• One of the first theories to explain the fact that all acids had similar reactions, was that of Arrhenius .
• This proposed that in aqueous solution all acids , to some extent (dependent on the strength of the acid), split up to form a hydrogen ion and an anion , i.e. for an acid HX:
HX (aq) H + (aq) + X
–
(aq)
• The hydrogen ion is hydrated, like all ions in aqueous solution, but some chemists prefer to show this reaction more explicitly with one water molecule forming a dative covalent bond to the hydrogen ion, to produce the H
3
O + ion (the hydronium ion ; also called hydroxonium ion and oxonium ion ). In these terms the above equation becomes
HX (aq) + H
2
O (l) H
3
O + (aq) + X
–
(aq)

3
+

• This also emphasises the fact that water is not an inert solvent , but is necessary for acid –base activity .
• Indeed solutions of acids in many non –aqueous solvents do not show acidic properties .
• For example a solution of hydrogen chloride in methylbenzene does not dissociate and hence, for example, it will not react with magnesium .
• Invoking the hydronium ion is useful in discussing some aspects of acid –base theory, such as conjugate acid –base pairs .
• But apart from this the simpler terminology of the hydrate proton / hydrogen ion, H + (aq), will be adopted in this book.

Reactions of Acids as reactions of
Hydrogen Ions
• The similar reactions of acids can be explained as all being reactions of the hydrogen ion.
• It is perhaps more accurate to write them as ionic equations, for example the reaction of an aqueous acid with magnesium can be written as:
Mg (s) + 2 H + (aq) Mg 2+ (aq) + H
2
(g)

• Bases were defined as substances that react with, and neutralise , acids to form water .
• Soluble bases ( alkalis ) form the hydroxide ion when dissolved in water,
• either because they are soluble and contain the hydroxide ion (as with NaOH ),
• or because they react with water to produce one (as with ammonia , carbonates and hydrogen carbonates ):
NaOH (aq) Na + (aq) + OH – (aq)
NH
CO
3
HCO
(aq) + H
3
2
–
3
–
2
O (l) NH
(aq) + H
2
(aq) CO
2
4
+ (aq) + OH
O (l) HCO
(aq) + OH
3
–
–
–
(aq)
(aq) + OH
(aq)
–
(aq)

• Aqueous acids and alkalis contain ions that are free to move, which explains why they conduct electricity to some extent.
• If an acid and an alkali are mixed the hydrogen ions and hydroxide ions react exothermically to form water:
H + (aq) + OH
–
(aq) H
2
O (l)
• This leaves the anion from the acid and the cation from the base in solution.
• If the water is then evaporated these combine to form a solid salt.
• For example if the acid were hydrochloric acid and the base sodium hydroxide:
Na + (aq) + Cl – (aq) NaCl (s)

• O (e.g. H
2
SO
4
) was originally thought to cause acidic properties. Later, H was implicated, but it was still not clear why CH
4 was neutral.
• Arrhenius made the revolutionary suggestion that some solutions contain ions & that acids produce H
3
O +
(hydronium) ions in solution.
Ionization
Cl H
+ O
H
H
H O
H
+
+
H
Cl
• The more recent Bronsted-Lowry concept is that acids are H + (proton) donors and bases are proton acceptors

Hydrochloric acid is neutralised by both sodium hydroxide solution and ammonia solution. In both cases, you get a colourless solution which you can crystallise to get a white salt
- either sodium chloride or ammonium chloride.
These are clearly very similar reactions. The full equations are:
The lone electron pair on the nitrogen (N) in ammonia is represented as a pair of dots. This electron pair forms the bond
+ with a proton H .
In the sodium hydroxide case, hydrogen ions from the acid are reacting with hydroxide ions from the sodium hydroxide - in line with the Arrhenius theory.

However, in the ammonia case, there don't appear to be any hydroxide ions!
You can get around this by saying that the ammonia reacts with the water it is dissolved in to produce ammonium ions and hydroxide ions:
This is a reversible reaction, and in a typical dilute ammonia solution, about 99% of the ammonia remains as ammonia molecules. Nevertheless, there are hydroxide ions there, and we can squeeze this into the Arrhenius theory.

Limitations of the Arrhenius theory
However, this same reaction also happens between ammonia gas and hydrogen chloride gas.
In this case, there aren't any hydrogen ions or hydroxide ions in solution - because there isn't any solution. The
Arrhenius theory wouldn't count this as an acid-base reaction, despite the fact that it is producing the same product as when the two substances were in solution. That's silly!

