Chemistry

Assume that the normal blood buffer contains 0.00080 M carbonic acid and 0.0085 M hydrogen carbonate, the pKa = 6.35 for carbonic acid and the volume of blood in the body is 7.00 L. The blood pH, due to disruption is now 7.20. What is the ratio of [HCO3−]/[H2CO3] now that the blood has been challenged? How many moles of hydrogen carbonate must be added to the blood to bring the carbonic acid/hydrogen carbonate ratio back to a normal pH = 7.4? How many milliliters of 0.10 M hydrogen carbonate must be added to the blood?

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  1. First I would check to see what the pH is at these concentrations and I like to work in millimoles so
    base = 0.0085 x 7.00L = 59.5 mmols
    acid = 0.008 x 7.00L = 5.6 mmols.

    pH = 6.35 + log(59.5/5.6) = 7.38

    I look at the second part of the problem by asking how much acid must I add to the "normal" blood to make it pH of 7.20 and that will be how much bicarbonate I must add to get it back from 7.20 to normal 7.38.
    ........HCO3^- + H^+ ==> H2CO3
    I.......59.5..............5.6
    add..............x..............
    C.......-x......-x........+x
    E......59.5-x....0.......5.6+x

    Plug into the HH equation and solve for x, then add to 5.6 and subtract from 59.5

    I like to check to see that this actually produces the desired pH. Of course if this is how much acid must be added to produce that "off beat" pH of 7.20 it will be the bicarbonate that must be added to get it back to normal.

    You can find the ratio of the base/acid at pH 7.20 from the HH equation. Just plug in the values.

    E........................59.5+x

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