This is chapter 9, and this space will hopefully house an introduction to the chapter.
Properties of acids and basesEdit
Properties of acids and bases in aqueous solutionEdit
A reaction between a strong acid and a strong base produces water and a salt. A reaction between an acid that is not strong and a base that is not strong (can be only 1) produces only a salt. For example:
Acids reacting with metals will produce hydrogen gas, e.g. .
Acids reacting with carbonates will produce water and carbon dioxide, e.g. .
Experimental properties of acids and basesEdit
When acids and bases neutralize, the reaction is noticeably exothermic (i.e., heat can be felt coming from the reaction). Obviously, they will have an effect on the color of indicators as described below. The hydrogen produced in the reaction of acids with metal will produce a 'pop' sound if a match is held to it, and the CO2 from the carbonate reaction will turn limewater a milky white when bubbled trough it.
Examples of some acids: , , , .
Examples of bases: , NH3, CH3COO-.
Brønsted-Lowry acids and basesEdit
For a compound to act as a B-L acid, it must have a hydrogen atom in it, which it is capable or losing while remaining fairly stable. A BL base must be capable of accepting a hydrogen ion while remaining relatively stable (or reacting to form a stable compound, e.g. water and a salt). Some compounds (such as water) may act as both, e.g. (H2O-> OH- or H3O+). These substances are called "amphoteric".
Acid base reactions always involve an acid-base conjugate pair. If the reactant is an acid, the matching product is its conjugate base, e.g. HCl/Cl-, CH3COOH/CH3COO-, NH4+/NH3.
The conjugate base will always have one less H atom than the acid (or the acid one more than the base). In compounds where there are many hydrogen atoms, the one which is held the weakest is generally the one which is lost, and this must be reflected in the writing of the compound, as in the CH3COOH example above.
Lewis defined acids as electron pair acceptors and bases as electron pair donors.
A simple proton-hydroxide reaction shows that this is equivalent to Brønsted-Lowry theory:
H+ + :OH- --> H-OH
The proton is accepted because the hydroxide has a pair of electrons with which it forms a covalent bond to the proton.
Lewis theory goes beyond Brønsted-Lowry theory to describe non-proton acid-base reactions:
F3B + :NH3 --> F3B-NH3
Ligands and nucleophiles can be classed as Lewis bases. Lewis acids include metal ions and electrophiles.
Ag+ + 2 :NH3 --> [Ag(NH3)2]+
Strong and weak acids and basesEdit
Strong and weak acids are defined by their ease of losing (or donating) a proton. A strong acid, when placed in water, will almost fully ionise/dissociate straight away, producing H3O+ ions from water. a weak acid will, however, only partially do this, leaving some unreacted acid remaining. This is set up as an equilibrium, and so when some of the H3O+ ions produced by a weak acid are reacted, LCP means that more of the acid will react to form H3O+ ions. This means that, given an equal number of moles of acid, they will be neutralized by the same amount of strong base, but their solutions will have different pH values. A weak base is the same as this, only it accepts protons and so produces OH- ions from water rather than H3O+. Any solution's ability to conduct electricity is defined by its ionic charges. As a result, a strong acid will produce more charged ions than a weak one, and so its solution will be a better electrical conductor than a weak acid. The same goes for strong/weak bases.
Examples of strong and weak acids and basesEdit
Strong acids : HCl, HNO3, H2SO4, HClO3, HBr .
Weak acids : CH3COOH, H2CO3. Strong bases : group 1 hydroxides (e.g. NaOH, etc.), BaOH.
Weak bases : NH3, CH3CH2NH2.
The strength of an acid or base can obviously be measured with a universal indicator or a pH meter, and also the rate of reaction. Hydrogen production with metals or CO2 with CaCO3 will reveal the strength of an acid. The relative acidities can also be found by neutralizing two acids with a strong base in the presence of an indicator.
The pH scaleEdit
pH values range up and down from 7 (being the neutral value of pure water at 25°C and 1 atmosphere). Lower pH value are acidic; higher values are basic. pH can be measured with a pH meter, or with pH paper (paper containing a mixture of indicators to cause a continuous color change). pH is a measure of the concentration of H3O+ ions.
If we have two solutions with their pH values, the lower one will be more acidic and the higher one will be more basic (though they could both still be basic/acidic with respect to water – pH 7).
9.4.3 : a change of 1 in the pH scale represents a 10 times change in the acidity or basicity of the solution (because it is a log scale). Concentration is equal to 10-pH or pH=-log10[H3O+] where [H3O+] is the concentration of hydrogen ions in solution.
Topic 18 is the additional HL material for Topic 9.
Brønsted-Lowry acids and bases
Calculations involving acids and basesEdit
This HL Sub-topic is Sub-topic 5 of SL Option A, with one difference: The SL Option does not require knowedge of Ka x Kb = Kw.
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Indicators change colour depending in the presence of acid or base. Here is a table of common indicators:
|Indicator||pKa||pH range||Acid colour||Alkali colour|
Table adapted from: Stark JG & Wallace, HG (1982) Chemistry Data Book p103 2nd ed, John Murray, London.