HCl Cl CH
3
COOH CH
3
COO -
NH
4
+ NH
3
HNO
3
NO
3
-
How does a conjugate pair differ?
H + transfer

• In this idea, the ionization of an acid by water is just one example of an acid-base reaction.
Cl H acid
+
+
H H
O H O + Cl
H H base conjugate acid conjugate base conjugate acid-base pairs
• Acids and bases are identified based on whether they donate or accept H + .
• “Conjugate” acids and bases are found on the products side of the equation. A conjugate base is the same as the starting acid minus H + .

• The usual contemporary definition of an acid is the
Brønsted-Lowry definition, that an acid is a substance that acts as a donor of hydrogen ions (a hydrogen ion of course consists of just a proton, so acids are also often referred to as ‘ proton donors ’).
• This means that when it dissolves in water it produces a solution containing hydrogen ions and hence fits with the Arrhenius definition , but extends this to other solvent systems so that reactions such as:
NH
4
+ + NH
2
– 2 NH
3
• In a non –aqueous solvent (such as liquid ammonia ) are also classified as acid –base reactions .
• According to the same definition, a base is a substance that acts as an acceptor of hydrogen ions
(‘ proton acceptor ’).

• For a species to act as an acid it must contain a hydrogen atom attached by a bond that is easily broken in many cases this hydrogen is attached to an oxygen atom.
• For a substance to act as a base , it must have a non –bonding electron pair that can be used to form a bond to a hydrogen ion.
• Usually this lone pair is on an oxygen or nitrogen atom.

• When an acid loses one hydrogen ion , the species produced is referred to as the conjugate base of the acid, for example, the conjugate base of H
2
SO
4 is HSO
4
–
.
• Similarly the species formed when a base gains one hydrogen ion is referred to as the conjugate acid of that base.
• The ammonium ion,NH
4
+ , is therefore the conjugate acid of ammonia, NH
3
.
• Acid –base reactions , which can be recognised because they involve the transfer of a hydrogen ion , therefore always involve two such acid –base pairs .

Consider ethanoic acid dissolving in water in these terms:
• It can be seen that the acid and its conjugate base in these two pairs:
• (CH
3
COOH / CH
3
COO
– and H
3
O + / H
2
O)
• differ only in the loss of a single hydrogen ion.

Identify the acid and base species in the following equation:
CO
3
2(aq) + H
2
O(l) HCO
3
(aq) + OH -
Acid species : H
2
O(l) HCO
3
(aq)
Base species : CO
3
2(aq) OH -

STRENGTH OF CONJUGATE ACID – BASE PAIRS
• If an acid is a strong acid (such as HCl), then its conjugate base (Cl – ) will be such a weak base that it can be considered non –basic so that the equilibrium below will lie fully to the right.
• As the strength of an acid (HB) decreases however the position of the equilibrium below shifts to the left , which is equivalent to an increase in the strength of its conjugate base (B
–
).
• Eventually with a strong base (e.g. OH – ) the equilibrium lies so far to the left that the conjugate acid (H
2 be regarded as non –acidic .
O) may

• Some species, like the water molecule and the hydrogensulfate ion ,
• can act as both acids and bases and are therefore described as amphiprotic :

An amphiprotic species is a species that can act as either an acid or a base (it can lose or gain a proton), depending on the other reactant.
Consider water:
H
2
O + CH
3
O  OH + CH
3
OH acid base
H
2
O + HBr  H
3
O + + Br base acid

• Most acids and bases only lose or gain one hydrogen ion and so are said to be monoprotic , but other acids and bases that can gain and that can lose more hydrogen ions and are said to be polyprotic .
• Sulfuric acid for example, is diprotic in aqueous solution because it can lose two hydrogen ions forming first the hydrogensulfate ion and then the sulfate ion:

• Similarly phosphoric(V) acid , found in cola drinks , is triprotic :
• Polyprotic acids and bases may be recognised because they form anions with more than one charge .
• Carbonic acid (aqueous carbon dioxide), for example, must be diprotic because it forms the carbonate ion , which has a charge of minus two (CO
3
2 – ).

