6.2 b: pH and pOH

pH and pOH

6.2 b: calculate pH, pOH, hydrogen ion concentration ([H+ ]) and hydroxide ion concentration ([OH– ]) for a range of solutions

d: write ionic equations to represent the dissociation of acids and bases in water, conjugate acid/base pairs in solution and amphiprotic nature of some salts, for example:

– sodium hydrogen carbonate

– potassium dihydrogen phosphate

​What we will learn:

• Recap - definitions of acids and bases

• recap - types of equations

• pH definition

• Calculating pH given H⁺ concentration

• pOH definition

• Calculating pOH given OH^- concentration

​Types of equations

• When we are dealing with neutralisation reactions, we can present them in several different equation types:

• Neutral species - complete equation with all reactants and products fully formed into compounds

• Complete ionic - complete equations with all reactants and products but split into ions

• Net ionic - only presents the chemical species that are involved in the reaction (in neutralisation reactions, only the water, as all other species are ionised into it)

​Definitions of acids and bases

Arrhenius and Bronsted-Lowry

Arrhenius:
- acids produce H+ in solutions
- bases produce OH- in solutions

Bronsted-Lowry
- acids are proton donors
- bases are proton acceptors
- 'proton' essentially meaning a H molecule - makes sense with Arrhenius because when a H loses an electron, only a proton is left in the nucleus

Which is the 'better' theory?

Bronsted- Lowry
- explains the behaviour of acids and bases outside of water, but also useful in other solvents
- explains the behaviour of bases without an OH
- allows use of conjugate acids and bases
- limitations are explained by Lewis theory - describes them as electron pair donors and acceptors

Conjugate acids and bases

Amphiprotic and polyprotic acids

Amphiprotic substances are able to both donate and accept a proton - therefore, can act as either an acid or a base
It will act as an acid when it reacts with a stronger base, and will act as a base if it reacts with a stronger acid

A polyprotic substance has the ability to donate more than one proton
Diprotic - can donate two. Triprotic - can donate three

Amphiprotic acids

Polyprotic acids

​pH definition

• log measurement of the amount of H+ in a solution

• pH = -log₁₀[H⁺]

• pH is measured on a logarithmic scale, so a single digit change really means a 10-fold change

• smaller pH value = higher concentration of H+ = more acidic

​pH definition

• there is a quantitative relationship between the pH and pOH that becomes self-evident when describing the autoionisation of water and observing the mathematical proof of Kw

• where Kw = [H+][OH-]

​Calculating pH

Open Ended

Calculate the pH of a substance with a [H+] of 0.0010

Open Ended

Calculate the [H+] of a substance with a pH of 4.35

​Calculating pOH

Open Ended

Calculate the pOH of a substance with a [OH-] of 6.3 x 10^-8 M

Open Ended

Calculate the [OH-] of a substance with a pOH of 7.9

​Self ionisation of water

• mostly water is happy by itself, but sometimes it self-ionises and gives away a H+ --> meaning that it becomes OH^- and the H⁺ will join with another water molecule to become H₃O⁺

• H_2​O⇌H₃​O⁺+OH⁻

• Kw=[H3​O⁺][OH⁻]=1.0×10⁻¹⁴ at 25°C

• because pH + pOH = 14

• Can use this to calculate ion concentrations in solution when you know one unknown

​Next up

• strengths of acids and bases