The Bohr model of the atom is the most recent, scientifically accepted model for the atom, shown below.

The model denotes negative electrons orbiting the positive nucleus (containing protons and neutrons) in discrete shells. In the Bohr model, energy levels refer to the specific orbits or shells where electrons exist around the nucleus of an atom. Each level corresponds to a discrete energy value electrons can occupy. Electrons cannot exist between energy levels, but can transition between these levels by absorbing or emitting photons.

The ground state in the Bohr model of the atom refers to the energy level an electron can occupy which is closest to the nucleus.

Ionisation refers to when an electron is given enough energy so that it can overpower the attraction of the positive nucleus and is removed from the atom. The minimum energy required to remove an electron is called the ionisation energy.

In the Bohr model of the atom, an electron at the ionisation level, when it has enough energy to have been removed from the atom, is taken to have zero potential energy.

A consequence of this is that all other energy levels have a negative potential energy value. Using this convention, calculations of the energy released/absorbed when an electron transitions energy level shows that going from a higher to a lower energy level results in a negative energy change ($E=E_{lower}-E_{higher}$), meaning that energy is released from the system in the emission of a photon, and going from a lower to a higher energy level results in a positive energy change ($E=E_{higher}-E_{lower}$), meaning that energy enters the system in the absorption of a photon.

The following relationships link the energy of electrons at different energy levels and the frequency of the radiation emitted/absorbed, with Planck’s constant, and can be used to solve problems involving those quantities.

$$ E_2-E_1=hf $$

$$ E=hf $$

Continuous emission spectra show the full range of colours in the visible spectrum.

These are produced by sources that emit white light (incandescent sources, e.g. filament lamps, stars, etc.) because of the multitude of electron energy level transitions (from a higher energy level to a lower energy level, resulting in the emission of a photon) possible in these sources which are continuously occurring, covering all wavelengths of light from violet to red, producing a continuous band of colour.

Line emission spectra show discrete lines of colour.

Line emission spectra are specific to the atoms giving off the light; a signature which shows the presence of an element. These are produced due to specific electron energy level transitions (from a higher energy level to a lower energy level, resulting in the emission of a specific photon) occurring within elements when electrons within them are excited, corresponding to individual wavelengths of light being emitted, to produce discrete lines of colour. Different transitions result in different photons being emitted.

Absorption spectra show a continuous spectra which is missing distinct wavelengths of light, shown as black in place of individual lines of colour.

Like, line emission spectra, these are a signature which is specific to the atoms absorbing the light and can be used to identify what element is present. They are produced as a result of specific electron energy level transitions (from lower to higher energy levels as they absorb photons of a certain frequency) when the element is exposed to white light, leading to specific colours being missing from the continuous emission spectrum where they have been absorbed. Different photons can be absorbed due to different transitions.