Aqueous Reactions and Solution Stoichiometry Solution Composition
Water possesses many unusual properties. One of the most important properties of water is its ability to dissolve a wide variety of substances. It may sound strange, but absolutely pure water can be considered corrosive due to its capacity to absorb other compounds and ions. Solutions in which water is the dissolving medium are called aqueous solutions. Limestone caves, for example, are formed by the dissolving action of water, and dissolved CO2, on solid Calcium Carbonate. The dissolved mineral is then deposited as stalagtites and stalagmites as the water evaporates: CaCO3(s) + H2O(l) + CO2(aq) Ca(HCO3)2(aq) Many physiological chemical reactions occur in aqueous solutions. How do we express solution composition? What are the chemical forms in which substances occur in aqueous solutions?
Solution Composition A solution is a homogenous mixture of two or more substances, consisting of 1. 2.
The solvent - usually the substance in greater concentration The other component(s) is (are) called the solute(s) - they are said to be dissolved in the solvent
When a small amount of NaCl is dissolved in a large quantity of water, we refer to the water as the solvent and the NaCl as the solute.
Molarity The term concentration is used to indicate the amount of solute dissolved in a given quantity of solvent or solution. The most widely used way of quantifying concentration in chemistry is molarity. The molarity (symbol M) of a solution is defined as the number of moles of solute in a liter volume of solution:
Dilution For convenience, solutions are either purchased or prepared in concentrated stock solutions which must be diluted prior to use.
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When we take a sample of a stock solution we have a certain number of moles of molecules in that sample. Dilution alters the molarity (i.e. concentration) of the solution but not the total number of moles of molecules in the solution (in other words, dilution does not create or destroy molecules).
One of the standard equations for determining the effects of dilution upon a sample is to set up an equation comparing (concentration)*(volume) before and after dilution. Since (concentration)*(volume) gives us the total number of moles in the sample, and since this does not change, this value before and after dilution are equal: (concentration)*(volume) = (concentration)*(volume) (moles/liter)*(liter) = (moles/liter)*(liter) moles = moles
Solution Stoichiometry For balanced chemical equations involving solutions we calculate the number of moles by knowing the concentration (moles/liter, or Molarity) and volume (in liters). Titrations How can we know the concentration of some solution of interest? One answer to this problem lies in the method of titration. In titration we will make use of a second solution known as a standard solution that has the following characteristics: 1. 2.
The second solution contains a chemical which reacts in a defined way, with known stoichiometry, with the solute of the first solution The concentration of the solute in this second solution is known.
Classic titrations include so-called acid-base titrations. In these experiments a solution of an acid with an unknown concentration is titrated with a solution of known concentration of base (or vice versa). For example, we may have a solution of hydrochloric acid (HCl) of unknown concentration and a standard solution of NaOH. To a fixed amount of the HCl solution is added incremental amounts of the NaOH solution until the acid is completely neutralized - i.e. a stoichiometrically equivalent quantity of HCl and NaOH have been combined. This is known as the equivalence point in the titration. By knowing the concentration of the standard solution, and the amount added to achieve stoichiometric equivalency, we can determine the amount of moles of HCl in the original sample volume.
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How do we know when we have reached the equivalence point in such a titration experiment? In this type of acid-base titration, so called indicator-dyes are used. For example phenolphthalein is colorless in acidic solutions and turns red in basic solutions. Thus, in the above experiment we will add a small amount of this indicator-dye and add base until we barely begin to see a color change to red. 25 mls of a solution of HCl with an unknown concentration is titrated with a standard solution of 0.5 M NaOH. The phenolphthalein indicator dye begins to turn color after the addition of 2.8 mls of standard solution. What is the concentration of the HCl? Balanced equation for the reaction: HCl + NaOH -> NaCl + H2O
of NaOH was added Since the stoichiometry of the NaOH and HCl is 1:1, the sample of HCl must have contained 0.0014 moles of HCl. The concentration of the HCl solution is therefore:
, or 0.056 M
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:
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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) 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:
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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.
Distinctions between concentrations of ionic compounds and resulting concentrations of ions in an aqueous solution
1.0 mole of NaCl dissolved in 1 liter of H2O results in a 1.0 molar solution of NaCl o When 1.0 mole of NaCl is dissolved it produces 1.0 mole of Na+ ion and 1.0 mole of Clion
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
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
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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)
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:
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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):
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-):
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(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
Some strong acids and bases: Strong Acids: HClO3
Chloric Acid
HBr
Hydrobromic Acid
HCl
Hydrochloric Acid
HI
Hydroiodic Acid
HNO3
Nitric Acid
HClO4
Perchloric Acid
H2SO4
Sulfuric Acid
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Strong Bases: LiOH, NaOH, KOH, RbOH, CsOH
Group 1A Metal Hydroxides
Ca(OH)2, Sr(OH)2, Ba(OH)2
The heavy Group 2A Metal Hydroxides
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
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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
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.
