## Acid reactions
1. acid + metal (exc. Cu, Hg, Ag) $\longrightarrow$ salt + $H_{2(g)}$
-2. acid + metal carbonate $\longrightarrow$ salt + $CO_{2(g)} + H_2O _{(l)}$
+2. acid + metal carbonate $\longrightarrow$ salt + $CO_{2(g)} + H_2O_{(l)}$
3. acid + metal hydrogen carbonate $\longrightarrow$ salt + $CO_{2(g)} + H_2O_{(l)}$
4. acid + metal sulfite $\longrightarrow$ salt + $SO_{2(g)} + H_2O_{(l)}$
5. acid + metal sulfide $\longrightarrow$ salt + $H_2S_{(g)}$
## Bronsted-Lowry theory
-> **Acid:** donates a proton ($H^+$ ion)
+> **Acid:** donates a proton ($H^+$ ion) r
> **Base:** accepts a proton from another substance
H atom is one proton and electron, so removing an electron leaves $H^+$ ion.
Ionic bases dissolving in $H_2O$ - ionic compounds dissociate into constituent ions. Not ionised.
e.g. $NaOH_{(s)}\stackrel{\mathrm{H_2O}}{\longrightarrow}Na^+_{(aq)}+OH^−_{(aq)}$
+
+## pH values
+
+Acid/base/neutral not equivalent to pH (logarithmic scale)
+
+In water at 25C: $[H_3O^+] \times [OH^-]=10^{-7} \therefore \text{pH}=7$
+
+$$\text{pH} = -\log[H_3O^+]=-\log[H^+]$$
+$$[H_3O^+]=10^{-\text{pH}}$$