From 2c9d76f085a66b9c5428915899a4164aa34ada7f Mon Sep 17 00:00:00 2001 From: eggy Date: Wed, 5 Jan 2022 22:14:15 -0500 Subject: [PATCH] chem: Add pH/pOH --- docs/sch4uz.md | 26 ++++++++++++++++++++++++++ 1 file changed, 26 insertions(+) diff --git a/docs/sch4uz.md b/docs/sch4uz.md index 57e9e8d..76a3f08 100644 --- a/docs/sch4uz.md +++ b/docs/sch4uz.md @@ -684,6 +684,32 @@ Practically, this means comparing the rate of reaction with a metal or water or ### pH and pOH +This section will assume Bronsted-Lowry theory. + +pH represents $\ce{[H3O+]}$ logarithmically on a scale from 0 to 14. + +$$ +\ce{pH = -\log\big[H3O+_{(aq)}\big]} \\ +\ce{pOH = -\log\big[OH-_{(aq)}\big]} +$$ + +!!! warning + The number of sigfigs in pH is equal to the number of digits **after the decimal place**. + +A solution is **neutral** (neither acidic nor basic) when $\ce{[H3O+] = [OH-]}$. This happens to be $\ce{pH = 7}$ at SATP. In pure water, this is true as a small number of water molecules react with each other. + +In an equilibrium reaction between an acid and a base, $\ce{K_c = \frac{[H3O+][OH-]}{[H2O]}}$, but water has a constant concentration, so the equilibrium of the two ions is represented with the **water ionisation constant** $K_w$ is used. +$$K_w = \ce{[H3O+][OH-] = 1.00\times10^{-14} @ SATP}$$ + +As temperature **increases**, $K_w$ increases, therefore changing the pH of neutrality, but this may not necessarily change the acidity of the solution as the ion concentration is still the same. + +As pH increases, $\ce{[H3O+]}$ decreases, so $\ce{[OH-]}$ must increase to keep $K_w$ constant and maintain equilibrium. +$$\ce{pK_w = pH + pOH}$$ + +At 25°C, $\ce{pK_w = 14.0000}$, so: +$$\ce{14 = pH + pOH}$$ + + ## Organic chemistry !!! definition