A hot lump of 39.4 g of iron at an initial temperature of 52.4 °C is placed in 50.0 mL of H2O initially at 25.0 °C and allowed to reach thermal equilibrium. What is the final temperature of the iron and water, given that the specific heat of iron is 0.449 J/(g·°C)? Assume no heat is lost to the surroundings.

By applying the concept of conservation of energy and equating the heat lost by iron to the heat gained by water, we can solve for the final agreed upon temperature of the system.

The final temperature of the iron and water can be calculated using the concept of conservation of energy. According to this concept, the heat gained by water will be equal to the heat lost by iron, in the absence of any heat loss to surroundings.

The heat lost by iron can be calculated by multiplying the mass of iron, the change in its temperature and the specific heat of iron. Similarly, the heat gained by water can be calculated by multiplying the mass of water, the change in its temperature and the specific heat of water (remember to convert milliliters to grams if density is 1 g/mL for water).

Setting these two calculations equal to each other, we can solve for the final temperature (which is the same for both water and iron in this scenario).

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