Question

Suppose that the microwave radiation has a wavelength of 12.4 cm . How many photons are required to heat 295 mL of coffee from 25.0 ∘C to 62.0 ∘C? Assume that the coffee has the same density, 0.997 g/mL , and specific heat capacity, 4.184 J/(g⋅K) , as water over this temperature range.

Answer #1

first we calculate the change in temperature.

T = 62 - 25

T = 37 c = 37
K

the mass of the water is given by

Mass of water = volume of water * density of water

M =295 ml * 0.997 g/mL

M = 294.115 g

the energy needed to change the temperature. is given by :-

Q = T * m * specific
heat constant

Q = 37 K * 294.115 g * 4.184 J/(gK)

Q = 45531.3549 J

energy required by a single photon

Energy required by a single photon (photon) = h * c /

E(photon) = 6.63x10^{-34} Js * 3x10^{8} m/s / 0.124
m

E(photon) = 1.604x10^{-24} J

no of photond required :-

no of photons = total energy / energy of single photon = E /
E(photon)

n = 45531.3549 J / 1.604x10^{-24} J

n = 2.8386 x10^{28}

Suppose that the microwave radiation has a wavelength of 11.2 cm
. How many photons are required to heat 255 mL of coffee from 25.0
∘C to 62.0 ∘C? Assume that the coffee has the same density, 0.997
g/mL , and specific heat capacity, 4.184 J/(g⋅K) , as water over
this temperature range.

Suppose that the microwave radiation has a wavelength of 12 cm .
How many photons are required to heat 255 mL of coffee from 25.0 ∘C
to 62.0 ∘C? Assume that the coffee has the same density, 0.997 g/mL
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temperature range.

± Using Microwave Radiation to Heat Coffee
Microwave ovens use microwave radiation to heat food. The
microwaves are absorbed by the water molecules in the food, which
is transferred to other components of the food. As the water
becomes hotter, so does the food.
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