Aqueous Reactions
Aqueous Reactions Properties of Solutes in Aqueous Solutions
Solvent versus Solute
Water has the ability to dissolve many different types of substances, resulting in a homogeneous mixture. In homogeneous mixtures involving water, water is considered to be the solvent: The typical molar concentrations of substances dissolved in water would be on the order of 10 -6 to 101 molar, thus, they are present at far lower molar concentration and are considered to be the solute.
How does water "dissolve" a solute? The polar nature of the water molecule
The Lewis Structure of water:
The central Oxygen has a tetrahedral geometry for the valence electron pairs Thus the H2O molecule will adopt a bent molecular geometry:
The Oxygen (3.5) is more electronegative than the Hydrogen (2.1), thus the O-H bond is polar covalent The bent geometry results in an overall dipole for the water molecule:
Thus, H2O can participate in the following types of non-covalent interactions: o London Dispersion Forces o Dipole-dipole interactions o Ion-dipole interactions (with ions)
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Hydrogen bonds (between other water molecules or with an appropriate solute)
It is the ability of water to participate in these diverse non-covalent interactions that allows water to "dissolve" a variety of solutes Ionic Compounds in Water Water can participate in ion-dipole interactions.
Water molecules will organize around an ion to orient the appropriate opposite partial charge of the water dipole:
Water molecules will separate, surround and disperse ions in an ionic solid:
Although H2O is a poor conductor of electricity, dissolved ions in an aqueous solution can conduct electricity. Thus, ionic aqueous solutions are known as electrolytes.
An electrolyte solution (a solution of ions) can be described by either the concentration of the ionic compound that was dissolved, or by the relative concentrations of the anion and cation components Molecular compounds in an aqueous solution
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Generally speaking, interaction with H2O will not break any covalent bonds. Thus, the ability of water to dissolve a molecular compound is based on non-covalent interactions of H2O with the molecular compound
Example: Methanol (CH3OH) mixed with water
The alcohol (-OH) group of methanol is similar to water in that the oxygen valence electron geometry is tetrahedral with two non-bonding pairs of electrons There is a polar bond between the hydrogen and oxygen of the alcohol group in methanol, similar to that in water:
Thus, water can form Hydrogen bonds with the alcohol group of methanol, and in this way the water molecules can separate, surround and disperse the molecules of methanol
Some molecular compounds interact so strongly with H2O that covalent bonds of the compound may be broken o Although the molecular compound in question may be neutral, the molecular fragments produced by the bond breakage may be oppositely charged ions. An example of this would be the uncharged molecular compounds H-Cl. Interaction with water is so strong that it results in the breakage of the H-Cl bond. This produces an H+ ion, and a Cl- ion. o Thus, although these types of molecular compounds may be neutral, in water they result in the production of ions and therefore, are electrolytes o Acids are one example of neutral compounds that ionize in aqueous solution (i.e. interact so strongly that a covalent bond is broken and ions are produced)
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Strong and weak electrolytes
Some compounds in aqueous solution dissociate completely into ions. This would include most ionic compounds, and some molecular compounds (like H-Cl) Other compounds only have a slight tendency to ionize in aqueous solutions. In other words, only a few of the molecules in solution will ionize, and most will remain as neutral compounds (although the compound is completely dissolved in water) Compounds that ionize completely are known as strong electrolytes Compounds that ionize only partially are known as weak electrolytes
For example, acetic acid only partially ionizes (i.e. is a weak electrolyte) when dissolved in H 2O. This ionization involves the breaking of a covalent bond between an oxygen and hydrogen atom in the acid:
The double arrow means that the reaction is significant in both directions o At any given moment, some of the anion form of acetic acid is combining with H+ cation to form a covalent bond and produce the neutral acetic acid o Likewise, at any given moment, some of the neutral acetic acid in aqueous solution is dissociating (i.e. ionizing) to form the anion and H+ cation o The balance (i.e. the chemical equilibrium) between these two processes determines the relative concentration of the neutral molecule and ion forms
Chemists use a double arrow to indicate the ionization of weak electrolytes, and a single arrow to indicate the ionization of strong electrolytes (i.e. in strong electrolytes, the ions have essentially no tendency to recombine to form the neutral compound)
Acids, Bases and Salts Acids Acids are substances that are able to ionize in aqueous solutions to form H+ ions (and an associated anion)
A Hydrogen atom consists of a single proton and a single electron (no neutron) Thus, an H+ ion is just a proton Acids are often referred to as "proton donors"
Different types of acids can ionize to release one or more protons
A monoprotic acid releases a single proton when it ionizes (e.g. hydrochloric acid):
A diprotic acid can release two protons when it ionizes (e.g. sulfuric acid):
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o o o
For sulfuric acid, the first ionization is complete (as indicated by the single arrow), therefore it is a strong electrolyte However, only some of the molecules of HSO4- undergo a second ionization. Thus, the deprotonization of HSO4-, and the protonization of SO42- are significant reactions (as indicated by the double arrows) Aqueous solutions of sulfuric acid will therefore contain some mixture of protons (H+), HSO4- and SO42-ions
Bases Bases are substances that react with (or accept) H + ions
A common example of a base is the hydroxide ion (OH-):
(Note that this represents the formation of a chemical bond between the ions to produce a water molecule)
Any substance that increases the concentration of OH-(aq) is a base
Some of the most common bases are metal hydroxides (e.g. NaOH, KOH, Ca(OH)2) o These are ionic compounds, that when dissolved in H2O release the metal ion and one or more hydroxide (OH-) ions
Other compounds can react with H2O in such a way that they are considered bases (even though they do not directly contribute a OH- ion)
Ammonia (NH3) reacts chemically with a H2O. It has a high affinity for the proton that can be obtained from a water molecule. This leaves a hydroxide ion (OH-) left over in solution:
Ammonia has "accepted a proton" from the water molecule, and the concentration of OH- ions has increased in solution, thus, ammonia is a base Note that the double arrows indicate that only some of the ammonia molecules will react with water to accept a proton. Thus, ammonia is a weak electrolyte.
Strong and Weak Acids and Bases
Acids and bases that ionize completely in solution are strong electrolytes, and are therefore known as strong acids or strong bases Likewise, acids and bases that only partially ionize in solution are weak electrolytes, and are known as weak acids or weak bases o If a chemical reaction depends upon the concentration of H + ions, then strong acids will be more chemically reactive than weak acids o However, the anion that is also released in the ionization of an acid can also be reactive. If a weak acid (e.g. HF) releases an anion that is highly reactive (e.g. F-), then the weak acid can also be chemically highly reactive
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Observations about acids and bases: 1. 2. 3. 4. 5.
The commonly used acids, hydrochloric, nitric and sulfuric, are all strong acids Several of the strong acids are a combination of hydrogen and a halogen (the exception is HF, which is a weak acid) There are not a lot of strong acids, most acids are weak acids There are not a lot of strong bases; the strong bases are metal hydroxides (group 1A and heavy group 2A metals) Ammonia (NH3) is a weak base
Based on the above discussion, here is a flow chart to help you decide if a compound is a strong or weak electrolyte (or a nonelectrolyte):
Neutralization reactions and salts When an acid solution and a base solution are mixed, a neutralization reaction occurs.
In general, a neutralization reaction between an acid and a metal hydroxide (strong base) produces water and a salt The term salt has come to mean any ionic compound whose cation comes from a base and whose anion comes from an acid
Example: aqueous hydrochloric acid mixed with aqueous sodium hydroxide
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Ionic Equations For strong electrolytes, where the molecules in a reaction completely dissociate to ionic forms in solution, there are two different ways we could think of writing the chemical reaction
In an example of the neutralization reaction of hydrochloric acid and the metal hydroxide of sodium (i.e. sodium hydroxide) the molecular equation for the reaction would be:
Since HCl and NaOH are strong electrolytes and will dissociate completely to their ionic forms, we can also write the reaction in terms of a complete ionic equation:
As with any equation, we can algebraically manipulate the complete ionic equation to cancel the same terms that are present as both reactants and products:
which yields the overall reaction of:
This is known as the net ionic equation for this particular reaction of strong electrolytes. It has the following characteristics o The net ionic equation only shows the ions and molecules directly involved in reaction. The other ions (Na+ and Cl- in this case) are called spectator ions. o Like any reaction, the overall charge on the reactant and products must be equal. In the above case the net charge on the reactants (H+ and OH-) is zero, and the net charge on the product (H2O) is zero.
This equation represents the net reaction of any strong acid with any strong base (i.e. a proton and hydroxide ion react to produce H2O)
Only reactions involving soluble, strong electrolytes can be written in ionic form
Metathesis Reactions In many aqueous reactions it seems that the reaction involves the ionic compounds swapping their ionic partners. For example, in the reaction involving the ionic compounds silver nitrate and potassium chloride we have:
The silver cation exchanges its nitrate anion partner for the chloride anion. Likewise, the potassium cation exchanges chloride anion for the nitrate anion:
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This swapping of ions in aqueous reactions can be symbolically represented as follows:
This type of reaction is known as a Metathesis reaction Note: metathesis is not pronounced "meta-thesis", but rather "meh-TATH-eh-sis" (apparently the Greeks prefer to pronounce it that way) There is something subtle in the above example that is important to note.
