The energy of an electron in an atom is negative, with the value zero corresponding to a free electron at rest. 'Least energy' implies the most negative energy value, corresponding to the most stable and tightly bound state.

1. Hydrogen (1s): The electron occupies the $n=1$ shell. The energy of the electron in the ground state of Hydrogen is $E_1 = -2.18 \times 10^{-18} \text{ J}$ (or $-13.6 \text{ eV}$).
2. Carbon (2p): The specified electron is in the $2p$ orbital ($n=2$). Although the nuclear charge ($Z=6$) is higher, the shielding by inner $1s$ electrons and the higher principal quantum number raise its energy compared to H($1s$). The first ionization enthalpy of Carbon is roughly $1086 \text{ kJ mol}^{-1}$ ($E \approx -11.3 \text{ eV}$), which is higher (less negative) than Hydrogen.
3. Sodium (3s): The outermost electron is in the $3s$ orbital ($n=3$). Due to shielding by inner shells ($n=1, 2$) and the higher $n$, it is loosely bound. Its ionization enthalpy is $496 \text{ kJ mol}^{-1}$ ($E \approx -5.1 \text{ eV}$).
4. Copper (4s): The specified electron is in the $4s$ orbital ($n=4$). Its ionization enthalpy is $745 \text{ kJ mol}^{-1}$ ($E \approx -7.7 \text{ eV}$).

Comparing the energy levels, the electron in the Hydrogen atom ($n=1$) has the most negative energy, thus the least amount of energy.