Identify the acid, base, conjugate acid, conjugate base, and conjugate acid-base pairs:
HC
2
H
3
O acid
2
(aq)
+ H
2 base
O
(l) 
C
2
H
3
O
2
– conjugate base
(aq)
+ H
3
O + (aq) conjugate acid conjugate acid-base pairs
OH base
– (aq)
+ HCO acid
3
– (aq) 
CO
3
2 – (aq) conjugate base
+ H
2
O
(l) conjugate acid conjugate acid-base pairs

HF
(aq)
+ SO
3
2 – (aq) 
F
– (aq)
+ HSO acid base conjugate base
3
– (aq) conjugate acid conjugate acid-base pairs
CO
3
2 – (aq)
+ HC
2
H
3
O
2 base acid
(aq) 
C
2
H
3
O
2
– conjugate base
(aq)
+ HCO
3
– (aq) conjugate acid conjugate acid-base pairs
H
3
PO acid
4
(aq)
+ OCl base
– (aq) 
H
2
PO
4
– (aq) conjugate base
+ HOCl
(aq) conjugate acid conjugate acid-base pairs

HCO acid
3
– (aq)
+ S 2 – (aq) 
HS
– (aq)
+ CO base conjugate acid
3
2 – (aq) conjugate base conjugate acid-base pairs
H
2
CO acid
3
(aq)
+ OH base
– (aq) 
HCO
3
– conjugate base
(aq)
+ H
2
O
(l) conjugate acid conjugate acid-base pairs
H
3
O + (aq)
+ HSO
3
– (aq) 
H acid base
O
(l)
+ H SO
(aq) conjugate acid-base pairs
OH – (aq) base
+ HSO acid
3
– (aq) 
H O
(l)
+ SO
3
2 – (aq) conjugate base conjugate acid-base pairs
For more lessons, visit www.chalkbored.com

• When a base accepts a proton from an acid it forms a covalent bond to the proton, but this differs from most covalent bonds in that both of the electrons come from the base , as the proton has no electrons to contribute to the bond.
• Covalent bonds of this sort are known as
‘ dative ’ or ‘ co –ordinate ’ covalent bonds, but are identical to other covalent bonds in every way but the origin of the electrons.
• Dative bonds are indicated in structural formulae by an arrow , pointing in the direction that the electrons are donated , rather than a line.

• Lewis therefore pointed out that an acid could be defined as ‘ a species that accepts a pair of electrons to form a dative bond ’.
• All Bronsted –Lowry acids are in fact Lewis acids ,
• for example when hydrogen chloride dissolves in water the water molecule forms a dative bond to the hydrogen ion to generate the hydronium ion:

• The term ‘ Lewis acid ’ is however usually reserved for a species that is not also a
Bronsted –Lowry acid.
• The substance that donates the electron pair to form the bond to these is known as a
‘ Lewis base ’.
• This extended the range of acid-base reactions (beyond those involving the transfer of a hydrogen ion) to include all reactions involving the formation of a dative bond .

3
• A common example of a Lewis acid is boron trifluoride , in which boron has only six electrons in its valence shell .
• This reacts with ammonia (which acts as a
Lewis base ) to give a compound containing a dative bond (note the arrow in the structural formula), in which the lone pair from the nitrogen completes the valence shell of the boron :

• Other common Lewis acids are compounds of elements in group 3 of the periodic table, such as aluminium chloride ( AlCl
3
) in which the element forms three covalent bonds , leaving a vacancy for two electrons in the valence shell .
• Any species that can accept an electron pair into its incomplete valence shell (e.g. CH
3
+ ) is, however, capable of acting as a Lewis acid .
• Similarly any species with a nonbonding electron pair ( all anions and indeed all molecules that are not hydrides of group 3 and group 4 elements) is capable of acting as a
Lewis base .

• All interactions to form‘ complex ions ’ are also Lewis acid –base reactions .
• In these the ligand acts as the Lewis base by donating a pair of electrons that is accepted by the central metal ion , which hence acts as a Lewis acid .
• For example:

Al(H
2
O)
6
3+

We are going to look in detail at the bonding in the complex ion formed when water molecules attach themselves to an aluminium ion to give Al(H
2
O)
6
3+

Start by thinking about the structure of a naked
. aluminium ion before the water molecules bond to it.

Aluminium has the electronic structure
1s
2
2s
2
2p
6
3s
2
3p x
1

When it forms an Al
3+
ion it loses the 3-level electrons to leave
1s
2
2s
2
2p
6


That means that all the 3-level orbitals are now empty. The aluminium uses of six of these to accept lone pairs from six water molecules.