Take a look at the state of the silver chloride ion pair. The AgCl(s) is a solid, and therefore the ions are not dispersed in solution. In fact, they precipitate out of solution. The precipitation effectively removes the AgCl ions from the solution, and this is the driving force for the observed metathesis reaction
The driving force for metathesis reactions is the removal of ions from solution What are the ways in which ions can be removed from solution and thus drive a metathesis reaction? 1. Certain ions can associate to form an insoluble precipitate (as with the formation of AgCl(s)) 2. Certain ions can chemically combine to form a neutral molecular compound (resulting in either a non-electrolyte, or a weak electrolyte). o o
Acid/base neutralization reactions that produce water from H + and OH- ions are an example of this. Being a non-electrolyte (or weak electrolyte) the formation of the molecular compound from the constituent ions is essentially an irreversible process
3. Certain ions can chemically combine to form a gas, and the gas physically escapes from the solution Precipitation Reactions Metathesis reactions that result in an insoluble precipitate are called precipitation reactions
Solubility refers to the amount of a substance that can be dissolved in a given quantity of water The solubility of an ionic compound determines whether it will precipitate or not Any substance with a solubility less than 0.01 moles/L will be considered as essentially insoluble
The reaction of KI and Pb(NO3)2:
The PbI2 is an insoluble ionic compound that will precipitate and drive the metathesis reaction:
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Introduction to Oxidation-Reduction Reactions Many metals, when placed in an acid solution, will react chemically to produce hydrogen gas. For example, zinc metal in an aqueous solution of hydrochloric acid will react as follows:
What is going on in this type of reaction?
The zinc starts out as a solid metal, i.e. the neutral (uncharged) elemental form of zinc. The zinc ends up as an ion (Zn2+)in an ionic compound involving chloride ion The hydrogen starts out as a hydrogen ion (H +) from the acid, and ends up in the neutral (covalently bonded) form of diatomic hydrogen Chloride ion (Cl-) is a spectator ion in this reaction
In order for these ionic changes to occur, electrons leave the metal (producing zinc cations) and join the hydrogen ions (producing neutral diatomic hydrogen). When an atom has become more positively charged (by LOSING electrons) chemists say that it has been OXIDIZED
In the above reaction, the zinc metal has become oxidized (producing Zn 2+ cations)
When an atom becomes more negatively charged (by GAINING electrons) chemists say that is has been REDUCED
In the above reaction, the hydrogen ions have been reduced (to produce neutral diatomic hydrogen gas)
Oxidation of the metal refers to the fact that this type of reaction was actually first characterized by studying the reaction of metals with oxygen:
Many metals react with oxygen in the air to form metal oxides
In this reaction the neutral metal loses electrons to the oxygen, forming a cation (Ca2+ in this case). The metal cation participates in forming an ionic solid with the oxide (O2-) anion. Thus, the metal becomes oxidized
When one substance in a reaction loses electrons, another substance in the reaction must gain them. The oxidation of a reactant is always accompanied by the reduction of another reactant in the reaction
These types of reactions are therefore called oxidation/reduction (or REDOX) reactions
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In predicting oxidation reactions between different metals, compare the relative ease with which the two metals can be oxidized
Of the two metals, the one that is relatively more difficult to oxidize will preferentially exist in the neutral elemental form (i.e. be reduced) The other metal will preferentially exist in the oxidized (cation) form (i.e. will give its electrons up to the other metal)
A ranking of the metals by the relative ease with which they can be oxidized, is known as an Activity Series. Hydrogen is also included in such lists so as to include the behavior of acids along with the metals:
At the top (the easiest to oxidize) are the alkali metals (group 1A) and the alkaline earth metals (group 2A) At the bottom (the least likely to oxidize) are the transition metals, and in particular, the metals that are typically found in jewelry or coins are at the very bottom
The Activity Series can be use to predict the outcome of reactions between metals and either metal salts or acids
Any metal in the list can be oxidized by the ions of any metal below it (or hydrogen ions if they fall below the metal in question). The corresponding ions of the other metal (or hydrogen) will become reduced to the elemental form
For example, silver is below copper in the Activity Series, therefore, copper metal will be oxidized by silver cations (resulting in the formation of copper cations and elemental silver)
The Solution Process
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Important characteristics of solutions:
They are homogenous mixtures Solutions may be gasses, liquids or solids Each substance in a solution is a component of the solution. Usually, the component with the highest concentration is termed the solvent (other components are termed solutes)
Most solutions we will deal with are those in a liquid state, where the solvent is H2O (i.e. aqueous solutions)
The liquid state, and the solid state, are known as condensed states In condensed states, the attractive forces between molecules are strong enough (in comparison to the temperature-induce kinetic energy) to hold neighboring molecules together. o In solids, the neighbors are held rigid o In liquids, the neighbor molecules can slide past each other
Homogenous mixtures (solutions) can form only when the following attractive forces are approximately equivalent:
Attraction between solvent and solute molecules Attraction of solute molecules for other solute molecules Attraction of solvent molecules for other solvent molecules
If the attractive forces of solute molecules for other solute molecules are greater than the attractive forces of solute molecules for water, then the solute will not dissolve
For ionic solids, the lattice energy describes the attractive forces between the solute molecules (i.e. ions) For an ionic solid to dissolve in water, the water-solute attractive forces has to be strong enough to overcome the lattice energy
The process known as solvation is where the solute-solvent interactions are strong enough separate, surround and disperse a solute
If the solvent is H2O, then solvation is referred to as hydration
Energy Changes and Solution Formation We have previously studied enthalpy changes associated with chemical reactions (e.g. combustion reactions, ΔHrxn) and with physical processes (e.g. changes of state, ΔHfusion, and the heating of matter in a specific state, Molar Heat Capacity)
The process of solvation involves energy changes also, known as the enthalpy of solvation (ΔHsolv) It is a physical process, not chemical
What are the different processes that contribute to the enthalpy of solvation ΔHsolv?
There is the energy (ΔH1) associated with dispersing the solutes. It is the lattice energy (here's an example for an ionic solid):
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In order to accommodate the dispersed solutes within the H 2O solution, the H2O molecules have to separate from one another to provide the necessary space (in other words, we have to disrupt solvent-solvent interactions). There is an energy (ΔH2) associated with this process:
And, finally, there is the energy (ΔH3) associated with the formation of solvent-solute interactions for the solvated solute molecules:
What is the overall energy associated with these three distinct energetic steps involved in solvation of a solute?
ΔH1, separating the solute molecules from each other, requires an input of energy to overcome the attractive forces holding the solute molecules together. Thus, ΔH1 will be positive in sign (endothermic). ΔH2, disrupting solvent-solvent interactions, to allow space for dispersed solute molecules, will also require the input of energy. Thus, ΔH2 will be positive in sign (endothermic). ΔH3, formation of solvent-solute interactions, will release energy. Thus, ΔH3 will be negative in sign (exothermic).
If the energy released by the formation of solvent-solute interactions (ΔH3) is greater than the sum of the energies required to disrupt solute-solute interactions and solvent-solvent interactions (ΔH1 + ΔH2), then the overall enthalpy of solvation (ΔHsolv) will be exothermic and energetically favorable and spontaneous:
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Some solvation enthalpies are actually positive (i.e. endothermic) and absorb heat from their surroundings (this process is used in cold packs for muscle sprains):
Although some heats of solvation are positive (e.g. the hydration of ammonium nitrate), their reactions proceed spontaneously.
Up to this point, when trying to predict whether a reaction or process is spontaneous we have considered the overall change in heat energy (enthalpy)
Processes in which the overall heat energy of the system decreases tend to be spontaneous (i.e. a negative value for the overall ΔH indicates spontaneity)
However, the hydration of Ammonium Nitrate is spontaneous, but has a slightly positive value for the overall enthalpy (it has absorbed heat energy). Clearly, something else is going on that forces the hydration of Ammonium Nitrate to occur
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The answer to this riddle lies in the behavior of collections of objects. Consider a field with a herd of sheep, some are black and some are white.
