What does electronic structure relate to

This page explores how you write electronic structures for atoms using s, p, and d notation. It assumes that you know about simple atomic orbitals - at least as far as the way they are named, and their relative energies.

If you haven't already read the page on atomic orbitals you should follow this link before you go any further. This isn't a mistake, but an effect of converting my original diagram into a lower quality gif image for efficient web use. UK syllabuses for 16 - 18 year olds tend to stop at krypton when it comes to writing electronic structures, but it is possible that you could be asked for structures for elements up as far as barium.

After barium you have to worry about f orbitals as well as s, p and d orbitals - and that's a problem for chemistry at a higher level. It is important that you look through past exam papers as well as your syllabus so that you can judge how hard the questions are likely to get. This page looks in detail at the elements in the shortened version of the Periodic Table above, and then shows how you could work out the structures of some bigger atoms. You must have a copy of your syllabus and copies of recent exam papers.

If you are studying a UK-based syllabus and haven't got them, follow this link to find out how to get hold of them. Hydrogen has its only electron in the 1s orbital - 1s 1and at helium the first level is completely full - 1s 2. Now we need to start filling the second level, and hence start the second period. Lithium's electron goes into the 2s orbital because that has a lower energy than the 2p orbitals.

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Lithium has an electronic structure of 1s 2 2s 1. Beryllium adds a second electron to this same level - 1s 2 2s 2. Now the 2p levels start to fill. These levels all have the same energy, and so the electrons go in singly at first. You wouldn't normally write them any differently from the other orbitals. You can see that it is going to get progressively tedious to write the full electronic structures of atoms as the number of electrons increases.

There are two ways around this, and you must be familiar with both. Shortcut 1: All the various p electrons can be lumped together. For example, fluorine could be written as 1s 2 2s 2 2p 5and neon as 1s 2 2s 2 2p 6. This is what is normally done if the electrons are in an inner layer. If the electrons are in the bonding level those on the outside of the atomthey are sometimes written in shorthand, sometimes in full. Don't worry about this.Electronic Spectroscopy relies on the quantized nature of energy states.

Given enough energy, an electron can be excited from its initial ground state or initial excited state hot band and briefly exist in a higher energy excited state. Electronic transitions involve exciting an electron from one principle quantum state to another.

Without incentive, an electron will not transition to a higher level. Only by absorbing energy, can an electron be excited. Once it is in the excited state, it will relax back to it's original more energetically stable state, and in the process, release energy as photons.

Often, during electronic spectroscopy, the electron is excited first from an initial low energy state to a higher state by absorbing photon energy from the spectrophotometer.

If the wavelength of the incident beam has enough energy to promote an electron to a higher level, then we can detect this in the absorbance spectrum. Once in the excited state, the electron has higher potential energy and will relax back to a lower state by emitting photon energy.

what does electronic structure relate to

This is called fluorescence and can be detected in the spectrum as well. Often, during electronic transitions, the initial state may have the electron in a level that is excited for both vibration and rotation. This can be true for the ground state and the excited state. In addition, due to the Frank Condon Factor, which describes the overlap between vibrational states of two electronic states, there may be visible vibrational bands within the absorption bands. Therefore, vibrational fine structure that can be seen in the absorption spectrum gives some indication of the degree of Frank Condon overlap between electronic states.

When interpreting the absorbance and fluorescence spectra of a given molecule, compound, material, or an elemental material, understanding the possible electronic transitions is crucial. Assigning the peaks in the absorption spectrum can become easier when considering which transitions are allowed by symmetry, the Laporte Rules, electron spin, or vibronic coupling.

Knowing the degree of allowedness, one can estimate the intensity of the transition, and the extinction coefficient associated with that transition. These guidelines are a few examples of the selection rules employed for interpreting the origin of spectral bands.

Only a complete model of molecular energy diagrams for the species under investigation can make clear the possible electronic transitions. Every different compound will have unique energy spacing between electronic levels, and depending on the type of compound, one can categorize these spacings and find some commonality.

Although surprises in science often lead to discovery, it is more fortuitous for the interpreter to predict the spectra rather than being baffled by the observation. The following section will discuss the interpretation of electronic absorption spectra given the nature of the chemical species being studied. This includes an understanding of the molecular or elemental electronic state symmetries, Russell-Sanders states, spin multiplicities, and forbidden and allowed transitions of a given species.

As the light passes through the monochrometer of the spectrophotometer, it hits the sample with some wavelength and corresponding energy.

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The ratio of the initial intensity of this light and the final intensity after passing through the sample is measured and recorded as absorbance Abs. When absorbance is measured at different wavelengths, an absorbance spectrum of Abs vs wavelength can be obtained. This spectra reveals the wavelengths of light that are absorbed by the chemical specie, and is specific for each different chemical.