Let's look at another neutralization reaction between a strong acid (Nitric acid) and a strong base (KOH):
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The molecular equation would be:
The complete ionic equation would be:
The net ionic equation would therefore be:
This is the same net ionic equation for the reaction of HCl with NaOH
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:
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.
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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 solubility of NaCl is around 10 moles/L (it is highly soluble and won't drive a metathesis reaction). The solubility of PbI2 is around 0.0012 moles/L (it is essentially insoluble, and can drive a metathesis reaction).
The reaction of KI and Pb(NO3)2:
The PbI2 is an insoluble ionic compound that will precipitate and drive the metathesis reaction:
Can we predict whether an ionic compound will be soluble or not?
If an ionic compound is insoluble it means that neighboring ions have an attraction for each other that is greater than the attraction of water for the ions (i.e. water molecules cannot separate, surround and disperse the ions in the ionic solid) Unfortunately, there are no clear rules for solubility based on physical properties of ions. However, some general behaviors of certain ions are observed:
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Ions that form soluble compounds NO3-, C2H3O2-, Group 1A ions (Li+, Na+, K+, Rb+, Cs+), NH4+
Exceptions
None (compounds involving any of these ions will be soluble)
Cl-, Br-, I-
Ionic compounds with these ions are soluble, except for compounds involving Ag+, Hg22+ and Pb2+
SO42-
Ionic compounds with these ions are soluble, except for compounds involving Ca2+, Sr2+, Hg22+ and Pb2+
Metathesis reactions in which a nonelectrolyte (or weak electrolyte) forms Ions can chemically combine to form a nonelectrolyte (or weak electrolyte)
Even though the nonelectrolyte may be soluble in aqueous solution, its formation is essentially irreversible. Thus, ions are effectively removed from solution by this irreversible process
The neutralization reaction of HCl and NaOH is an example of this type of reaction:
The net ionic equation for this neutralization reaction is:
The formation of the covalent compound (H2O) from the proton and hydroxide ions is essentially irreversible and drives the metathesis reaction (even though we would consider H 2O to be "highly soluble" in H2O)
Metathesis reactions in which a gas forms If a possible metathesis reaction involves the formation of a gas (and the gas is not particularly soluble in H2O) the loss of the gas can drive the metathesis reaction (i.e. the ions react to form a gas, and the gas is lost - therefore, it is an irreversible process)
Typical gasses that can form from ionic compounds include H 2S (hydrogen sulfide - smells like rotten eggs) and CO2 (carbon dioxide)
Formation of carbon dioxide from carbonic acid
The bicarbonate ion (HCO3-) can combine with a proton to produce carbonic acid (H2CO3):
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Carbonic acid in water is unstable and decomposes to form water and carbon dioxide gas:
The carbon dioxide gas is lost, and thus the formation of carbonic acid is irreversible and drives the metathesis reaction
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 (Ca 2+ in this case). The metal cation participates in forming an ionic solid with the oxide (O2-) anion. Thus, the metal becomes oxidized
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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
In addition to oxidation by O2, metals can be oxidized by both acids and salts
We saw an example of the oxidation of a metal (Zn) by an acid (HCl) above:
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In this case, the metal is oxidized by aqueous protons to produce zinc cation, and the hydrogen is correspondingly reduced to neutral diatomic hydrogen gas Metals can also be oxidized by aqueous solutions of various salts:
o
In this case solid iron can be oxidized by aqueous nickel ions to produce iron cations, and the nickel ions are correspondingly reduced to produce nickel metal
Again, whenever one substance is oxidized, another substance is reduced in a redox reaction In the above example of the oxidation of iron by nickel:
Why doesn't the reaction go the other way around? Why doesn't iron oxidize nickel? Can we predict which metal will be preferentially reduced and which will be oxidized when elemental or ionic forms of different metals are combined?
The different metals differ in the ease with which they can be oxidized
Some metals, e.g. Li and K, are easy to oxidize Other metals, e.g. platinum (Pt) and gold (Au) are difficult to oxidize
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:
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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)
Metals can react with acids to form diatomic hydrogen gas and metal cations (i.e. the oxidation of the metal and the reduction of the hydrogen ion). Which metals can participate in such reactions?
Metals that are higher than H+ ions on the Activity Series can be oxidized by the H+ ions (i.e. can react with acids to form hydrogen gas) Metals below H+ ions on the Activity Series cannot be oxidized by H +, and therefore, will not react with acids to produce hydrogen gas (the metals below hydrogen represent the group of metals used in coins and jewelry)
The Solution Process Important characteristics of solutions:
They are homogenous mixtures
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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
1996 Michael Blaber
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