If we let the sheep roam around and then take an aerial picture of the field, the black and white sheep would be randomly distributed:
Suppose the farmer wants to separate the sheep. A sheepdog will have to do work to separate them:
Not only that, but the dog will have to keep working to keep them separated, because their natural tendency would be to randomly distribute throughout the field We can postulate that if energy must be expended to keep the sheep from being randomly distributed, then if we reverse the process (i.e. start with sheep in an ordered arrangement and let them become randomly distributed) then energy is somehow released
The same situation is true with collections of molecules. In the case of the solvation of ammonium nitrate we start with a crystal of the ionic solid being placed in a container with water:
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When the crystal initially dissolves all the ions, though hydrated, are not randomly distributed throughout the water in the container - they are concentrated in the vicinity of the original crystal The solvated ions are therefore initially not in a random distribution throughout the container. They are initially in a more organized or ordered state (i.e. concentrated in one location in the container) and will naturally want to become more disordered or randomly distributed Processes in which the disorder of the system increases tend to occur spontaneously
Thus the increase in disorder is a driving force that can overcome the slight positive enthalpy associated with the hydration of Ammonium Nitrate
The term for the degree of disorder in a system is ENTROPY
A system with high entropy is disordered A system with low entropy is ordered Systems tend to go from low to high entropy The entropy of the universe is increasing (just check out your kitchen or bathroom for evidence of this)
Ways of Expressing Concentration Dilute versus Concentrated Solutions:
Some liquid cleaning solutions are sold in "concentrated" form, and the instructions require you to "dilute" them prior to use A concentrated solution is one where the solute is present in high concentration A dilute solution is one where the solute is present in low concentration
"high" and "low" are relative terms and do not give us any precise information regarding the actual concentration of solute in a solution.
We need a formal description of the concentration of a solute in a solution
Mass Percentage
A formal description of concentration that refers to the ratio of the mass of the component of interest (e.g. solute) with the entire mass of the sample:
For example, a 100g aqueous solution containing NaCl is evaporated to dryness, leaving only the NaCl. The NaCl is weighed and contains 5g. What was the mass percentage of NaCl in the original solution?
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For very dilute solutions (again, a relative term) the concentrations are often expressed as "parts per million" (abbreviated as ppm) ppm of component is given as:
A mass percentage of 1% would be equal to 10,000 ppm 1 ppm would be equal to .0001%
For really, really dilute solutions, the concentrations of solute are sometimes expressed as parts per billion (abbreviated as ppb) ppb of component is given as:
Mole Fraction, Molarity and Molality
It is quite common to refer to concentration of a solute in terms of the number of moles of solute
Mole Fraction:
The symbol X is used to refer to the mole fraction of a component in a sample A subscript is sometimes added to indicate the component being referred to. For example, the mole fraction of lead ion in a sample might be indicated by
The sum of the moles fractions of all components in a solution (including the solution itself) must equal 1.0 Molarity:
The symbol M is used to refer to the molarity of a solute in a solution Molarity defines the concentration of a solute in terms of the number of moles of solute and the total volume of solution
Molality:
The symbol m is used to refer to the molality of a solute in solution
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Molality defines the concentration of a solute in terms of the number of moles of solute and the mass of the solvent component
Saturated Solutions and Solubility
When a salt (NaCl) crystal is initially place in a sample of H 2O the solution is devoid of hydrated Na+ and Cl- ions:
As the water molecules surround, separate and disperse the Na+ and Cl- ions, the solution becomes populated by the hydrated ions:
The dispersed ions in the solution will collide with water molecules, the surfaces of the container and potentially other ions as well. If the original crystal has not completely dissolved, then dispersed ions can also collide with the remaining crystal. These collisions can result in the incorporation of the ions back into the NaCl crystal lattice:
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Thus, there are two opposing processes that can potentially occur: 1. 2.
Dissolving of the crystal, resulting in hydration of the individual Na + and Cl- ions, and Collision of dispersed ions resulting in an increase in the crystal mass (a process also known as Crystallization)
Note that when a hydrated ion collides with a crystal and is incorporated into the crystal lattice, that the waters that are hydrating the ion are released (i.e. the exact reverse of the hydration process) These two opposing processes of dissolving and crystallization can be represented as follows:
If the rate of dissolution is greater than the rate of crystallization, then the crystals of NaCl in the solvent will get smaller If the rate of crystallization is greater than the rate of dissolution, then the crystals of NaCl in the solvent will get larger If the rates of the two opposing processes are equal, then the size of the NaCl crystals will remain unchanged and the system is said to be in a dynamic equilibrium
A solution that is in dynamic equilibrium with undissolved solute is said to be saturated (i.e. no more solute will dissolve into the solvent under the current conditions)
The concentration of solute present in the solution under conditions of saturation is known as the solubility of that solute
At higher temperatures, usually more solute can be dissolved, and the solubility is higher.
Under some conditions it is possible to produce a supersaturated solution of a solute
Solute is dissolved to saturation at a high temperature The solution is carefully and slowly cooled to a lower temperature (the idea is not to induce the formation of tiny crystals that can serve to nucleate crystal growth)
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At the lower temperature, the concentration of solute is higher than the equilibrium concentration at that temperature. The introduction of a "seed" crystal will stimulate rapid crystal formation
Factors Affecting Solubility Factors that can affect solubility:
Properties of solute Properties of solvent Temperature Pressure (Gases)
Solute-Solvent interactions Gases
Gases are gases because the attractive forces are typically weak - involve primarily London dispersion forces Consequently, attractive forces (basis of solubility) between gas molecules and solvent are also primarily London dispersion forces London dispersion forces increase with increasing size and mass of molecules involved
Thus, the solubility of gas molecules typically increases with increasing mass of the gas molecules Some gases appear to be a lot more soluble than what their mass would predict Higher than normal solubilites of gases are an indication that a chemical reaction (in addition to the physical process of dissolution) is occuring Polar Solutes in Polar Solvents Polar solutes tend to dissolve readily in polar solvents
Interactions between polar solutes are typically dipole-dipole (or Hydrogen-bonds) Interactions between molecules of a polar solvent are also dipole-dipole (or Hydrogen-bonds) Thus, the energies associated with disrupting solute-solute interactions and solvent-solvent interactions are approximately equivalent Entropic forces can subsequently drive the dissolution process Polar liquids tend to dissolve readily in polar solvents
Pairs of liquids that mix in any proportion are termed miscible. Liquids that do not mix are termed immiscible.
Ethanol contains a hydroxyl (OH) functional group that is similar in structure to water. Attractive forces between ethanol molecules include Hydrogen bonds, like the attractive forces between water molecules. Ethanol is miscible in H2O Octane (gasoline) molecules contain only C-H and C-C bonds and are essentially non-polar molecules. Attractive forces between octane molecules include primarily London dispersion forces. Octane is not miscible in H2O Ethanol is an alcohol (contains an OH functional group). Octanol is also an alcohol (it contains a single OH functional group). However, it contains a string of 8 carbon groups compared to the two in ethanol. The carbon groups cannot participate in Hydrogen bonding (only dispersion forces).
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The single OH group in octanol is not enough to provide solubility for octanol, and octanol is essentially immiscible in H2O. The observation that if similar attractive forces exits between solute-solute molecules and solvent-solvent molecules results in miscible solutions (i.e. the ability of the solvent to dissolve the solute), has led to the following generalization: "Like dissolves like" (in other words, substances with similar intermolecular attractive forces tend to be soluble in one another) Pressure Effects on Gases and Solubility
Increasing the pressure (at constant T) results in more collisions of the gas molecules, per unit time, with the surface of the solvent. This results in greater solubility.
Temperature effects on Gases and Solubility
Kinetic energy increases with increasing temperature Thus, increasing the temperature reduces the solubility of gas molecules in a solvent
Temperature effects on Solid Solutes and Solubility
Thus, increasing temperature increases the solubility of solid solutes
Colligative Properties Various kinds of solutes (e.g. NaCl, ethylene glycol) added to H 2O result in a decrease in the freezing temperature, as well as an increase in the boiling temperature, of H2O.
Solutes added to H2O are a useful mechanism with which to prevent water-cooled machines from freezing in winter and boiling over in summer Salt added to the roads can prevent the freezing of water in winter
These effects upon the physical properties of water are termed "colligative properties" and are primarily dependent upon the amount of solute added, and are relatively independent of the type of solute Colligative properties depend upon the collective effect of the number of solute particles in a solution Raoult's Law The vapor pressure of the solute-containing solution (PA) is equal to the vapor pressure of the pure solvent (PA0) times the mole fraction of the solvent (XA)
An ideal solution is a solution that obeys Raoult's Law Variations in the strength of the solvent-solute and solvent-solvent interactions can result in vapor pressures that deviate from ideal solution behavior o If solvent-solute interactions are really strong (e.g. involve Hydrogen bonding) then the vapor pressure of a solution might be lower than ideal o If the solvent-solvent interactions are really strong then the vapor pressure of a solution might be higher than ideal
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Boiling Point Elevation Here is a typical vapor pressure diagram for H2O:
The addition of a solute results in a lowering of the vapor pressure. At any given temperature, the vapor pressure of a solution of H2O and a non-volatile solute will have a lower vapor pressure. What does this mean with regard to the liquid/gas phase transition? o At any given temperature, the vapor pressure of the liquid is lower o This means that bubbles that form will have a lower vapor pressure and will now collapse, where previously they remained (i.e. the sample boiled) o The lower vapor pressure means the sample will boil at this temperature, but only if the pressure is lowered:
Thus, at a constant pressure, the boiling point of an aqueous solution of a non-volatile solute will be higher than pure H2O:
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The increase in the boiling point (ΔT b) relative to that of the pure solvent is directly proportional to the number of solute particles per mole of solvent molecules o Molality (m) represents the number of moles of solute per kg of the solvent component (whereas molarity, M, relates to volume of total solution). Thus, molality gives information about the number of moles of solute per mole of solvent o Therefore, ΔTb is proportional to molality (and not molarity): ΔTb = Kb * m o
Kb is called the molal boiling-point-elevation constant, and is dependent only upon the choice of solvent
Freezing Point Depression
If an aqueous solution of ethanol is frozen, the water selectively freezes as a pure substance and the ethanol (with a lower freezing point) is squeezed out as pure ethanol. Consequently the vapor pressure diagram representing the vapor/solid equilibrium (vapor pressure of solid) for an aqueous salt solution will be identical to that of pure H2O:
Thus, at a constant pressure, the melting point for a aqueous salt solution will be lower than pure H2O:
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Thus, the addition of a solute will both raise the boiling point, and lower the melting point of an aqueous solution
The decrease in the freezing point temperature, ΔT f, is directly proportional to the molality of the solute: ΔTf = -Kf * m
The value of Kf, the molal freezing-point-depression constant is a characteristic only of the solvent o For water, the value of Kf is 1.86°C/m o Thus, a 1 molal solution of solute will decrease the freezing point of water by 1.86°C (i.e. ΔTf = -1.86°) o A 0.5 molal solution of NaCl will produce 1.0 molal of solute (remember, it dissociates into two ions)
Osmosis Some materials, like cellophane and some biological membranes, have tiny pores. These pores are large enough to allow solvent, like H2O, to freely pass across the membrane, but may prevent the passage of larger solute molecules
Such materials are termed semi-permeable membranes
If two solutions, with different solute concentrations, are separated by a semi-permeable membrane, there can be a net flow of solvent across the membrane
Like effusion with gas molecules, the rate of movement of solvent across the membrane is a function of the concentration of solvent and the kinetic energy Both solutions are at the same temperature, thus have the same kinetic energy. However, solvation of solute molecules means there are less free solvent molecules to pass through the membrane (i.e. those solvent molecules involved in hydrating solute are not free to pass through the membrane) On the side with the pure solvent, more molecules of solvent per unit time can pass through the membrane
Therefore, there will be a net flow of solvent molecules from the side with pure solvent to the side containing solute
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Here is how such an effect can be demonstrated:
At the beginning of the experiment, here is how things might look in a 'U' shaped glass tube with two solutions separated by a semi-permeable membrane:
After equilibrium is reached, here is how the apparatus and solutions might look:
There will be a net flow of solvent across the membrane, into the salt solution, until the height differential (π) results in a pressure that offsets the flow of solvent. This pressure, π, is termed the osmotic pressure The osmotic pressure behaves very much like pressure in the ideal gas equation: πV = nRT π = (n/V)RT
Where n = moles of solute, T is temperature (K), R is the gas constant and V is the volume of solution (n/V) is the number of moles of solute per volume of solution. This is also the molarity of the solution (M). Thus: π = MRT
The osmotic pressure is just a function of the molar concentration of solute
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Relative osmotic pressures
Two solutions with equal osmotic pressures are termed isotonic If a solution has a lower osmotic pressure than another, it is termed hypotonic If a solution has a higher osmotic pressure than another it is termed hypertonic
Summary of Colligative Properties 1. Raoult's Law The vapor pressure of the solute-containing solution (PA) is equal to the vapor pressure of the pure solvent (PA0) times the mole fraction of the solvent (XA) 2. Boiling Point Elevation ΔTb = Kb * m The increase in boiling point is equal to the molality of the solute times the molal boiling point elevation constant, Kb, (unique for each solvent) 3. Freezing Point Depression ΔTf = Kf * m The decrease in the freezing point temperature is equal to the molality of the solute times the molal freezing-point-depression constant, Kf, (unique for each solvent) 4. Osmotic Pressure π = MRT The osmotic pressure is equal to the molar concentration of solute times the gas constant times the temperature in Kelvin Determination of Molar Mass Using Colligative Properties Colligative properties relate either the mole fraction, molal concentration, or molar concentration of a solute to a measurable change in a physical property of a solution
Thus, we have a way to quantitate the number of moles of a solute based upon one of the four colligative properties If we know the quantity in grams of a solute added to a solution (to produce the measurable colligative effect), then we have both the number of moles of the solute and the associated mass Molar mass has the units of grams/mole; it describes the mass associated with 1 mole of the compound (i.e. solute)
Colloids The solutes in solutions that we have been considering up to this point are ions or small molecules
They form homogeneous solutions with the solvent They do not slowly settle out of solution, or sink to the bottom of the solution, over a period of time Gravitational forces are small compared to the kinetic energy of the molecules in solution
As solutes get larger, at some point they my start to settle out, or sink to the bottom of the solvent
Gravitational forces are greater in comparison to the kinetic energy of the solution In this case, we no longer have a homogeneous solution, but rather, a heterogeneous mixture
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Another property of "large" solute molecules, or high-molecular weight components in a mixture, is the interaction of such molecules with visible light.
Particles can scatter light when their physical dimensions are similar to the wavelength of the light Light scattering of this type will be manifest as the sample appearing "cloudy" Visible light has wavelengths of from ~400nm to ~750nm Particles with physical dimensions (i.e. diameter) of at least 400nm to 750nm can scatter visible light and will therefore appear cloudy A Carbon-Carbon bond has dimensions of approximately 0.15nm, therefore, a carbon-containing molecule would need to have on the order of several thousand carbons in a chain to be of a size to scatter visible light
Large solute molecules that are still small enough not to settle out Between the tiny solutes we have been considering up to this point, and solutes that are so large that they settle out of solution, are homogenous mixtures involving "big" solutes
These solutions are termed "colloidal dispersions", or just "colloids" Colloids are somewhere between a homogenous solution and a heterogenous mixture o They are small enough to where random collisions keep them dispersed throughout the solution, and the won't settle out due to the effects of gravity, but they are not really dissolved by the solvent o They can be detected by their light scattering properties (i.e. their large molecular size). The scattering of light indicates that dispersed throughout the solvent is a "solute" that is actually comprised of large chunky bits. Although too small to settle out, their presence is considered to represent a heterogenous mixture, rather than a homogenous solution Colloidal dispersions can be gasses, liquids or solids. Here are some examples:
Phase of Colloidal Dispersion (i.e. "solvent" phase)
Colloidal "Solute" phase
Official Name
Example
Gas
Liquid
Aerosol
Fog
Gas
Solid
Aerosol
Smoke
Liquid
Gas
Foam
Whipped cream
Liquid
Liquid
Emulsion
Homogenized Milk
Liquid
Solid
Sol
Paint
Solid
Gas
Solid foam
Marshmallow
Solid
Liquid
Solid emulsion
Butter
Solid
Solid
Solid Sol
Ruby glass
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Hydrophilic Colloids
Hydrophilic means "water loving" Hydrophilic colloids have solutes with structural groups exposed on their surface that are able to hydrogen bond with water (electronegative groups with or without hydrogen covalently attached) Since "like dissolves like" hydrophilic colloids form aqueous colloidal dispersions
Hydrophobic Colloids
Hydrophobic means "water fearing" Hydrophobic colloids have solutes with surface groups that cannot hydrogen bond, and typically involving groups that can only interact via dispersion forces Hydrophobic colloids cannot be prepared in water - unless some type of chemical alteration is done to the solute. o If the solute can adsorb ions onto its surface, then it may be able to interact strong enough with water. (Note: adsorb means to stick to the surface). o Another strategy is to combine the solute with another molecule that has two distinct ends to it. One end is hydrophobic in nature, the other hydrophilic. The hydrophobic end binds to the hydrophobic solute, and the hydrophilic end can interact with water (and solubilize the hydrophobic solute). Examples of this are soap molecules, and the digestive juices called bile (helps to dissolve fats in the diet in the aqueous environment of our bodies).
Removing colloidal solutes (or colloidal particles) This can be a little tough to do because you cannot centrifuge them out, and often cannot filter them out. Some times heat or addition of electrolytes can cause colloid particle to clump together - or coagulate
Heat increases the kinetic energy and rate of intermolecular collisions between colloidal particles. Sometimes they have a natural tendency to stick together. Thus, larger aggregates are built up that eventually settle out of solution or can be filtered out. If some colloidal solutes are solubilized in aqueous solution by surface ions, then the addition of counter ions can cause such particles to stick together (by eliminating the charge repulsion of likecharges) As seen with the section on osmosis, semi-permeable membranes can also separate colloids. This is the basis of kidney dialysis.
