One of the syllabus statements in my syllabus was: "Each orbital [in an atom] has a defined energy state for a given electronic configuration and chemical environment."
There's nothing wrong with this statement from my perspective.
It is common and in fact necessary for spectroscopists to discriminate between orbitals and states. A common approach when trying to determine the allowed energies of electrons in atoms is to define/calculate orbital characteristics by solving for a single electron. For an atom, we call this a "hydrogenic" system, because hydrogen has only one electron, but the approach can work for molecules as well. One-electron orbitals have a characteristic energy and probability density. However with the exception of hydrogen, atoms have more than one electron. What you would do then is fill electrons into the orbitals determined for a one-electron system, but the electrons in a multi-electron system interact with each other. These electron-electron interaction energies can be treated for after the fact using various variational or perturbation methods.
Thus we define a State as a particular arrangement or configuration of electrons, and the State Energy differs from the Orbital Energy because of the specific electron-electron interactions in the given electron configuration.
Example:
In helium you have two electrons. The simplest way to solve this problem is to determine the hydrogenic orbitals for a nuclear core charge of 2, and then fill in both electrons into this orbital. Then you need a way to estimate, because it cannot be determined analytically, how much the true energy differs of this orbital differs from that of a hydrogen orbital by virtue of the fact that the two electrons interact with each other. There are many theoretical methods available to do this.
In the Ground State, which is the lowest energy state, both electrons are in a "1S orbital", where "1S" is a designation for a hydrogen orbital. No, it is not energetically the same as a hydrogen 1S orbital, but it shares many of the same characteristics (shape, say) of a hydrogen 1S orbital, so we usually preserve the nomenclature. The spins of these electrons are opposite due to the Pauli Exclusion rule.
In an Excited State, one of these electrons is promoted to a higher lying orbital, say the "2S orbital". Now you have one electron in the 2S orbital and one electron in the 1S orbital - this is your electron configuration. As in the Ground State, this Excited State has a specific associated energy based on the interaction energy between the two electrons, which is different from the interaction energy between the electrons when they are in the Ground State. Morever, because the two electrons are in different orbitals, they no longer need to spin pair. Thus for this configuration you actually have two States - one in which the spins are parallel and one in which they are antiparallel. (Actually, the spin designations are a bit more complicated than this, but in the interest of keeping things simple...) As you might guess, the mutual orientations of the spins also have different energetic interactions. These two States are usually called singlet and triplet states. And then, you actually have coupling between the spin states and the orbital angular momentum of the electrons, which leads to even more possible states.
The important thing is that each of these States has its own unique, quantized energy value, and the number of states depends on what the electron (orbital) configuration is, because this defines what the potential interactions are. It is transitions between these Electron States that are measured by various spectroscopy techniques, and which also give rise to, for example, atomic and astronomical spectral patterns that are characteristic of the elements (i.e., it's how we determine the chemical composition of faraway stars). You might imagine that the number of possible States quickly gets very large, and spectroscopists have developed bookkeeping methods, such as Term Symbols, that can be used to quickly refer to what States are involved in a given spectroscopic transition.
http://en.wikipedia.org/wiki/Term_symbolThe "Chemical Environment" aspect of the statement I think just refers to the fact that the energies of the various (atomic) states are sensitive to other nearby electrostatic systems. If you have two atoms and bring them close enough together, the electrons will begin to interact with each other (and the electrons in one atom will interact with the nucleus of the other). Alternatively, it can mean the solvent your atom/molecule is dissolved in, which can greatly affect the state energies. (Look at the solvent-dependence of fluorescence color, for example.) Another good example that is relevant is the Zeeman Effect, or the associated Stark Effect.
http://en.wikipedia.org/wiki/Zeeman_EffectIf you put some systems in an external magnetic field, previously degenerate States (i.e., states that have the same energy) will split because they interact with the external field differently. In the aforementioned example of helium, I called one of the Excited States a "triplet". It is called a triplet because it actually involves three states that are usually degenerate. These states are distinguished by the direction of the total electron spin. (One in which the spins of the two electrons are both "up", one in which they are both "down" and one in which they are antiparallel.) Normally, it doesn't matter which way the spin vector is going. But it does matter in an external field, where one state becomes slightly higher in energy, and state becomes slightly lower in energy, and the other stays the same. Triplet = three. Spectroscopically a transition to this state from the Ground State would look like one peak normally, but would look like a triplet when performed in an external magnetic field. This is a good example of how the external environment can impact the energies of atomic or molecular States, even if it's not usually a factor that is important to calculating the properties of an orbital.