HBr and HF are both monoprotic Arrhenius acids—that is, in aqueous solution, they dissociate and ionize to give hydrogen ions. A strong acid ionizes completely; a weak acid ionizes partially.
In this case, HBr, being a strong acid, would ionize completely in water to yield H+ and Br- ions. However, HF, being a weak acid, would ionize only to a limited extent: some of the HF molecules will ionize into H+ and F- ions, but most of the HF will remain undissociated.
pH is, by definition, a measurement of the concentration of hydrogen ions in solution (pH = -log[H+]). A higher concentration of hydrogen ions gives a lower pH, while a lower concentration of hydrogen ions gives a higher pH. At 25 °C, a pH of 7 indicates a neutral solution; a pH less than 7 indicates an acidic solution; and a pH greater than 7 indicates a basic solution.
If we have equal concentrations of HBr and HF, then the HBr solution will have a greater concentration of hydrogen ions in solution than the HF solution. Consequently, the pH of the HBr solution will be less than the pH of the HF solution.
Choice A is incorrect: Strong acids like HBr dissociate completely, not partially.
Choice B is incorrect: While the initial concentration of HBr and HF are the same, the H+ concentration in the HBr solution is greater. Since pH is a function of H+ concentration, the pH of the two solutions cannot be the same.
Choice C is correct: A greater H+ concentration gives a lower pH value. The HBr solution has the greater H+ concentration. Thus, the pH of the HBr solution would be less than that of the HF solution.
Choice D is incorrect for the reason why choice C is correct.
<span>All scientists try to base their conclusions on the experiments they have conducted and they have repeated them to be sure of it. </span>
The given question is incomplete. The complete question is :
Carbon tetrachloride can be produced by the following reaction:
Suppose 1.20 mol 
of and 3.60 mol of 
were placed in a 1.00-L flask at an unknown temperature. After equilibrium has been achieved, the mixture contains 0.72 mol of 
. Calculate equilibrium constant at the unknown temperature.
Answer: The equilibrium constant at unknown temperature is 0.36
Explanation:
Moles of 
= 1.20 mole
Moles of 
= 3.60 mole
Volume of solution = 1.00 L
Initial concentration of = 
Initial concentration of = 
The given balanced equilibrium reaction is,
Initial conc. 1.20 M 3.60 M 0 0
At eqm. conc. (1.20-x) M (3.60-3x) M (x) M (x) M
The expression for equilibrium constant for this reaction will be,
![K_c=\frac{[S_2Cl_2]\times [CCl_4]}{[Cl_2]^3[CS_2]}](https://tex.z-dn.net/?f=K_c%3D%5Cfrac%7B%5BS_2Cl_2%5D%5Ctimes%20%5BCCl_4%5D%7D%7B%5BCl_2%5D%5E3%5BCS_2%5D%7D)
Now put all the given values in this expression, we get :
Given :Equilibrium concentration of , x = 
Thus equilibrium constant at unknown temperature is 0.36
Answer:
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