Many electronic transitions can be visible in the spectrum if the energy of the incident light matches or surpasses the quantum of energy separating the ground state and that particular excited state. An example of an absorbance spectrum is given below.

Here we can see the effect of temperature and also the effect of solvents on the clarity of the spectrum. We can see from these anonymous compounds that decreasing the temperature allows the vibrational fine structure to emerge. With a higher temperature, the vibrational transitions become averaged in the spectrum due to the presence of vibrational hot bands and Fermi Resonance, and with this, the vibrational fine structure is lost at higher temperatures.

The effect that the solvent plays on the absorption spectrum is also very important. It is clear that polar solvents give rise to broad bands, non-polar solvents show more resolution, though, completely removing the solvent gives the best resolution. This is due to solvent-solute interaction. The solvent can interact with the solute in its ground state or excited state through intermolecular bonding. For example, a polar solvent like water has the ability of hydrogen bonding with the solute if the solute has a hydrogen bonding component, or simply through induced dipole-dipole interactions.Custom Search.

As you I just discussed in the Spectral Lines page, electrons fall to lower energy levels and give off light in the form of a spectrum. These spectral lines are actually specific amounts of energy for when an electron transitions to a lower energy level. If you assume the energy levels of an atom to be a staircase; if you roll a ball down the stairs the ball only has a few "steps" that it can stop on.

what does electronic structure relate to

This is the same situation an electron is in. Electrons can only occupy specific energy levels in an atom. It most be on an energy level if it is in the atom. There is no in between. This is why you get lines and not a "rainbow" of colors when electrons fall. Jahann Balmer in derived an equation to calculate the visible wavelengths that the hydrogen spectrum displayed. These electrons are falling to the 2nd energy level from higher ones. This transition to the 2nd energy level is now referred to as the "Balmer Series" of electron transitions.

Johan Rydberg use Balmers work to derived an equation for all electron transitions in a hydrogen atom. Here is the equation:. If we wanted to calculate energy we can adjust R by multipling by h planks constant and c speed of light.The electron arrangement of an element is related to its position on the periodic table. The electron arrangement of sodium 2. The number of electron shells is therefore the same across a period but increases when moving down a group.

Electronic band structure

The electron arrangements of atoms help explain the properties of elements and the structure of the periodic table. When substances react, it is only the outer electrons in the atoms that are involved - the nature of the reactions which occur depend upon the number of outer shell electrons in the reacting atoms.

So, elements in the same group have similar chemical properties because they have the same number of electrons in their outer shell.

The electron arrangements of the first three elements in group The atoms of all group 1 elements have similar chemical properties and reactions because they all have one electron in their outer shell. Similarly, the atoms of all group 7 elements have similar chemical properties and reactions to each other, because they all have seven electrons in their outer shell. Electron arrangements and the periodic table The electron arrangement of an element is related to its position on the periodic table.

Electronic structure feature Link to the periodic table Number or numbers of shells Period number Number of electrons in outermost shell Group number Total number of electrons in all shells Atomic number The electron arrangement of sodium 2.

Electron arrangements and properties of elements The electron arrangements of atoms help explain the properties of elements and the structure of the periodic table.

Electron configuration

The electron arrangements of the first three elements in group 1: The atoms of all group 1 elements have similar chemical properties and reactions because they all have one electron in their outer shell. Number or numbers of shells. Period number. Number of electrons in outermost shell. Group number. Total number of electrons in all shells. Atomic number.Now that you are familiar with the structure of the atom, we can further explore how the structure and the light generated from each element are related.

In order to understand this relationship we need to look at the model of the atom proposed by Neils Bohr. Bohr stated that the structure of an atom had specific energy levels in which the electrons were located around the nucleus. Whenever an electron moves from one of these energy levels to another it must either gain or lose some energy. If the electron gains energy this is called an absorbance and if it lost energy this is called an emittance. Different elements have different energy levels so that is why different elements emit or absorb different amounts wavelengths of light.

Before going on we need to define a couple of terms often used to describe light:. Wavelengths found in the electromagnetic spectrum range of light can be measured in units as large as 10 3 meters radio waves to 10 meters gamma waves. For the wavelengths of visible light the light we see in color the most common units used are nanometers 10 -9 meters and Angstroms 10 meters.

In the case of light, frequency refers to the number of times a wavelength is repeated per second. The relationship between wavelength and frequency is related through the speed of light. If you know the frequency you can easily convert to wavelength using the speed of light and vice versa.

