In this lesson, we will learn:
- To understand the experimental evidence leading to our current understanding of electrons.
- To recall the shape and nature of early electron orbitals and energy level diagrams.
- To fill electron orbitals for the first two rows of elements correctly according to the Pauli and Aufbau principles.
- There are many methods (they'll be seen soon) which we use to find out the shape and structure of chemical compounds, but why do atoms even form molecules in the first place? What causes different molecules to have different shapes?
These points might lead you to think that it's not the whole atoms that are influencing the shape or bonding of a larger molecule – it's the electrons and how many there are.
Recall the shapes of some molecules from C11.4.5: Molecular Geometry:
- BH3 is trigonal planar, but NH and PH3, with the same number of atoms, don't form the same shape as BH3.
- CH4 is tetrahedral and isoelectronic molecules (same number of electrons) like NH3 and water H2O still form the same tetrahedral structure without four H atoms, using their lone electron pairs instead.
- Some atoms do not form molecules; for some reason they are stable enough to exist as individual atoms, like helium (He).
- A lot of experiments have been performed and data collected on how electrons behave in atoms and molecules.
One of the major lines of experiment looks at the light an element emits (when electrified or heated in some way) called atomic emission spectroscopy.
When many metal samples are heated by flame, they emit light of a consistent color/character – very much like a fingerprint for each metal. Some include:
Experiments of this nature led to the discovery of rubidium and cesium, both of which are named after the Latin root of their flame colors.
Helium was also discovered by its spectral lines found in a solar eclipse – Helios means sun in Greek.
- Sodium: orange-yellow
- Strontium: red
- Copper: blue-green
- Potassium: lilac
- Magnesium: intensely bright colorless flame.
Experiments on hydrogen atoms, having only one electron, found it also has a distinct emission spectrum. Each time a particular element was experimented on in this way, its same unique emission spectrum was observed.
This led to the conclusion that all electrons in atoms have specific energy levels, or quantized energy:
What these "quantized" levels mean is that at any one time, all electrons are in a particular "energy level" state with a precise amount of energy associated with it.
- The emission spectra are produced when an element is energized (by heat or electric current) and its electrons are move from a stable low energy 'ground' state to a less stable higher energy excited state. Think of a person jumping in the air, moving literally from the ground state to mid-air, an unstable higher energy temporary state.
- When returning to the stable ground state, energy (light) of a precise frequency is released which is related to the gap in energy between the two states.
- This light has frequency in the visible region of the electromagnetic spectrum – that's why we see flame colors when metals are heated!
Like going up a tall building and taking the elevator; it will only stop at the precise floor levels. You can stop at the 5th or 6th floor but you cannot stop between them! The elevator's stops are quantized.
- These energy 'levels' lead us to Heisenberg's uncertainty principle: we cannot precisely know the location AND the momentum of an electron at the same time. Because the momentum is related to energy - which we do know - we cannot know the location of an electron!
This has major consequences for how we think about what an electron is and where it is:
In short, don't draw orbits, draw orbitals. Orbitals are like houses for electrons; they aren't always in there, but they spend a lot of their time there because that's where they live!
- Drawing electrons as particles in 'electron shells' orbiting around a nucleus is WRONG. The previous electron configurations showing electron 'orbiting' around the nucleus are inaccurate – they suggest we know where the electrons are and we don't! We know their energy, so we cannot know their location.
- Electrons are better thought of as a 'cloud' of electron density where they are most likely to be found, called an orbital.
The first, lowest energy orbital is called a 1s orbital.
The 1s orbital is spherical and covers the area closest to the positive nucleus. This is the lowest energy orbital because a negative electron will experience greatest attraction to a positive nucleus the closer it is.
- 1 refers to an energy level; the lower it is, the more stable it is and the closer to the nucleus it is.
- S refers to shape – think s for sphere.
All orbitals can hold two electrons, no more. This covers the electrons for hydrogen and helium, which you can show using an energy-level diagram.
- After the first two electrons are filled in the 1s orbital, the third and fourth electrons are found in the 2s orbital. This is where the third electron of Li and third/fourth for Be are found.
The 2s orbital is different from 1s however, because it has a node. A node is an electron 'dead zone' - there is no chance of the electron being found here.
- Again, s means spherical but 2 shows a higher energy level. This orbital is higher in energy than 1s because the electrons are normally further away from the nucleus.
Because this separates the orbital into two halves, it is convention to draw these halves as opposing regions (red/blue, +/-, shaded/unshaded). This is just because we describe electrons as part wave in nature.
- The 2nd energy level also has orbitals where the node is not spherical, but in a single plane. These are called p orbitals, and there are three in each energy level - one each for the x, y and z axes they occupy. They are considered "lobe shaped"
- With two electrons each, this gives the electron configurations for the elements boron through neon. See the diagram below: