A spontaneous process is one that occurs naturally and without the need for external intervention, under appropriate conditions. In other words, we can say it happens on its own under specific circumstances.
While a nonspontaneous process does not occur naturally and requires some external influence or energy to take place.
Four examples of both processes are:
Spontaneous processes:
Water flowing downhill: When water flows from a higher place to a lower place, like a river running downhill, it naturally happens due to gravity, without needing extra energy.
Ice melting at room temperature: Ice will naturally melt into water above 0°C. It doesn’t need anything extra to make this happen.
A rock rolling down a hill: If a rock is placed at the hilltop, it will naturally roll down due to gravity, without needing any external force to make it move.
Rusting of iron: When iron is exposed to air and moisture, it will slowly rust (form iron oxide) over time. This is a spontaneous process that occurs naturally.
Water flowing downhill – Spontaneous Process
Non-spontaneous processes:
Water flowing uphill: For water to flow uphill, we need to pump it or apply energy (like in a fountain). Water doesn’t flow uphill by itself.
Freezing of water at room temperature: Water will only freeze into ice if we put that below 0°C. At room temperature, this does not happen naturally, and energy must be removed from the water to make it freeze.
Unmixing of a solution: If we mix sugar in water, the sugar dissolves. For the sugar to separate back out on its own, energy would need to be added, such as by evaporating the water.
Charging a battery: A battery cannot charge itself. We need to connect it to an external power source to add energy and recharge it.
This concept comes under Thermochemistry. This involves energy transfer and phase changes as the ice undergoes different stages—from being solid at lower temperatures to becoming liquid at higher temperatures.
To understand how heat is involved in these changes, we need to break down the underlying processes involved in temperature change, phase change, and the heat energy required at each step.
1. Sensible Heat and Temperature Change
Sensible heat is the heat energy required to change the temperature of a substance without changing its phase. In this case, we heat ice from -10°C to 0°C, which means the temperature of the solid ice increases.
The formula gives the heat required to achieve this temperature change:
Q=m⋅c⋅ΔT
Where:
Where, Q = heat energy
m = mass
c = specific heat
ΔT = temperature change
Ice’s specific heat is about 2.09 J/g°C. This means that to increase the temperature of 1 gram of ice by 1°C, 2.09 joules of energy are required.
2. Latent Heat and Phase Change (Melting Ice)
When the ice reaches 0°C, it will melt into water. Melting is a phase change, where the substance changes from solid to liquid. This transition occurs without a temperature change but requires a specific amount of energy to break the bonds between the ice molecules.
This energy is called latent heat (the heat required for a phase change), and for ice, it’s referred to as latent heat of fusion. The formula for latent heat is:
Q = n ΔH
Where, n= number of moles
ΔH = heat of fusion or vaporization depending on phase change
ΔH fusion= 6.02 kJ/mol
The temperature does not change during this process instead all the energy goes into breaking the bonds between the ice molecules, allowing them to transition into liquid form.
3. Sensible Heat in Water and Further Heating
Once the ice melts and turns into water, we can increase the temperature of the liquid water. Heating water requires sensible heat, just as it did for the ice; however, it involves a different value for the specific heat capacity.
Therefore, in this case, we heat the water from 0°C to 10°C. Consequently, the heat energy required for this temperature change is again calculated using:
Q=m⋅c(water)⋅ΔT
Where:
Q = heat energy
m = mass
c = specific heat
ΔT = temperature change
Cliquid= 4.21 J/g°C
Given:
Mass of ice =25.0 g
Temperature of ice = -10° C
Mass of water = 25.0 g
Temperature of water = +10° C
Step-by-step-solution
Since the molar mass of ice and water is the same and the given mass is also the same, the number of moles of both will be equal.
Step 1: To Find moles of ice:
Molar mass =18.015 g/mol
Moles=25.0 g x \(\frac{1mol}{18.05 g}\) = 1.388 mol ice
Moles of ice = Moles of water =1.388 mol
Step 2: Initially, Ice has a temperature of -10° C and it is converted to liquid water at a temperature of +10° C.
As a result, we will calculate three different heat values.
Heat required for temperature change from -10° C to 0° C, i.e. Q1
The heat required for phase change at 0° C solid ice to liquid water, i.e. Q2
Heat required for a temperature change from 0° C liquid water to +10° C liquid water, i.e. Q3
Calculation of three different heat:
1. To find heat required for a temperature change from -10° C to 0°C (Q1): Q1=m(ice) × C(ice) × [T (final) – T(initial)]
Q1 = 25.0 g × 2.09 J/g °C × [0° C-(-10° C)] =522 J or 0.522 kJ
2. To find heat required for temperature change at 0° C solid ice to 0°C liquid water(Q2): Q2=n(ice) × ΔH(fusion)
Q2= 1.388 mol × 6.02 kJ/g °C = 8.36 kJ
3. To find heat required for a temperature change from 0 °C to +10°C (Q3): Q3=m(water) × C(water) × [T (final) – T(initial)]
Q3= 25.0 g × 4.21 J/g °C × [10° C-(0° C)] =1052J 0r 1.05 kJ
Step :3
Therefore, the total heat required for the complete transition is the sum of the heat required for each individual transition step.
Q(Total) = Q1 + Q2 + Q3
= 0.522 kJ + 8.36 kJ + 1.05 kJ
= 9.93 kJ
Hence, the total heat required to convert 25.0 g of ice at – 10° C to 25.0 g of water at – 10° C is 9.93 kJ.
