute particles are considerable, from 1 to 1000 nm, we no longer have a true solution, we have a colloid [also called a colloidal dispersion or colloidal system).
Suspensions are mixtures with the solute particle size being too large for the solution to be stable. This system is unstable and separates into layers or phases.
The basic characteristics of colloids [derived from the size of the particles]:
· They scatter light and appear turbid, cloudy, or milky. The Tyndall effect is the scattering of a beam of light, allowing the beam to be visible without the particles being visible.
· They are stable and do not separate out upon standing/time.
Colloids are stable because:
· The colloid particles are small enough to be moved by the random, chaotic Brownian motion of the solution.
· The colloid particle are solvated with a water layer that prevents the individual particles from colliding and forming a larger particle.
· The colloid particle tends to acquire a charge [positive or negative depending on the material involved]. This charge tends to have the particles repel each other.
Most natural colloids are hydrophilic (water loving). Hydrophobic (water fearing) colloids have to be stabilized to occur in water. Hydrophobic colloids can be stabilized by adsorbing ions onto the surface of the colloid, or they can be stabilized with emulsifying agents (detergents) which have both a hydrophilic and hydrophobic ends. The hydrophobic ends can bond with the hydrophobic particle leaving the hydrophilic end to bond with the water.
Colloids cannot be filtered so the colloid particles must be made to stick together to form particles large enough to filter (a process called coagulation). Different processes can be used to cause the colloids to coagulate depending on type of colloid.
Henry's Law
Sg = kPg
%(w/w) = (mass solute/mass solution)(100%)
or mass of solute in 100 g of solution.
ppm solute = mg solute/kg solution
ppm solute = mg solute/kg solution
c1 = n1/nTotal
M = n/V
Molality (m) = mol solute
kg solvent
M1V1 = M2V2
Psoln = XsolventPosolvent
Xsolvent is the mole fraction of the solvent
Posolvent = vapor pressure of pure solvent
PTotal = PA + PB = XAPoA + XBPoB
∆T = Kbmsolute
Kb = molal boiling-point elevation constant
∆T = Kfmsolute
Kf = molal freezing-point depression constant
π = MRT
where
π is osmotic pressure in atm
M is molarity
R is gas constant
T is absolute temperature
DT = Kbimsolute
DT = Kfimsolute
p = iMRT