Topic 9 - Acids and Bases

From KstructIB

[edit] 9.1 Properties of acids and bases

[edit] 9.1.1

Properties of acids and bases in aqueous solutions of stuff. (Note, the term alkali refers to a base dissolved in water).

Indicators : Change color depending on whether they’re in acidic or basic conditions. Each one’s different, so I suppose I’d better list some common ones…

	Acid	Base
Methyl orange	Red	Yellow
Bromophenol blue	Yellow	Blue
Methyl red	Red	Yellow
Bromothymol blue	Yellow	Blue
Phenolphtalein	Colorless	Red

Each one change color as a different pH, and so there will be cases where one is useful and others are not. (This info may not really be necessary is SL)

Reaction of acids with bases : They will often produce water, and the remaining components will combine to form a salt.
e.g. HCl + NaOH -> H₂O + NaCl.

Acids with metals : Will produce hydrogen
e.g. 2HCl + Mg -> MgCl₂ + H₂.

Acids with carbonates : Will produce water and CO₂ and a Salt

e.g. 2HCl + CaCO₃ -> CO₂ + H₂O + CaCl₂.

[edit] 9.1.2

Experimental properties of acids and bases :

• When acids and bases neutralise, the reaction is noticeably exothermic (ie heat can be felt coming from the reaction).
• Obviously, they will have an effect on the color of indicators as described above.
• The hydrogen produced in the reaction of acids with metal will produce a ‘pop’ sound if a match is held to it
• The CO₂ from the carbonate reaction will turn limewater a milky white when bubbled trough it.

Examples of some acids :

• HCl
• CH₃COOH
• H₂SO₄
• NH₄⁺

Examples of bases :

• NaOH
• NH₃
• CH₃COO^-.

[edit] 9.2 Bronsted-Lowry acids and bases

[edit] 9.2.1

According to the bronsted-lowry theory, acids are defined as proton (H⁺ ion) donators and bases are defined as proton acceptors.

[edit] 9.2.2

For a compound to act as a BL acid, it must have a hydrogen atom in it, which it is capable of 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 i.e. a water and a salt).

Some compounds (such as water) may act as both a BL acid and a BL base i.e. (H₂O-> OH^- or H₃O⁺)

[edit] 9.2.3

Acid base reactions always involve an acid-base conjugate pair. One is an acid and one is its conjugate base.

• HCl and Cl^-
• CH₃COOH and CH₃COO^-
• NH₄⁺ and NH₃

[edit] 9.2.4

The conjugate base will always have one less H atom that 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.

[edit] 9.3 Strong and weak acids and bases

[edit] 9.3.1

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 mols 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 is charges ions in it. A strong acid will produce more charged ions than a weak one, and so it’s solution will be a better electrical conductor than a weak acid. The same goes for strong/weak bases.

[edit] 9.3.2

Strong acids :

• HCl
• HNO₃
• H₂SO₄

Weak acids :

• CH₃COOH
• H₂CO₃

Strong bases :

• group 1 hydroxides (i.e. NaOH etc)
• BaOH

Weak bases :

• NH₃
• CH₃CH₂NH₂

[edit] 9.3.3

The strength of an acid or base can obviously be measured with an indicator (universal) or a pH meter. Also the rate of reaction measured by hydrogen production with metals or CO₂ with CaCO₃ will reveal the strength of an acid.

The relative acidities (I’m assuming that means diprotic or whatever) can also be found by neutralizing two acids with a strong base in the presence of an indicator.

[edit] 9.4 The pH scale

[edit] 9.4.1

pH vales range up and down from 7 (being the neutral value of pure water at 20c and 1 atm). Lower pH values 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 dissociation of an acid or base, and also of the concentration of that acid / base (actually its related to the concentration of H₃O⁺ ions).

[edit] 9.4.2

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).

[edit] 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’s a logarithmic scale). Concentration is proportional to 10^pH.