and bicarbonate ion parts in order to return to the preferred equilibrium

And bicarbonate ion parts in order to return to the

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and bicarbonate ion parts in order to return to the preferred equilibrium balance by producing bicarbonate in the process of shifting the equilibrium to the right. Reaction observed: H 2 CO 3 (aq) + H 2 O ( l ) + NaHCO 3 HCO 3 - (aq) + H 3 O + (aq) + Na (aq) (r) The pH of the solution increases, once again, to 7.69 as expected. The addition of sodium bicarbonate stimulates the production of bicarbonate in the solution which in result, acting as a weak base, increases the pH of the solution. The equilibrium concentrations are altered by the addition of sodium bicarbonate that dissociates into its sodium ion and bicarbonate ion parts in order to return to the preferred equilibrium balance by producing bicarbonate in the process of shifting equilibrium to the right. The condition of the body at this point is called acidosis and is dangerous if the buffer system is not able to neutralize such state. Reaction observed: H 2 CO 3 (aq) + H 2 O ( l ) + NaHCO 3 HCO 3 - (aq) + H 3 O + (aq) + Na (aq) (s) Adding a pellet of solid carbon dioxide caused the pH of the solution to be lowered to 7.38. During the process of pH decrease, continuous water fog was seen flowing out of the beaker. The dry ice, carbon dioxide, reacts with water to produce carbonic acid which, in nature, causes the decrease in solution’s pH. The equilibrium shifts to the production of carbonic acid because an increased amount of carbon dioxide on the left side needs to be accounted for in the equilibrium system, and therefore, carbonic acid is formed. Reaction observed: 2 H 2 O ( l ) + CO 2 (g) H 2 CO 3 (aq) + H 2 O ( l ) (t) The pH of the solution, following the addition of ammonium chloride, decreased by an increment of 0.01 to 7.37. Addition of ammonium chloride, being an acid, donates its hydrogen to contribute to the production of carbonic acid when comes in contact with bicarbonate. In such pathway, the equilibrium shifts towards the formation of carbonic acid as to return to preferred equilibrium position. Reaction observed: HCO 3 - (aq) + H 3 O + (aq) + NH 4 Cl (aq) H 2 CO 3 (aq) + H 2 O ( l ) + NH 3 (aq) + Cl - Conclusion: Part 1: The addition of species to a system at equilibrium demonstrated the system’s ability to favour forward and reverse reactions. This indicates that the products formed in previous reactions, such as [Cu(H 2 O) 4 ] 2+ (aq), are not guaranteed to remain in the form of product. Part 2: Multiple equilibrium systems have been demonstrated by forming a combination of several equilibria in one solution. The change in concentration of certain species, and involvement of other factors, resulted in equilibrium shifts and reaction direction changes to minimize the effect of change, as detailed in Le Chatelier’s principle.
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Part 3: The bicarbonate buffer system in blood was simulated and tested. We observe the manipulation of the concentration of H 3 O + via the amphoteric ion, HCO 3 - (aq), that can donate and/or accept H + ions.
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  • Fall '14
  • Pell, Wendy

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