We previously looked at the Bohr model of the atom to describe atomic spectra. The Bohr model was an important step in conceptualizing the quantization of energy for electrons in an atom. We can more fully explain the emission spectrum of an atom using quantum theory. We follow a similar idea in that when an electron in a hydrogen atom absorbs energy, it jumps to a higher energy level and is said to be in an excited state. When the electron jumps back down to a lower energy level, it emits a photon with an energy equal to the difference in the two energy levels. For hydrogen, the energy levels are determined by the principal quantum number (n). Mathematically, we can calculate the change in energy for the electron by subtracting the initial energy from the final energy of the electron. The change in energy is equal to -2.18 times 10 to the -18 times the quantity one divided by the final principal quantum number squared minus one divided by the initial principal quantum number squared. The change in energy is measured in joules. If the change in energy for the electron is positive, the electron has absorbed energy. This means the electron has jumped from a lower energy level to a higher energy level. If the change in energy for the electron is negative, the electron has released energy in the form of a photon. This means the electron has jumped from a higher energy level to a lower energy level. If we use this equation for hydrogen, we will obtain four wavelengths in the visible spectrum corresponding to transitions from n=3, 4, 5, or 6 to n=2. Later on, we will learn that the energy levels for multi-electron atoms are adjusted due to sub-level splitting.