 # Reaction Order and Rate Law

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Data, Calculations, and Questions

A. Calculate the initial and final concentrations as needed to complete Tables 1 and 2.

Data Table 1: Varying the Concentration of 1.0 M HCl

| | | | |Concentrations |
| |# Drops |# Drops |# Drops |Initial |

|# Drops |# Drops |# Drops |Initial |Initial |Final |Final |Reaction Time (sec) |Reaction | |Well # |HCl |Water |Na2S2O3 |HCl |Na2S2O3 |HCl |Na2S2O3 |Trial 1 |Trial 2 |Avg |Rate (sec-1) | |1 |8 |0 |12 |1 M |0.3 M |0.4 |0.18 |18.4 |16.3 |17.35 |0.0576 | |2 |8 |6 |6 |1 M |0.15 |0.4 |0.0045 |37.1 |37.9 |37.5 |0.0267 | |3 |8 |8 |4 |1 M |0.1 |0.4 |0.02 |107.2 |106.6 |106.9 |0.0093 | |

B. Calculate the average reaction time for each reaction by adding the times for the two trials and dividing by 2.

C. Calculate the reaction rate by taking the inverse of the average reaction time, i.e., 1 divided by the average reaction time.

1. Use table 1 to determine the reaction order for HCl.
2. Use table 2 to determine the reaction order for Na2S2O3.

Remember, you want to see what happens to the reaction rate when you double the concentration of one reactant while the second reactant remains unchanged. In Part
1, we varied the concentration of HCl while we kept the concentration of Na2S2O3 the same. In Part 2 we varied the concentration of Na2S2O3 while keeping the concentration of HCl the same. These are experimental data and results will be different from some of the nice, even numbers you saw on textbook problems. For example, in this experiment you may double the concentration of a reactant but the reaction rate may increase anywhere from 1.7 times to 2.4

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