It re-organises (hybridises) the 3s, the three 3p, and two of the 3d orbitals to produce six new orbitals all with the same energy.

You might wonder why it chooses to use six orbitals rather than four or eight or whatever.

Six is the maximum number of water molecules it is possible to fit around an aluminium ion (and most other metal ions).

By making the maximum number of bonds, it releases most energy and so becomes most energetically stable.

Only one lone pair is shown on each water molecule.
The other lone pair is pointing away from the aluminium and so isn't involved in the bonding. The resulting ion looks like this:
Dotted arrows represent lone pairs coming from water molecules behind the plane of the screen or paper.
Wedge shaped arrows represent bonds from water molecules in front of the plane of the screen or paper.

• Several metal hydroxides such as Zn(OH)
2 and Al(OH)
3 are amphoteric.
• When these behave as acids the central metal ion acts as a Lewis acid and accepts a lone pair from the hydroxide ion, which acts as a Lewis base.
• For example :
• Al(OH)
3
(s) + OH
–
(aq) Al(OH)
4
–
(aq)
Lewis acid Lewis base anion from acid


Any oxygen-containing substance which will produce an acid, when dissolved in water, is called an ACID ANHYDRIDE.

When sulfur dioxide (SO
2
) is dissolved in water, sulfurous acid is formed.
SO
2
+ H
2
O H
2
SO
3

When sulfur trioxide (SO
3
) is dissolved in water, some sulfuric acid is formed.
SO
3
+ H
2
O H
2
SO
4

• Carbon dioxide dissolved in water is in equilibrium with carbonic acid:
CO
2
+ H
2
O ⇌ H
2
CO
3
• The hydration equilibrium constant at 25 °C is K h
= 1.70
×10 −3 : hence, the majority of the carbon dioxide is not converted into carbonic acid and stays as CO
2 molecules .

How a hydrogen carbonate act as a base?
+
HCO
3
–
(aq) + H
2
O(l) H
2
CO
3
(aq) +OH
–
(aq)
H
2
CO
3
(aq) CO
2
(aq) + H
2
O(l)
HCO
3
–
(aq) CO
2
(aq) + OH
–
(aq)


Any oxygen-containing substance which will produce a base, when dissolved in water, is called a BASIC ANHYDRIDE.

If sodium oxide (Na
2
O) is added to water, sodium hydroxide, a base, is formed.
Na
2
O + H
2
O 2 NaOH

ANHYDRIDES

Anhydride means "without water", so anhydrides may be classified as acids or bases with all the water removed.

It might be more accurate and understandable if we say anhydrides are acids or bases with all the hydrogen removed in the form of water.

Given a particular acid or base, we can determine its anhydride by removing two hydrogen atoms and one oxygen atom from its molecule (or molecules). For example:
2 HClO Cl
2
O + H
2
O
Ca(OH)
2
CaO + H
2
O

The acid anhydride for hypochlorous acid, HClO, is
Cl
2
O. The basic anhydride for calcium hydroxide is
CaO.


In predicting anhydrides,

ENOUGH H
2
O UNITS MUST BE REMOVED TO
LEAVE THE ANHYDRIDE WITHOUT ANY
HYDROGEN .

To form the anhydride for an acid like H
3
PO
4
, remove three water molecules from two phosphoric acid molecules to produce the anhydride, P
2
O
5
.
2 H
3
PO
4
P
2
O
5
+ 3 H
2
O

Arrhenius only in water
•
• Acids – produce H +
Bases - produce OH -
Bronsted-Lowry any solvent
•
• Acids
Bases
– donate H
– accept H
+
+
Lewis used in organic chemistry, wider range of substances
• Acids – accept e pair
• Bases – donate e pair

The hydrogen ion in aqueous solution
H + + H
2
O  H
3
O + (hydronium ion)
Arrhenius
HCl NaOH
Bronsted-Lowry
Lewis
HCl
BF
3
HCN
:NH
3
NH
3

Acids are corrosive chemicals with a sour taste
(note: you should never taste chemicals as many are poisonous.).
All acids have certain chemical characteristics in common:
• They form solutions with a pH < 7 , so that indicators which change colour at about this pH give the same reaction with all acids, for example, they turn blue litmus red.
• They react with active metals (those above hydrogen in the reactivity series ) to give a salt and hydrogen gas.
For example, sulfuric acid reacts with magnesium to give magnesium sulfate and hydrogen:
H
2
SO
4
(aq) + Mg (s) MgSO
4
(aq) + H
2
(g) acid metal salt hydrogen