© 2000 Dr. Michael Blaber
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Solvent versus Solute
Water has the ability to dissolve many different types of substances, resulting in a homogeneous mixture. In homogeneous mixtures involving water, water is considered to be the solvent: The typical molar concentrations of substances dissolved in water would be on the order of 10 -6 to 101 molar, thus, they are present at far lower molar concentration and are considered to be the solute.
How does water "dissolve" a solute? The polar nature of the water molecule
The Lewis Structure of water:
The central Oxygen has a tetrahedral geometry for the valence electron pairs Thus the H2O molecule will adopt a bent molecular geometry:
The Oxygen (3.5) is more electronegative than the Hydrogen (2.1), thus the O-H bond is polar covalent The bent geometry results in an overall dipole for the water molecule:
Thus, H2O can participate in the following types of non-covalent interactions: o London Dispersion Forces o Dipole-dipole interactions o Ion-dipole interactions (with ions)
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o
Hydrogen bonds (between other water molecules or with an appropriate solute)
It is the ability of water to participate in these diverse non-covalent interactions that allows water to "dissolve" a variety of solutes Ionic Compounds in Water Water can participate in ion-dipole interactions.
Water molecules will organize around an ion to orient the appropriate opposite partial charge of the water dipole:
Water molecules will separate, surround and disperse ions in an ionic solid:
Although H2O is a poor conductor of electricity, dissolved ions in an aqueous solution can conduct electricity. Thus, ionic aqueous solutions are known as electrolytes.
An electrolyte solution (a solution of ions) can be described by either the concentration of the ionic compound that was dissolved, or by the relative concentrations of the anion and cation components Molecular compounds in an aqueous solution
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Generally speaking, interaction with H2O will not break any covalent bonds. Thus, the ability of water to dissolve a molecular compound is based on non-covalent interactions of H2O with the molecular compound
Example: Methanol (CH3OH) mixed with water
The alcohol (-OH) group of methanol is similar to water in that the oxygen valence electron geometry is tetrahedral with two non-bonding pairs of electrons There is a polar bond between the hydrogen and oxygen of the alcohol group in methanol, similar to that in water:
Thus, water can form Hydrogen bonds with the alcohol group of methanol, and in this way the water molecules can separate, surround and disperse the molecules of methanol
Some molecular compounds interact so strongly with H2O that covalent bonds of the compound may be broken o Although the molecular compound in question may be neutral, the molecular fragments produced by the bond breakage may be oppositely charged ions. An example of this would be the uncharged molecular compounds H-Cl. Interaction with water is so strong that it results in the breakage of the H-Cl bond. This produces an H+ ion, and a Cl- ion. o Thus, although these types of molecular compounds may be neutral, in water they result in the production of ions and therefore, are electrolytes o Acids are one example of neutral compounds that ionize in aqueous solution (i.e. interact so strongly that a covalent bond is broken and ions are produced)
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Strong and weak electrolytes
Some compounds in aqueous solution dissociate completely into ions. This would include most ionic compounds, and some molecular compounds (like H-Cl) Other compounds only have a slight tendency to ionize in aqueous solutions. In other words, only a few of the molecules in solution will ionize, and most will remain as neutral compounds (although the compound is completely dissolved in water) Compounds that ionize completely are known as strong electrolytes Compounds that ionize only partially are known as weak electrolytes
For example, acetic acid only partially ionizes (i.e. is a weak electrolyte) when dissolved in H 2O. This ionization involves the breaking of a covalent bond between an oxygen and hydrogen atom in the acid:
The double arrow means that the reaction is significant in both directions o At any given moment, some of the anion form of acetic acid is combining with H+ cation to form a covalent bond and produce the neutral acetic acid o Likewise, at any given moment, some of the neutral acetic acid in aqueous solution is dissociating (i.e. ionizing) to form the anion and H+ cation o The balance (i.e. the chemical equilibrium) between these two processes determines the relative concentration of the neutral molecule and ion forms
Chemists use a double arrow to indicate the ionization of weak electrolytes, and a single arrow to indicate the ionization of strong electrolytes (i.e. in strong electrolytes, the ions have essentially no tendency to recombine to form the neutral compound)
Acids, Bases and Salts Acids Acids are substances that are able to ionize in aqueous solutions to form H+ ions (and an associated anion)
A Hydrogen atom consists of a single proton and a single electron (no neutron) Thus, an H+ ion is just a proton Acids are often referred to as "proton donors"
Different types of acids can ionize to release one or more protons
A monoprotic acid releases a single proton when it ionizes (e.g. hydrochloric acid):
A diprotic acid can release two protons when it ionizes (e.g. sulfuric acid):
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o o o
For sulfuric acid, the first ionization is complete (as indicated by the single arrow), therefore it is a strong electrolyte However, only some of the molecules of HSO4- undergo a second ionization. Thus, the deprotonization of HSO4-, and the protonization of SO42- are significant reactions (as indicated by the double arrows) Aqueous solutions of sulfuric acid will therefore contain some mixture of protons (H+), HSO4- and SO42-ions
Bases Bases are substances that react with (or accept) H + ions
A common example of a base is the hydroxide ion (OH-):
(Note that this represents the formation of a chemical bond between the ions to produce a water molecule)
Any substance that increases the concentration of OH-(aq) is a base
Some of the most common bases are metal hydroxides (e.g. NaOH, KOH, Ca(OH)2) o These are ionic compounds, that when dissolved in H2O release the metal ion and one or more hydroxide (OH-) ions
Other compounds can react with H2O in such a way that they are considered bases (even though they do not directly contribute a OH- ion)
Ammonia (NH3) reacts chemically with a H2O. It has a high affinity for the proton that can be obtained from a water molecule. This leaves a hydroxide ion (OH-) left over in solution:
Ammonia has "accepted a proton" from the water molecule, and the concentration of OH- ions has increased in solution, thus, ammonia is a base Note that the double arrows indicate that only some of the ammonia molecules will react with water to accept a proton. Thus, ammonia is a weak electrolyte.
Strong and Weak Acids and Bases
Acids and bases that ionize completely in solution are strong electrolytes, and are therefore known as strong acids or strong bases Likewise, acids and bases that only partially ionize in solution are weak electrolytes, and are known as weak acids or weak bases o If a chemical reaction depends upon the concentration of H + ions, then strong acids will be more chemically reactive than weak acids o However, the anion that is also released in the ionization of an acid can also be reactive. If a weak acid (e.g. HF) releases an anion that is highly reactive (e.g. F-), then the weak acid can also be chemically highly reactive
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Observations about acids and bases: 1. 2. 3. 4. 5.
The commonly used acids, hydrochloric, nitric and sulfuric, are all strong acids Several of the strong acids are a combination of hydrogen and a halogen (the exception is HF, which is a weak acid) There are not a lot of strong acids, most acids are weak acids There are not a lot of strong bases; the strong bases are metal hydroxides (group 1A and heavy group 2A metals) Ammonia (NH3) is a weak base
Based on the above discussion, here is a flow chart to help you decide if a compound is a strong or weak electrolyte (or a nonelectrolyte):
Neutralization reactions and salts When an acid solution and a base solution are mixed, a neutralization reaction occurs.
In general, a neutralization reaction between an acid and a metal hydroxide (strong base) produces water and a salt The term salt has come to mean any ionic compound whose cation comes from a base and whose anion comes from an acid
Example: aqueous hydrochloric acid mixed with aqueous sodium hydroxide
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Ionic Equations For strong electrolytes, where the molecules in a reaction completely dissociate to ionic forms in solution, there are two different ways we could think of writing the chemical reaction
In an example of the neutralization reaction of hydrochloric acid and the metal hydroxide of sodium (i.e. sodium hydroxide) the molecular equation for the reaction would be:
Since HCl and NaOH are strong electrolytes and will dissociate completely to their ionic forms, we can also write the reaction in terms of a complete ionic equation:
As with any equation, we can algebraically manipulate the complete ionic equation to cancel the same terms that are present as both reactants and products:
which yields the overall reaction of:
This is known as the net ionic equation for this particular reaction of strong electrolytes. It has the following characteristics o The net ionic equation only shows the ions and molecules directly involved in reaction. The other ions (Na+ and Cl- in this case) are called spectator ions. o Like any reaction, the overall charge on the reactant and products must be equal. In the above case the net charge on the reactants (H+ and OH-) is zero, and the net charge on the product (H2O) is zero.
This equation represents the net reaction of any strong acid with any strong base (i.e. a proton and hydroxide ion react to produce H2O)
Only reactions involving soluble, strong electrolytes can be written in ionic form
Metathesis Reactions In many aqueous reactions it seems that the reaction involves the ionic compounds swapping their ionic partners. For example, in the reaction involving the ionic compounds silver nitrate and potassium chloride we have:
The silver cation exchanges its nitrate anion partner for the chloride anion. Likewise, the potassium cation exchanges chloride anion for the nitrate anion:
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This swapping of ions in aqueous reactions can be symbolically represented as follows:
This type of reaction is known as a Metathesis reaction Note: metathesis is not pronounced "meta-thesis", but rather "meh-TATH-eh-sis" (apparently the Greeks prefer to pronounce it that way) There is something subtle in the above example that is important to note.
