calculate the pH of a solution prepared by dissolving 2.05 g of sodium acetate, CH3COONa, in 78.0 mL of 0.25 M acetic acid, CH3COOH(aq). Assume the volume change upon dissolving the sodium acetate is negligible. ka of CH3COOH is 1.75 x 10-5
Moles of CH3COONa= 2.0582.0= 0.025 molMoles~ of~ CH3COONa= ~\cfrac{2.05}{82.0} = ~0.025 ~molMoles of CH3COONa= 82.02.05= 0.025 mol
Volume V=78.0 mL=0.078 LVolume~ V = 78.0 ~mL = 0.078~ LVolume V=78.0 mL=0.078 L
[CH3COONa]= 0.0250.078= 0.32 M[CH_3COONa]=~ \cfrac{0.025}{0.078} = ~0.32~M[CH3COONa]= 0.0780.025= 0.32 M
[CH3COOH]=0.25 M[CH_3COOH] = 0.25~ M[CH3COOH]=0.25 M
pH= −log(Ka)+ log([CH3COONa][CH3COOH])pH =~ -log(K_a) +~ log\bigg(\cfrac{[CH_3COONa]}{[CH_3COOH]}\bigg)pH= −log(Ka)+ log([CH3COOH][CH3COONa])
pH= −log(1.75∗ 10−5)+ log(0.320.25)pH =~ -log(1.75 *~10^{-5}) + ~log\bigg(\cfrac{0.32}{0.25}\bigg) \\pH= −log(1.75∗ 10−5)+ log(0.250.32)
pH=4.757+0.079=4.836pH= 4.757 + 0.079 = 4.836pH=4.757+0.079=4.836
Answer: pH = 4,836
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