Calorimetry help! I need to know why, in this video,the enthalpy calculated using the calorimetry equation ∆H=mc∆T is HIGHER than the enthalpy...

I watched the video and have an explanation for the two parts that confused you.

The host conducted an experiment in which he combined 100 g of 1.00 M HCl and 100 g of 1.00 M NaOH in a calorimeter, a container designed to prevent heat loss to the surroundings. He measured the temperature both before and after combining the two solutions. After the solutions were combined the temperature increased due to the heat released by the neutralization reaction. He used the following formula to calculate the heat change:

q = ∆H = sm∆T, where:

∆H = heat released or absorbed

s = specific heat capacity, which is 4.184 J/g-K for water

m = mass of substance absorbing or releasing heat

∆H = (4.184 J/g-k)(200. g)(7.4 K) = 6192.32 J = 6.2 kJ

The host explained that the value is positive because the temperature increases. In other words, the final temperature minus the initial temperature is positive. 

Here's the reason why the results give a positive value:

The reaction between HCl and NaOH is exothermic, meaning that heat is released. The heat released by the chemical reaction is absorbed by the water in which the chemicals are dissolved. The amount of heat lost by the reactants in forming products equals the amount of heat gained by the water, but is opposite in sign. What was calculated was the heat gained by the water, thus the positive sign. The heat of reaction is the opposite, -6.2 kJ. This wasn't pointed out.

The second point of confusion was the comparison of the values obtained using ∆H=sm∆T and Hess's Law. These two values should be equal but aren't due to experimental error. The Hess's Law calculation, which gave a value of -5.67 kJ, is the theoretical value for the neutralization reaction. The value was calculated using standard heats of formation of the products and reactants. The experimental value varies from the theoretical value by 9.3%. 

The host pointed out two possible sources of error that could account for the difference between the two values:

1. The specific heat capacity of pure water was used to calculate the heat absorbed by the solution, even though the solution was salt water and not pure water.

2. The heat absorbed by the calorimeter wasn't accounted for. 

The first reason is a better explanation. Error due to the calorimeter absorbing heat would produce a result that's smaller that the theoretical value, not larger. 

I hope this helps you better understand the problem worked out in the video.