Ori - All four metal chloride salts are highly soluble in water. Thus it should be very easy to prepare clear aqueous solutions of them. If you say that mildly acidic conditions are also optional then everything should be fine even without addition of any other chemicals. For example, according to the link cited below "Magnesium chloride dissolves in water to give a faintly acidic solution (pH = approximately 6)":

https://www.chemguide.co.uk/inorganic/period3/chlorides.html

Aqueous solutions of ZnCl2 have a pH of ca. 4:

https://www.cs.mcgill.ca/~rwest/wikispeedia/wpcd/wp/z/Zinc_chloride.htm

When not absolutely pure, CaCl2 sometimes contains a minor amount of CaO which could cause a slightly basic reaction. In that case you can add a few drops of HCl solution to redissolve it. I would not suggest to add any NaOH, ammonia, buffer solutions etc.

why don't you use of PBS (phosphate buffer saline)?

You can use ammonia to increase the pH, however, their respective complexes will form, but no precipitation of insoluble species will occur.

Calculate and plot the solubility curves of the hydroxides: these cations can be hydrolysed, you might need acidic conditions to avoid the precipitation of solid hydroxides. Then calculate the amount of HCl solution needed... and try the dissolution (of your chloride salts) by stiring eventually overnight, eventually heating a little, eventually adding little excesses of HCl. Measure the pH but do not adjust it by adding NaOH: this could induce local precipitation of hydroxides that will be difficult to dissolve back to equilibrium conditions. Using a buffer might induce complexation of your dications by the basic form of the buffer. Good chemists never adjust pH: they titrate or use known amounts of reactants to exactly know what is in the solution and to calculate the resulting pH. pH measurement is a control

Simetimes ammonium acetate buffer (pH~7) is better then others

Ori - All four metal chloride salts are highly soluble in water. Thus it should be very easy to prepare clear aqueous solutions of them. If you say that mildly acidic conditions are also optional then everything should be fine even without addition of any other chemicals. For example, according to the link cited below "Magnesium chloride dissolves in water to give a faintly acidic solution (pH = approximately 6)":

https://www.chemguide.co.uk/inorganic/period3/chlorides.html

Aqueous solutions of ZnCl2 have a pH of ca. 4:

https://www.cs.mcgill.ca/~rwest/wikispeedia/wpcd/wp/z/Zinc_chloride.htm

When not absolutely pure, CaCl2 sometimes contains a minor amount of CaO which could cause a slightly basic reaction. In that case you can add a few drops of HCl solution to redissolve it. I would not suggest to add any NaOH, ammonia, buffer solutions etc.

Role of Divalent Metal Ions on Activity and Stability of Thermostable Dipeptidase from Bacillus stearothermophilus Hong-Yon eRO.t Katsuyuki T ANIZA WA, * and Kenji SODA

https://www.tandfonline.com/doi/pdf/10.1271/bbb.61.1688?needAccess=true

Thank you all for your insightful answers !

However, if I would like to get a specific pH for all my solutions, (eg. pH 6) how can I obtain it?

I understand that NaOH will induce precipitations. Maybe I can slowly titrate with NaOH, try to re-dissolve it, filter the solution and measure the final concentration.

Dear Ori Brookstein as I mentioned in my previous answer, salts like magnesium chloride already have a pH around 6. Instead of waiting for more answeres here or doing calculations, why don't you just try it out? Just dissolve the desired amount of your metal chloride in a suitable amount of water and measure the pH? If necessary, add a few drops of HCL solution in order to adjust the pH to ca. 6. This should be a matter of half an hour instead of waiting several weeks for more suitable answers. Adding NaOH should alsobe unnecessary.

Dear Ori Brookstein ,

You have selected nicely two types of coordinated metal ions, the so-called borderline/soft ions (like Zn2+ and Cu2+) that coordinate with nitrogen- and sulphur-containing groups found in His and Cys, and the so-called hard cations (like Ca2+ and Mg2+) that use oxygen ligands as the preferred coordination sites (especially Asp and Glu). See for example: Gregory, D. S., Martin, A. C., Cheetham, J. C., & Rees, A. R. (1993). The prediction and characterization of metal binding sites in proteins. Protein Engineering, Design and Selection, 6(1), 29-35.

Protein metal ion interaction studies need to be performed in a proper buffer. Buffer constituents are required to maintain protein integrity, prevent protein aggregation etc. (unless you work with peptides). Examples of good buffers are: 20 mM Hepes/50 mM NaCl (pH 7.0) or a simple 5 mM Tris (pH 7.2) buffer.

All salts you mentioned are nicely soluble in water, so at 50 mM concentrations there are no problems to be expected. As said by Frank T. Edelmann just give it a go without a buffer as well as with a buffer (in my view the preferred one). Using a buffer for your protein will mean that you can simply add a bit of your ion solution (50 mM is in terms of metal ligand binding pretty high).

Best regards.

I. About speciation at solution of zinc chloride in water (acidic), you may want to check my posts elsewhere at this forum: https://www.researchgate.net/post/when_i_prepare_a_zinc_chloride_solution_a_white_precipitate_is_shown_plz_can_someone_tell_me_why_and_how_i_can_avoid_this_problem https://www.researchgate.net/post/Inconsistent_results_titrating_zinc_chloride-hydrochloric_acid_soldering_fluxes

II. About pH of calcium chloride solution:

https://www.researchgate.net/post/Whats-the-best-choice-to-modify-the-PH-of-calcium-chloride-hexahydrate

https://www.researchgate.net/post/Why-is-my-anhydrous-calcium-chloride-not-going-into-solution

III. On magnesium chloride solutions: https://www.researchgate.net/post/I_tried_to_prepare_a_MgCl2_4M_solution_using_anhydrous_magnesium_chloride_overnight_but_doesnt_dissolve_completely_in_water_Any_ideas

https://www.researchgate.net/post/How-to-make-accurate-solutions-of-MgCl2-6H2O

IV. On precipitation of Cu(OH)2 from Cu(II) aq. sol.: https://www.researchgate.net/post/Is_it_possible_to_ppt_out_CuOH2_at_acidic_pH_4-6

https://www.researchgate.net/post/Hi-everyone-why-is-there-a-precipitate-in-my-01M-CuSo4-aq-solution https://www.researchgate.net/post/How_to_adjust_the_pH_of_copperII_sulfate_to_7