Take a look at the state of the silver chloride ion pair. The AgCl(s) is a solid, and therefore the ions are not dispersed in solution. In fact, they precipitate out of solution. The precipitation effectively removes the AgCl ions from the solution, and this is the driving force for the observed metathesis reaction
The driving force for metathesis reactions is the removal of ions from solution What are the ways in which ions can be removed from solution and thus drive a metathesis reaction? 1. Certain ions can associate to form an insoluble precipitate (as with the formation of AgCl(s)) 2. Certain ions can chemically combine to form a neutral molecular compound (resulting in either a non-electrolyte, or a weak electrolyte). o o
Acid/base neutralization reactions that produce water from H + and OH- ions are an example of this. Being a non-electrolyte (or weak electrolyte) the formation of the molecular compound from the constituent ions is essentially an irreversible process
3. Certain ions can chemically combine to form a gas, and the gas physically escapes from the solution Precipitation Reactions Metathesis reactions that result in an insoluble precipitate are called precipitation reactions
Solubility refers to the amount of a substance that can be dissolved in a given quantity of water The solubility of an ionic compound determines whether it will precipitate or not Any substance with a solubility less than 0.01 moles/L will be considered as essentially insoluble
The reaction of KI and Pb(NO3)2:
The PbI2 is an insoluble ionic compound that will precipitate and drive the metathesis reaction:
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Introduction to Oxidation-Reduction Reactions Many metals, when placed in an acid solution, will react chemically to produce hydrogen gas. For example, zinc metal in an aqueous solution of hydrochloric acid will react as follows:
What is going on in this type of reaction?
The zinc starts out as a solid metal, i.e. the neutral (uncharged) elemental form of zinc. The zinc ends up as an ion (Zn2+)in an ionic compound involving chloride ion The hydrogen starts out as a hydrogen ion (H +) from the acid, and ends up in the neutral (covalently bonded) form of diatomic hydrogen Chloride ion (Cl-) is a spectator ion in this reaction
In order for these ionic changes to occur, electrons leave the metal (producing zinc cations) and join the hydrogen ions (producing neutral diatomic hydrogen). When an atom has become more positively charged (by LOSING electrons) chemists say that it has been OXIDIZED
In the above reaction, the zinc metal has become oxidized (producing Zn 2+ cations)
When an atom becomes more negatively charged (by GAINING electrons) chemists say that is has been REDUCED
In the above reaction, the hydrogen ions have been reduced (to produce neutral diatomic hydrogen gas)
Oxidation of the metal refers to the fact that this type of reaction was actually first characterized by studying the reaction of metals with oxygen:
Many metals react with oxygen in the air to form metal oxides
In this reaction the neutral metal loses electrons to the oxygen, forming a cation (Ca2+ in this case). The metal cation participates in forming an ionic solid with the oxide (O2-) anion. Thus, the metal becomes oxidized
When one substance in a reaction loses electrons, another substance in the reaction must gain them. The oxidation of a reactant is always accompanied by the reduction of another reactant in the reaction
These types of reactions are therefore called oxidation/reduction (or REDOX) reactions
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In predicting oxidation reactions between different metals, compare the relative ease with which the two metals can be oxidized
Of the two metals, the one that is relatively more difficult to oxidize will preferentially exist in the neutral elemental form (i.e. be reduced) The other metal will preferentially exist in the oxidized (cation) form (i.e. will give its electrons up to the other metal)
A ranking of the metals by the relative ease with which they can be oxidized, is known as an Activity Series. Hydrogen is also included in such lists so as to include the behavior of acids along with the metals:
At the top (the easiest to oxidize) are the alkali metals (group 1A) and the alkaline earth metals (group 2A) At the bottom (the least likely to oxidize) are the transition metals, and in particular, the metals that are typically found in jewelry or coins are at the very bottom
The Activity Series can be use to predict the outcome of reactions between metals and either metal salts or acids
Any metal in the list can be oxidized by the ions of any metal below it (or hydrogen ions if they fall below the metal in question). The corresponding ions of the other metal (or hydrogen) will become reduced to the elemental form
For example, silver is below copper in the Activity Series, therefore, copper metal will be oxidized by silver cations (resulting in the formation of copper cations and elemental silver)
The Solution Process
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Important characteristics of solutions:
They are homogenous mixtures Solutions may be gasses, liquids or solids Each substance in a solution is a component of the solution. Usually, the component with the highest concentration is termed the solvent (other components are termed solutes)
Most solutions we will deal with are those in a liquid state, where the solvent is H2O (i.e. aqueous solutions)
The liquid state, and the solid state, are known as condensed states In condensed states, the attractive forces between molecules are strong enough (in comparison to the temperature-induce kinetic energy) to hold neighboring molecules together. o In solids, the neighbors are held rigid o In liquids, the neighbor molecules can slide past each other
Homogenous mixtures (solutions) can form only when the following attractive forces are approximately equivalent:
Attraction between solvent and solute molecules Attraction of solute molecules for other solute molecules Attraction of solvent molecules for other solvent molecules
If the attractive forces of solute molecules for other solute molecules are greater than the attractive forces of solute molecules for water, then the solute will not dissolve
For ionic solids, the lattice energy describes the attractive forces between the solute molecules (i.e. ions) For an ionic solid to dissolve in water, the water-solute attractive forces has to be strong enough to overcome the lattice energy
The process known as solvation is where the solute-solvent interactions are strong enough separate, surround and disperse a solute
If the solvent is H2O, then solvation is referred to as hydration
Energy Changes and Solution Formation We have previously studied enthalpy changes associated with chemical reactions (e.g. combustion reactions, ΔHrxn) and with physical processes (e.g. changes of state, ΔHfusion, and the heating of matter in a specific state, Molar Heat Capacity)
The process of solvation involves energy changes also, known as the enthalpy of solvation (ΔHsolv) It is a physical process, not chemical
What are the different processes that contribute to the enthalpy of solvation ΔHsolv?
There is the energy (ΔH1) associated with dispersing the solutes. It is the lattice energy (here's an example for an ionic solid):
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In order to accommodate the dispersed solutes within the H 2O solution, the H2O molecules have to separate from one another to provide the necessary space (in other words, we have to disrupt solvent-solvent interactions). There is an energy (ΔH2) associated with this process:
And, finally, there is the energy (ΔH3) associated with the formation of solvent-solute interactions for the solvated solute molecules:
What is the overall energy associated with these three distinct energetic steps involved in solvation of a solute?
ΔH1, separating the solute molecules from each other, requires an input of energy to overcome the attractive forces holding the solute molecules together. Thus, ΔH1 will be positive in sign (endothermic). ΔH2, disrupting solvent-solvent interactions, to allow space for dispersed solute molecules, will also require the input of energy. Thus, ΔH2 will be positive in sign (endothermic). ΔH3, formation of solvent-solute interactions, will release energy. Thus, ΔH3 will be negative in sign (exothermic).
If the energy released by the formation of solvent-solute interactions (ΔH3) is greater than the sum of the energies required to disrupt solute-solute interactions and solvent-solvent interactions (ΔH1 + ΔH2), then the overall enthalpy of solvation (ΔHsolv) will be exothermic and energetically favorable and spontaneous:
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Some solvation enthalpies are actually positive (i.e. endothermic) and absorb heat from their surroundings (this process is used in cold packs for muscle sprains):
Although some heats of solvation are positive (e.g. the hydration of ammonium nitrate), their reactions proceed spontaneously.
Up to this point, when trying to predict whether a reaction or process is spontaneous we have considered the overall change in heat energy (enthalpy)
Processes in which the overall heat energy of the system decreases tend to be spontaneous (i.e. a negative value for the overall ΔH indicates spontaneity)
However, the hydration of Ammonium Nitrate is spontaneous, but has a slightly positive value for the overall enthalpy (it has absorbed heat energy). Clearly, something else is going on that forces the hydration of Ammonium Nitrate to occur
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The answer to this riddle lies in the behavior of collections of objects. Consider a field with a herd of sheep, some are black and some are white.