• They react with bases , such as metal oxides and hydroxides to form a salt and water .
For example, copper(II) oxide dissolves in nitric acid to form a solution of copper(II) nitrate and water:
2 HNO
3
(aq) + CuO (s) Cu(NO
3
)
2
(aq) + H
2
O (l) acid metal oxide salt water
Phosphoric acid reacts with sodium hydroxide to form sodium phosphate and water:
H
3
PO
4
(aq) + 3 NaOH (aq) Na
3
PO
4
(aq) + 3 H
2
O (l) acid metal hydroxide salt water

• They react with metal carbonates and hydrogen carbonates to give a salt , water and carbon dioxide , which appears as effervescence
(bubbles).
For example, hydrochloric acid will react with zinc carbonate to form zinc chloride,water and carbon dioxide:
2 HCl (aq) + ZnCO
3
(s) ZnCl
2
(aq) + H
2
O (l) + CO
2
(g) acid metal carbonate salt water carbon dioxide
Ethanoic acid reacts with sodium hydrogencarbonate to form sodium ethanoate, water and carbon dioxide.
CH
3
COOH (aq) + NaHCO
3
(aq) NaCH
3
COO (aq) + H
2
O (l) + CO
2
(g) acid metalh ydrogen salt water carbon carbonate dioxide

•
• Originally a base was considered to be any substance that reacted with an acid to neutralise it, but now the term has more precise meanings.
• The most common bases are the oxides , hydroxides and carbonates of metals , but a number of other compounds, such as ammonia and amines also act as bases.
• Aqueous solutions of bases , known as alkalis (for example aqueous sodium hydroxide), have a slippery feel and a bitter taste
(though, again, you should never taste them).

All bases have certain chemicalcharacteristics in common:
• If they are soluble in water they give a solution with pH>7
, so that they will all have a similar effect on indicators that change colour at about this pH, for example, they turn red litmus blue
.
• They react with acids to form a salt
. For example, calcium oxide will react with hydrochloric acid to form calcium chloride and water:
CaO (s) + 2 HCl (aq) CaCl
2
(aq) + H
2
O (l) base acid salt water

BASES
ACIDS
• Taste sour
• Taste bitter
• Affect indicators
• Affect indicators e.g. Turn litmus e.g.Turn litmus
• Electrolytes (conduct
• Electrolytes electricity)
• pH between 0 and <7
(conduct electricity)
• pH between >7 and 14
• Neutralize Acids
(Antacids)
• Neutralize Bases
• React with active metals
– Fe, Zn
• Feel soapy or slippery
(react with fats to make soap)

Acids
HCl
H
2
SO
4
HNO
3
Juices, Soda
Bases
NaOH
Ca(OH)
2
KOH
Soap,
Ammonia,
Baking Soda

hydrochloric acid gastric juices sodium hydroxide vinegar lemon juice tomato juice coffee milk of magnesia pure water rain water blood baking soda wine oven cleaner salty water limewater coca cola toothpaste bleach urine

When an acid is dissolved in water, what kind of solution could it be called ?
When a base is dissolved in water, what kind of solution could it be called ?
When a solution is neither acidic nor alkaline, what could it be called ?
ACIDIC
ALKALINE
NEUTRAL
What name is given to chemicals which can show whether a solution is acidic, alkaline or neutral ?
INDICATORS

• Strong acids are those which are almost completely dissociated (ionised) in dilute aqueous solution:
HX (aq) H + (aq) + X
–
≈ 0% ≈ 100%
(aq)
• This means that such solutions are good conductors of electricity , owing to the presence of mobile ions .
• Hydrochloric acid is a typical example of a strong acid:
HCl (aq) H + (aq) + Cl
–
(aq)
• Other common strong acids include sulfuric acid
(H
2
SO
4
) and nitric acid (HNO
3
).
H
2
SO
4
HNO
3
(aq) H + (aq) + HSO
4
–
(aq) H + (aq) + NO
3
–
(aq)
(aq)

• Generally speaking, in strong acids the hydrogen is bonded either to a more electronegative element (such as Cl , Br and I in HX ), or to an oxygen bonded to a nonmetal (as in H
2
SO
4
).
• In the binary hydrogen halides , the acid strength increases down the group :
HI > HBr > HCl
• Note that HF is a weak acid , contrary to expectations and unlike the other hydrogen halides .
• This is because, although HF is a highly polarized bond , it is also a very strong bond.
• Another factor is that the HF molecule , unlike the other hydrogen halides, can strongly hydrogen bond to water and this stabilises the undissociated molecule .

