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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. The solvent - usually the substance in greater concentration

2. 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. The second solution contains a chemical which reacts in a defined way, with known

stoichiometry, with the solute of the first solution

2. 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 Cl

-

ion

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 H2O. 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 For sulfuric acid, the first ionization is complete (as indicated by the single arrow),

therefore it is a strong electrolyte

o However, only some of the molecules of HSO4- undergo a second ionization. Thus, the

deprotonization of HSO4-, and the protonization of SO4

2- are significant reactions (as

indicated by the double arrows)

o Aqueous solutions of sulfuric acid will therefore contain some mixture of protons (H+),

HSO4- and SO4

2-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. The commonly used acids, hydrochloric, nitric and sulfuric, are all strong acids

2. Several of the strong acids are a combination of hydrogen and a halogen (the exception is HF,

which is a weak acid)

3. There are not a lot of strong acids, most acids are weak acids

4. There are not a lot of strong bases; the strong bases are metal hydroxides (group 1A and heavy

group 2A metals)

5. 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 Acid/base neutralization reactions that produce water from H+ and OH

- ions are an

example of this.

o 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

Exceptions

NO3-, C2H3O2

-, Group 1A

ions (Li+, Na

+, K

+, Rb

+,

Cs+), NH4

+

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+, Hg2

2+ and Pb

2+

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 H2O 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 H2S (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 Zn2+

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

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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:

o 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