If we let the sheep roam around and then take an aerial picture of the field, the black and white sheep would be randomly distributed:
Suppose the farmer wants to separate the sheep. A sheepdog will have to do work to separate them:
Not only that, but the dog will have to keep working to keep them separated, because their natural tendency would be to randomly distribute throughout the field We can postulate that if energy must be expended to keep the sheep from being randomly distributed, then if we reverse the process (i.e. start with sheep in an ordered arrangement and let them become randomly distributed) then energy is somehow released
The same situation is true with collections of molecules. In the case of the solvation of ammonium nitrate we start with a crystal of the ionic solid being placed in a container with water:
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When the crystal initially dissolves all the ions, though hydrated, are not randomly distributed throughout the water in the container - they are concentrated in the vicinity of the original crystal The solvated ions are therefore initially not in a random distribution throughout the container. They are initially in a more organized or ordered state (i.e. concentrated in one location in the container) and will naturally want to become more disordered or randomly distributed Processes in which the disorder of the system increases tend to occur spontaneously
Thus the increase in disorder is a driving force that can overcome the slight positive enthalpy associated with the hydration of Ammonium Nitrate
The term for the degree of disorder in a system is ENTROPY
A system with high entropy is disordered A system with low entropy is ordered Systems tend to go from low to high entropy The entropy of the universe is increasing (just check out your kitchen or bathroom for evidence of this)
Ways of Expressing Concentration Dilute versus Concentrated Solutions:
Some liquid cleaning solutions are sold in "concentrated" form, and the instructions require you to "dilute" them prior to use A concentrated solution is one where the solute is present in high concentration A dilute solution is one where the solute is present in low concentration
"high" and "low" are relative terms and do not give us any precise information regarding the actual concentration of solute in a solution.
We need a formal description of the concentration of a solute in a solution
Mass Percentage
A formal description of concentration that refers to the ratio of the mass of the component of interest (e.g. solute) with the entire mass of the sample:
For example, a 100g aqueous solution containing NaCl is evaporated to dryness, leaving only the NaCl. The NaCl is weighed and contains 5g. What was the mass percentage of NaCl in the original solution?
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For very dilute solutions (again, a relative term) the concentrations are often expressed as "parts per million" (abbreviated as ppm) ppm of component is given as:
A mass percentage of 1% would be equal to 10,000 ppm 1 ppm would be equal to .0001%
For really, really dilute solutions, the concentrations of solute are sometimes expressed as parts per billion (abbreviated as ppb) ppb of component is given as:
Mole Fraction, Molarity and Molality
It is quite common to refer to concentration of a solute in terms of the number of moles of solute
Mole Fraction:
The symbol X is used to refer to the mole fraction of a component in a sample A subscript is sometimes added to indicate the component being referred to. For example, the mole fraction of lead ion in a sample might be indicated by
The sum of the moles fractions of all components in a solution (including the solution itself) must equal 1.0 Molarity:
The symbol M is used to refer to the molarity of a solute in a solution Molarity defines the concentration of a solute in terms of the number of moles of solute and the total volume of solution
Molality:
The symbol m is used to refer to the molality of a solute in solution
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Molality defines the concentration of a solute in terms of the number of moles of solute and the mass of the solvent component
Saturated Solutions and Solubility
When a salt (NaCl) crystal is initially place in a sample of H 2O the solution is devoid of hydrated Na+ and Cl- ions:
As the water molecules surround, separate and disperse the Na+ and Cl- ions, the solution becomes populated by the hydrated ions:
The dispersed ions in the solution will collide with water molecules, the surfaces of the container and potentially other ions as well. If the original crystal has not completely dissolved, then dispersed ions can also collide with the remaining crystal. These collisions can result in the incorporation of the ions back into the NaCl crystal lattice:
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Thus, there are two opposing processes that can potentially occur: 1. 2.
Dissolving of the crystal, resulting in hydration of the individual Na + and Cl- ions, and Collision of dispersed ions resulting in an increase in the crystal mass (a process also known as Crystallization)
Note that when a hydrated ion collides with a crystal and is incorporated into the crystal lattice, that the waters that are hydrating the ion are released (i.e. the exact reverse of the hydration process) These two opposing processes of dissolving and crystallization can be represented as follows:
If the rate of dissolution is greater than the rate of crystallization, then the crystals of NaCl in the solvent will get smaller If the rate of crystallization is greater than the rate of dissolution, then the crystals of NaCl in the solvent will get larger If the rates of the two opposing processes are equal, then the size of the NaCl crystals will remain unchanged and the system is said to be in a dynamic equilibrium
A solution that is in dynamic equilibrium with undissolved solute is said to be saturated (i.e. no more solute will dissolve into the solvent under the current conditions)
The concentration of solute present in the solution under conditions of saturation is known as the solubility of that solute
At higher temperatures, usually more solute can be dissolved, and the solubility is higher.
Under some conditions it is possible to produce a supersaturated solution of a solute
Solute is dissolved to saturation at a high temperature The solution is carefully and slowly cooled to a lower temperature (the idea is not to induce the formation of tiny crystals that can serve to nucleate crystal growth)
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At the lower temperature, the concentration of solute is higher than the equilibrium concentration at that temperature. The introduction of a "seed" crystal will stimulate rapid crystal formation
Factors Affecting Solubility Factors that can affect solubility:
Properties of solute Properties of solvent Temperature Pressure (Gases)
Solute-Solvent interactions Gases
Gases are gases because the attractive forces are typically weak - involve primarily London dispersion forces Consequently, attractive forces (basis of solubility) between gas molecules and solvent are also primarily London dispersion forces London dispersion forces increase with increasing size and mass of molecules involved
Thus, the solubility of gas molecules typically increases with increasing mass of the gas molecules Some gases appear to be a lot more soluble than what their mass would predict Higher than normal solubilites of gases are an indication that a chemical reaction (in addition to the physical process of dissolution) is occuring Polar Solutes in Polar Solvents Polar solutes tend to dissolve readily in polar solvents
Interactions between polar solutes are typically dipole-dipole (or Hydrogen-bonds) Interactions between molecules of a polar solvent are also dipole-dipole (or Hydrogen-bonds) Thus, the energies associated with disrupting solute-solute interactions and solvent-solvent interactions are approximately equivalent Entropic forces can subsequently drive the dissolution process Polar liquids tend to dissolve readily in polar solvents
Pairs of liquids that mix in any proportion are termed miscible. Liquids that do not mix are termed immiscible.
Ethanol contains a hydroxyl (OH) functional group that is similar in structure to water. Attractive forces between ethanol molecules include Hydrogen bonds, like the attractive forces between water molecules. Ethanol is miscible in H2O Octane (gasoline) molecules contain only C-H and C-C bonds and are essentially non-polar molecules. Attractive forces between octane molecules include primarily London dispersion forces. Octane is not miscible in H2O Ethanol is an alcohol (contains an OH functional group). Octanol is also an alcohol (it contains a single OH functional group). However, it contains a string of 8 carbon groups compared to the two in ethanol. The carbon groups cannot participate in Hydrogen bonding (only dispersion forces).
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The single OH group in octanol is not enough to provide solubility for octanol, and octanol is essentially immiscible in H2O. The observation that if similar attractive forces exits between solute-solute molecules and solvent-solvent molecules results in miscible solutions (i.e. the ability of the solvent to dissolve the solute), has led to the following generalization: "Like dissolves like" (in other words, substances with similar intermolecular attractive forces tend to be soluble in one another) Pressure Effects on Gases and Solubility
Increasing the pressure (at constant T) results in more collisions of the gas molecules, per unit time, with the surface of the solvent. This results in greater solubility.
Temperature effects on Gases and Solubility
Kinetic energy increases with increasing temperature Thus, increasing the temperature reduces the solubility of gas molecules in a solvent
Temperature effects on Solid Solutes and Solubility
Thus, increasing temperature increases the solubility of solid solutes
Colligative Properties Various kinds of solutes (e.g. NaCl, ethylene glycol) added to H 2O result in a decrease in the freezing temperature, as well as an increase in the boiling temperature, of H2O.
