Oxygenated tap water is rich in transition metals including Fe and Mn ions. Citric is a source of H+ and a good chelate and can drive a redox reaction in the presence of oxygen and H+ proceeding as follows:
$$\ce{4 Fe^2+/Mn^2+ + O2 + 2 H+ --> 4 Fe^3+/Mn^3+ + 2 OH-}$$
There is also a likely equilibrium reactions that can be effective in recycling ions (to continue the reaction, as occurs in natural waters) in the presence of citrate acting a chelate for, say, ferric:
$$\ce{Fe^2+ + Mn^3+ <=> Fe^3+ + Mn^2+}$$
$$\ce{Cu+ + Fe^3+ <=> Fe^2+ + Cu^2+}$$
Note: The above are active charge equilibrium reactions, referred to as a metal redox couple, and support the role of transition metals in their lower valence states engaging in notable radical creation formation via Fenton and Fenton-type (non-iron case) reactions. Further, in the presence of lab or sunlight, some further recycling of metal ions is possible.
Reference: See my comments and sources cited here.
The reaction with oxygen is electrochemical in nature with an inception period and proceeds with time.
Note: $\ce{H+}$ is consumed, so the $\mathrm{pH}$ would be expected to rise.
Here is also an electrochemistry reference, in particular, Table 2, where the first listed half-cell reaction corresponds to the above after adding two $\ce{H+}$ to both sides (creating water as a product). However, as this reaction can produce basic salts, as I have discussed previously, I believe it is more informative. Further, I do prefer my rendition, as per my prior link supplied above, it could also display a possible radical chemistry underpinning to the underlying half-cell reaction mechanics.