Atomic Structure

On this page:

  • Atomic models
  • Periodic trends
  • Spectroscopy

Atomic Models

Democritus

John Dalton

John Dalton

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In Ancient Greece, Democritus made the first claims about the atom. Democritus said that atoms are tiny, indestructible particles of matter that make up everything. 

John Dalton

John Dalton

John Dalton

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In 1803, Dalton suggested that atoms were indivisible and each element is made of its own, unique atom. He also theorized that atoms are not created or destroyed in chemical reactions.

JJ Thomson

John Dalton

Ernest Rutherford

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In the Cathode Ray Tube experiment, Thomson discovered the electron and created the Plum Pudding model. This model shows that negatively charged electrons are combined with the "soup" of positive charge that makes up the atom.

Ernest Rutherford

Ernest Rutherford

Ernest Rutherford

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In the Gold Foil experiment, Rutherford discovered the nucleus of the atom. He showed that the nucleus has a positive charge while the electrons have a negative charge and float around the nucleus. 

Niels Bohr

Ernest Rutherford

Erwin Schrodinger

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Bohr discovered that electrons orbit the nucleus. According to the Bohr model, each electron is set in a fixed orbit around the nucleus. Unfortunately, the Bohr model is most accurate only for the Hydrogen atom.

Erwin Schrodinger

Ernest Rutherford

Erwin Schrodinger

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Schrodinger's atomic model uses quantum mechanics to show that each atom has a region of electron density where an electron is most likely to be present. 

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Periodic Trends

Periodic trends are properties of the elements of the periodic table which are based on the structure of the atoms that make up each element. The periodic table was designed with these trends in mind, so it is always helpful to refer to the periodic table when comparing periodic trends of certain atoms. 


Two important concepts to understand for periodic trends are Coulomb's Law and effective nuclear charge. Coulomb's Law states that two charged particles apply a force on each other, and this force is greatest when the charges are close together. In an atom, this means that the positively charged nucleus has less of an attraction on electrons that are farther away, and a greater attraction on electrons that are closer. Effective nuclear charge states that internal electrons essentially "block" the charge of the nucleus from reaching the electrons, so an atom with more internal electrons will have valence electrons that experience a lesser effective nuclear charge. An increase in the number of protons in an atom without changing the number of internal  electrons (like moving to the right across a period) increases the effective nuclear charge on the valence electrons.


Listed below are examples of periodic trends, how they change throughout the periodic table, and why the structure of the atom results in this trend.


Atomic size: the measure of how large an individual atom is

  • Increases down a group as more electron shells are added which become farther away from the nucleus
  • Increases to the left across a period since the increased number of protons in the nucleus moving to the right results in a greater effective nuclear charge on the electrons which pulls them closer to the nucleus 

Ionic radius: the size of an ion of an element

  • Cations (positive charge) become smaller since they lose an electron (or more) to ionize
  • Anions (negative charge) become larger since they gain an electron (or more) to ionize.

Electronegativity: the pull an atom applies on a shared electron in a compound

  • Increases up a group and to the right across a period (except for the noble gasses) as Fluorine is the most electronegative element

Electron affinity: the change in energy when an atom accepts another electron

  • Higher electron affinity means an atom wants an electron, lower electron affinity means an atom does not want an electron. 
  • Increases to the right as elements are closer to forming an octet or filling their valence shell of electrons
  • No clear pattern is seen with electron affinity going down a group as the value varies throughout the periodic table

Ionization energy: the force that needs to be exerted in order to remove an electron

  • Increases to the right across periods as more protons in the nucleus result in a higher effective nuclear charge on the electrons that attracts them with a stronger force and makes the electrons more difficult to remove
  • Increases up groups as electrons get farther away in higher energy levels and Coulomb's Law says that the attractive force is reduced. Atoms higher in their groups have fewer energy levels, so the electrons are closer to the nucleus and are harder to remove
  • Exceptions include Oxygen and Boron as half-filled or fully-filled orbitals are more stable than atoms where there is a single extra election present
  • Both first ionization energy and subsequent ionization energies can be measured, and increased levels of ionization energy are always larger. As more electrons are removed, there is less repulsion between the electrons so they are harder to take away from the atom.



Spectroscopy

Photoelectron Spectroscopy

Photoelectron Spectroscopy

Photoelectron Spectroscopy

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Photoelectron spectroscopy is a way to identify the electron configuration of an atom or ion based on the Photoelectric Effect (which Einstein won his Nobel Prize for) and ionization energy. The graph above is an example of the output of the PES process. On the horizontal axis, the ionization energy is plotted. On the vertical axis, the number of electrons is plotted. Each spike correlates to an electron orbital: 1s2, 2s2, 2p6, and 3s2, for example, as shown above. Since there are 12 electrons present, this graph corresponds with magnesium. Be careful, since the orientation of the horizontal axis can be swapped, so be sure to check the units and numbers. 

Mass Spectroscopy

Photoelectron Spectroscopy

Photoelectron Spectroscopy

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Mass Spectroscopy is used to determine the mass of  isotopes and their relative abundance. Any type of particle can be bombarded with electrons and shot through a charged, curved tube. The heavier  particles are deflected less, while the lighter particles are deflected more. The output allows us to determine the frequency of certain masses of particle which helps determine the most abundant isotope.

UV/Vis Spectroscopy

Photoelectron Spectroscopy

UV/Vis Spectroscopy

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UV/Vis Spectroscopy is an important step in the process to determine concentrations of colored solutions using ultraviolet or visible light. This process helps us determine the absorbtivity constant which is used in 

the Beer - Lambert Law, A=abc or Absorbtivity = absorbtivity constant * path length of light * concentration. 

Worksheet

Atomic Structure Worksheet and Answer Key (pdf)

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