Solutes added to H2O are a useful mechanism with which to prevent water-cooled machines from freezing in winter and boiling over in summer Salt added to the roads can prevent the freezing of water in winter
These effects upon the physical properties of water are termed "colligative properties" and are primarily dependent upon the amount of solute added, and are relatively independent of the type of solute Colligative properties depend upon the collective effect of the number of solute particles in a solution Raoult's Law The vapor pressure of the solute-containing solution (PA) is equal to the vapor pressure of the pure solvent (PA0) times the mole fraction of the solvent (XA)
An ideal solution is a solution that obeys Raoult's Law Variations in the strength of the solvent-solute and solvent-solvent interactions can result in vapor pressures that deviate from ideal solution behavior o If solvent-solute interactions are really strong (e.g. involve Hydrogen bonding) then the vapor pressure of a solution might be lower than ideal o If the solvent-solvent interactions are really strong then the vapor pressure of a solution might be higher than ideal
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Boiling Point Elevation Here is a typical vapor pressure diagram for H2O:
The addition of a solute results in a lowering of the vapor pressure. At any given temperature, the vapor pressure of a solution of H2O and a non-volatile solute will have a lower vapor pressure. What does this mean with regard to the liquid/gas phase transition? o At any given temperature, the vapor pressure of the liquid is lower o This means that bubbles that form will have a lower vapor pressure and will now collapse, where previously they remained (i.e. the sample boiled) o The lower vapor pressure means the sample will boil at this temperature, but only if the pressure is lowered:
Thus, at a constant pressure, the boiling point of an aqueous solution of a non-volatile solute will be higher than pure H2O:
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The increase in the boiling point (ΔT b) relative to that of the pure solvent is directly proportional to the number of solute particles per mole of solvent molecules o Molality (m) represents the number of moles of solute per kg of the solvent component (whereas molarity, M, relates to volume of total solution). Thus, molality gives information about the number of moles of solute per mole of solvent o Therefore, ΔTb is proportional to molality (and not molarity): ΔTb = Kb * m o
Kb is called the molal boiling-point-elevation constant, and is dependent only upon the choice of solvent
Freezing Point Depression
If an aqueous solution of ethanol is frozen, the water selectively freezes as a pure substance and the ethanol (with a lower freezing point) is squeezed out as pure ethanol. Consequently the vapor pressure diagram representing the vapor/solid equilibrium (vapor pressure of solid) for an aqueous salt solution will be identical to that of pure H2O:
Thus, at a constant pressure, the melting point for a aqueous salt solution will be lower than pure H2O:
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Thus, the addition of a solute will both raise the boiling point, and lower the melting point of an aqueous solution
The decrease in the freezing point temperature, ΔT f, is directly proportional to the molality of the solute: ΔTf = -Kf * m
The value of Kf, the molal freezing-point-depression constant is a characteristic only of the solvent o For water, the value of Kf is 1.86°C/m o Thus, a 1 molal solution of solute will decrease the freezing point of water by 1.86°C (i.e. ΔTf = -1.86°) o A 0.5 molal solution of NaCl will produce 1.0 molal of solute (remember, it dissociates into two ions)
Osmosis Some materials, like cellophane and some biological membranes, have tiny pores. These pores are large enough to allow solvent, like H2O, to freely pass across the membrane, but may prevent the passage of larger solute molecules
Such materials are termed semi-permeable membranes
If two solutions, with different solute concentrations, are separated by a semi-permeable membrane, there can be a net flow of solvent across the membrane
Like effusion with gas molecules, the rate of movement of solvent across the membrane is a function of the concentration of solvent and the kinetic energy Both solutions are at the same temperature, thus have the same kinetic energy. However, solvation of solute molecules means there are less free solvent molecules to pass through the membrane (i.e. those solvent molecules involved in hydrating solute are not free to pass through the membrane) On the side with the pure solvent, more molecules of solvent per unit time can pass through the membrane
Therefore, there will be a net flow of solvent molecules from the side with pure solvent to the side containing solute
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Here is how such an effect can be demonstrated:
At the beginning of the experiment, here is how things might look in a 'U' shaped glass tube with two solutions separated by a semi-permeable membrane:
After equilibrium is reached, here is how the apparatus and solutions might look:
There will be a net flow of solvent across the membrane, into the salt solution, until the height differential (π) results in a pressure that offsets the flow of solvent. This pressure, π, is termed the osmotic pressure The osmotic pressure behaves very much like pressure in the ideal gas equation: πV = nRT π = (n/V)RT
Where n = moles of solute, T is temperature (K), R is the gas constant and V is the volume of solution (n/V) is the number of moles of solute per volume of solution. This is also the molarity of the solution (M). Thus: π = MRT
The osmotic pressure is just a function of the molar concentration of solute
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Relative osmotic pressures
Two solutions with equal osmotic pressures are termed isotonic If a solution has a lower osmotic pressure than another, it is termed hypotonic If a solution has a higher osmotic pressure than another it is termed hypertonic
Summary of Colligative Properties 1. Raoult's Law The vapor pressure of the solute-containing solution (PA) is equal to the vapor pressure of the pure solvent (PA0) times the mole fraction of the solvent (XA) 2. Boiling Point Elevation ΔTb = Kb * m The increase in boiling point is equal to the molality of the solute times the molal boiling point elevation constant, Kb, (unique for each solvent) 3. Freezing Point Depression ΔTf = Kf * m The decrease in the freezing point temperature is equal to the molality of the solute times the molal freezing-point-depression constant, Kf, (unique for each solvent) 4. Osmotic Pressure π = MRT The osmotic pressure is equal to the molar concentration of solute times the gas constant times the temperature in Kelvin Determination of Molar Mass Using Colligative Properties Colligative properties relate either the mole fraction, molal concentration, or molar concentration of a solute to a measurable change in a physical property of a solution
Thus, we have a way to quantitate the number of moles of a solute based upon one of the four colligative properties If we know the quantity in grams of a solute added to a solution (to produce the measurable colligative effect), then we have both the number of moles of the solute and the associated mass Molar mass has the units of grams/mole; it describes the mass associated with 1 mole of the compound (i.e. solute)
Colloids The solutes in solutions that we have been considering up to this point are ions or small molecules
They form homogeneous solutions with the solvent They do not slowly settle out of solution, or sink to the bottom of the solution, over a period of time Gravitational forces are small compared to the kinetic energy of the molecules in solution
As solutes get larger, at some point they my start to settle out, or sink to the bottom of the solvent
Gravitational forces are greater in comparison to the kinetic energy of the solution In this case, we no longer have a homogeneous solution, but rather, a heterogeneous mixture
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Another property of "large" solute molecules, or high-molecular weight components in a mixture, is the interaction of such molecules with visible light.
Particles can scatter light when their physical dimensions are similar to the wavelength of the light Light scattering of this type will be manifest as the sample appearing "cloudy" Visible light has wavelengths of from ~400nm to ~750nm Particles with physical dimensions (i.e. diameter) of at least 400nm to 750nm can scatter visible light and will therefore appear cloudy A Carbon-Carbon bond has dimensions of approximately 0.15nm, therefore, a carbon-containing molecule would need to have on the order of several thousand carbons in a chain to be of a size to scatter visible light
Large solute molecules that are still small enough not to settle out Between the tiny solutes we have been considering up to this point, and solutes that are so large that they settle out of solution, are homogenous mixtures involving "big" solutes
These solutions are termed "colloidal dispersions", or just "colloids" Colloids are somewhere between a homogenous solution and a heterogenous mixture o They are small enough to where random collisions keep them dispersed throughout the solution, and the won't settle out due to the effects of gravity, but they are not really dissolved by the solvent o They can be detected by their light scattering properties (i.e. their large molecular size). The scattering of light indicates that dispersed throughout the solvent is a "solute" that is actually comprised of large chunky bits. Although too small to settle out, their presence is considered to represent a heterogenous mixture, rather than a homogenous solution Colloidal dispersions can be gasses, liquids or solids. Here are some examples:
Phase of Colloidal Dispersion (i.e. "solvent" phase)
Colloidal "Solute" phase
Official Name
Example
Gas
Liquid
Aerosol
Fog
Gas
Solid
Aerosol
Smoke
Liquid
Gas
Foam
Whipped cream
Liquid
Liquid
Emulsion
Homogenized Milk
Liquid
Solid
Sol
Paint
Solid
Gas
Solid foam
Marshmallow
Solid
Liquid
Solid emulsion
Butter
Solid
Solid
Solid Sol
Ruby glass
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Hydrophilic Colloids
Hydrophilic means "water loving" Hydrophilic colloids have solutes with structural groups exposed on their surface that are able to hydrogen bond with water (electronegative groups with or without hydrogen covalently attached) Since "like dissolves like" hydrophilic colloids form aqueous colloidal dispersions
Hydrophobic Colloids
Hydrophobic means "water fearing" Hydrophobic colloids have solutes with surface groups that cannot hydrogen bond, and typically involving groups that can only interact via dispersion forces Hydrophobic colloids cannot be prepared in water - unless some type of chemical alteration is done to the solute. o If the solute can adsorb ions onto its surface, then it may be able to interact strong enough with water. (Note: adsorb means to stick to the surface). o Another strategy is to combine the solute with another molecule that has two distinct ends to it. One end is hydrophobic in nature, the other hydrophilic. The hydrophobic end binds to the hydrophobic solute, and the hydrophilic end can interact with water (and solubilize the hydrophobic solute). Examples of this are soap molecules, and the digestive juices called bile (helps to dissolve fats in the diet in the aqueous environment of our bodies).
Removing colloidal solutes (or colloidal particles) This can be a little tough to do because you cannot centrifuge them out, and often cannot filter them out. Some times heat or addition of electrolytes can cause colloid particle to clump together - or coagulate
Heat increases the kinetic energy and rate of intermolecular collisions between colloidal particles. Sometimes they have a natural tendency to stick together. Thus, larger aggregates are built up that eventually settle out of solution or can be filtered out. If some colloidal solutes are solubilized in aqueous solution by surface ions, then the addition of counter ions can cause such particles to stick together (by eliminating the charge repulsion of likecharges) As seen with the section on osmosis, semi-permeable membranes can also separate colloids. This is the basis of kidney dialysis.
© 2000 Dr. Michael Blaber
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