Questions on chapter 8 - SL

An alkali refers to a base that is dissolved in water.

Indicators : Change color depending on whether they’re in acidic or basic conditions.

	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.

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.

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 through it.

Examples of some acids :

• HCl
• CH₃COOH
• H₂SO₄
• NH₄⁺

• NaOH
• NH₃
• CH₃COO^-.

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

For a compound to act as a bronsted-lowry 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 abronsted-lowry acid and a BL base i.e. (H₂O-> OH^- or H₃O⁺)

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

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

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 can be described by an equilibrium reaction with H₃O⁺ (or H⁺) ions on the right-hand side.

Strength of an acid: How easily it dissolves to produce H₃O⁺ ions.
Concentration of an acid: How many moles of acid per dm³ the liquid contains.

If one compares two solutions of different acids that have the same concentration and temperature, the stronger acid will have more ions in the solution. Ions lead electricity so a stronger acid will have a higher conductivity. Reactions of an acid with metals, bases and carbonates will also go faster if the acid is strong because there are more H₃O⁺ ions produced that can react. Stronger acids have furthermore a lower pH value i.e. a larger concentration of H₃O⁺ ions.

If one compares two solutions of different bases that have the same concentration and temperature, the stronger alkali will have more ions in the solution. Ions lead electricity so a stronger alkali will have a higher conductivity. Reactions of an alkali will also go faster if the alkali is strong because there are more OH^- ions produced that can react. Stronger alkalis have furthermore a lower pOH value (higher pH value) i.e. a larger concentration of OH^- ions.

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

Strong acids :

• Hydrochloric acid - HCl
• Nitric acid - HNO₃
• Sulphuric acid - H₂SO₄
• Hydrobromic acid - HBr
• Perchloric acid - HClO₄

• CH₃COOH
• H₂CO₃
• C₆H₈O₇

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

• NH₃
• CH₃NH₂
• CH₃CH₂NH₂

pH values ranges from 0 to 14 (7 being the neutral value of pure water). 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 (or with other words 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).

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.