The full electromagnetic spectrum is generally shown with both measurements given:. The wavelengths and frequencies of the light emitted by an atom its emission spectrum is determined by its electronic structure. In the Bohr model shown above you can see that as each electron moves from a higher energy level orbit to a lower one, a different color is emitted. The numbers shown above the colors are the wavelengths that correspond to the color.

Each shade of color has a unique wavelength based on the unique distance and energy. Bohr also assumed that the electron can change from one allowed orbit to another:. Therefore, if the emitted radiation from a falling electron produces light and has a defined energy, then it must have a correspondingly defined frequency or wavelength:.

Gary, I would like to stop this animation when it gets to the page on "what does this have to do with the sun? Where R H is called the Rydberg constant and has a value of This relationship was defined in part by another scientist, Max Planck. In Planck's assumption, radiant energy is emitted in small bursts, known as "quanta". This equation is later found to be true for all EM radiant energy emitted or absorbed.In Chapter 3 "Atoms, Molecules, and Ions"we introduced the periodic table as a tool for organizing the known chemical elements.

A periodic table is shown in Figure 8. The elements are listed by atomic number the number of protons in the nucleusand elements with similar chemical properties are grouped together in columns.

Why does the periodic table have the structure it does? The answer is rather simple, if you understand electron configurations: the shape of the periodic table mimics the filling of the subshells with electrons. Let us start with H and He. These two elements make up the first row of the periodic table see Figure 8. Figure 8. The next two electrons, for Li and Be, would go into the 2 s subshell.

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For the next six elements, the 2 p subshell is being occupied with electrons. On the right side of the periodic table, these six elements B through Ne are grouped together Figure 8. The next subshell to be filled is the 3 s subshell.

what does electronic structure relate to

The elements when this subshell is being filled, Na and Mg, are back on the left side of the periodic table Figure 8. Next, the 3 p subshell is filled with the next six elements Figure 8. Instead of filling the 3 d subshell next, electrons go into the 4 s subshell Figure 8. The 4 s subshell is filled before the 3 d subshell.

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This is reflected in the structure of the periodic table. After the 4 s subshell is filled, the 3 d subshell is filled with up to 10 electrons. This explains the section of 10 elements in the middle of the periodic table Figure 8. The 3 d subshell is filled in the middle section of the periodic table. And so forth. As we go across the rows of the periodic table, the overall shape of the table outlines how the electrons are occupying the shells and subshells.

The first two columns on the left side of the periodic table are where the s subshells are being occupied. Because of this, the first two rows of the periodic table are labeled the s block The columns of the periodic table in which s subshells are being occupied.

Similarly, the p block The columns of the periodic table in which p subshells are being occupied. It could be part of the main body, but then the periodic table would be rather long and cumbersome. The periodic table is separated into blocks depending on which subshell is being filled for the atoms that belong in that section. The electrons in the highest-numbered shell, plus any electrons in the last unfilled subshell, are called valence electrons The electrons in the highest-numbered shell, plus any electrons in the last unfilled subshell.

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The inner electrons are called core electrons. The valence electrons largely control the chemistry of an atom. Their electron configurations abbreviated for the larger atoms are as follows, with the valence shell electron configuration highlighted:.

They all have a similar electron configuration in their valence shells: a single s electron. Because much of the chemistry of an element is influenced by valence electrons, we would expect that these elements would have similar chemistry— and they do.

The Electronic Structure Of The Atom

The organization of electrons in atoms explains not only the shape of the periodic table but also the fact that elements in the same column of the periodic table have similar chemistry.

The same concept applies to the other columns of the periodic table. Elements in each column have the same valence shell electron configurations, and the elements have some similar chemical properties.In atomic physics and quantum chemistry, the electron configuration is the arrangement of electrons in an atom, molecule, or other physical structure e.

Like other elementary particles, the electron is subject to the laws of quantum mechanics, and exhibits both particle-like and wave-like nature. Formally, the quantum state of a particular electron is defined by its wavefunction, a complex-valued function of space and time. According to the Copenhagen interpretation of quantum mechanics, the position of a particular electron is not well defined until an act of measurement causes it to be detected.

The probability that the act of measurement will detect the electron at a particular point in space is proportional to the square of the absolute value of the wavefunction at that point. Electrons are able to move from one energy level to another by emission or absorption of a quantum of energy, in the form of a photon.

Electronic Structure and the Periodic Table

Because of the Pauli exclusion principle, no more than two electrons may exist in a given atomic orbital; therefore an electron may only leap to another orbital if there is a vacancy there. Knowledge of the electron configuration of different atoms is useful in understanding the structure of the periodic table of elements. In bulk materials this same idea helps explain the peculiar properties of lasers and semiconductors.

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