• In the oxyacids the strength of the acid increases with
• the electronegativity of the non-metal :
H
2
S O
4 acid. is a strong acid, H
3
P O
4 a weaker
• and the number of oxygens present
(HNO
3 acid).
is a strong acid, HNO
2 a weaker

In oxyacids, in which an OH is bonded to another atom, Y , the more electronegative
Y is, the more acidic the acid.

For a series of oxyacids, acidity increases with the number of oxygens .

• The more polar the H-X bond and/or the weaker the H-X bond , the more acidic the compound.
• Acidity increases from left to right across a row and from top to bottom down a group.

• A weak acid is one which is only slightly dissociated into ions in dilute aqueous solution:
HA (aq) H+ (aq) + A – (aq)
≈ 99% ≈ 1%
• A typical example of a weak acid is ethanoic acid, where the undissociated acid is in equilibrium with the ions:
CH
3
COOH (aq) H + (aq) + CH
3
COO
–
• Almost all organic acids are weak acids .
(aq)
• Similarly aqueous carbon dioxide behaves as a weak acid:
CO
2
(aq) + H
2
O (l) H + (aq) + HCO
3
–
(aq)

Other common inorganic weak acids are:
• aqueous sulfur dioxide ( SO
2
)
) (analogous to aqueous
CO
2
• hydrofluoric acid ( HF , as noted previously)
• hydrocyanic acid ( HCN )
• intermediate ions of polyprotic acids (e.g. HSO
4
– )
• cations formed by weak bases (such as NH
4
+ )
• Hydrated ions of small highly charged metal ions (e.g.
Al 3+ (aq) ) also act as weak acids.

Differentiation of Strong and Weak Acids
Strong and weak acids can be differentiated by comparing solutions of equal concentrations. The concentration of hydrogen ions in the solution of the weak acid will be considerably lower, giving rise to a number of differences that may be tested experimentally:
• A weak acid has a higher pH than a strong acid of equal concentration.
• Weak acids do not conduct electricity as well as strong acids of equal concentration , but they conduct electricity better than water .
• Weak acids react more slowly in typical acid reactions
(such as those with a carbonate to give carbon dioxide or with an active metal to give hydrogen gas) than strong acids of equal concentration.

• In the same way a strong base is one which is completely dissociated into ions in aqueous solution,
• like sodium hydroxide and barium hydroxide :
NaOH (aq) Na + (aq) + OH
–
(aq)
Ba(OH)
2
(aq) Ba 2+ (aq) + 2OH
–
(aq)

• With weak bases an equilibrium exists between the base and the hydroxide ions so that, for example, ammonia is only partially converted to the hydroxide ion in aqueous solution:
NH
3
(aq) + H
2
O (l) NH
4
+ (aq) + OH
–
(aq)
• The closely related amines, such as ethylamine
(C
2
H
5
NH
2
) also act as weak bases.
C
2
H
5
NH
2
(aq) + H
2
O (l) C
2
H
5
NH
3
+ (aq) + OH –
• The anions formed by weak acids (such as the
(aq) carbonate ion, ethanoate and phosphate ions) also act as weak bases,for example:
CO
3
2 – (aq) + H
2
O (l) HCO
3
– (aq) + OH – (aq)

Differentiation of Strong and Weak
Bases
• Methods for differentiating strong and weak bases are similar to those for strong and weak acids; for solutions of equal concentration
• a strong base will have a higher pH and
• a greater electrical conductivity .

In chemistry care must be taken to use the terms
• strong and weak (meaning fully and partially dissociated ) correctly and not as synonyms for
• concentrated and dilute (meaning containing a large or small number of moles in a given volume ) as is done in everyday speech.
• The ‘chemical’ use of the term is also to be found, in ‘ strong electrolyte ’ and ‘ weak electrolyte ’.
• The term electrolyte means forming ions in aqueous solution allowing it to conduct electricity.

• The term strong electrolyte refers to a substance that is completely converted to ions in aqueous solution (such as salts, strong acids and strong bases) whilst weak electrolyte refers to those only partially converted to ions (such as weak acids and bases).
• Note that only a very small fraction (< 1 in 108) of molecules in pure water is split up into ions, so it is a very weak electrolyte and hence a poor conductor of electricity.