Why is measuring the size of an atom difficult

Atomic size is defined in several different ways and these different definitions often produce some variations in the measurement of atomic sizes. Because it is so difficult to measure atomic size from the nucleus to the outermost edge of the electron cloud, chemists use other approaches to get consistent measurements of atomic sizes.

One way that chemists define atomic size is by using the atomic radius. The atomic radius is one-half the distance between the centers of a homonuclear diatomic molecule a diatomic molecule means a molecule made of exactly two atoms and homonuclear means both atoms are the same element.

The figure below represents a visualization of the atomic size definition. How do we measure the size of the atom? Ernest Rutherford is famous for his experiments bombarding gold foil with alpha particles. The gold foil experiment by Rutherford, first done in , is of particular interest to us in this unit because it was this experiment that first gave science an approximate measurement for the size of the atom.

Let's now look at the atomic radii or the size of the atom from the top of a family or group to the bottom. Take, for example, the Group 1 metals. Each atom in this family and all other main group families has the same number of electrons in the outer energy level as all the other atoms of that family.

Each row period in the periodic table represents another added energy level. When we first learned about principal energy levels, we learned that each new energy level was larger than the one before. Energy level 2 is larger than energy level 1, energy level 3 is larger than energy level 2, and so on.

Therefore, as we move down the Periodic Table from period to period, each successive period represents the addition of a larger energy level. It becomes apparent that as we move downward through a family of elements, that each new atom has added another energy level and will, therefore, be larger.

One other contributing factor to atomic size is something called the shielding effect. The protons in the nucleus attract the valence electrons in the outer energy level because of opposite electrostatic charges. The strength of this attraction depends on the size of the charges, the distance between the charges, and the number of electrons in-between the nucleus and the valence electrons.

The core electrons shield the valence electrons from the nucleus. The presence of the core electrons weakens the attraction between the nucleus and the valence electrons. This weakening of the attraction is called the shielding effect. The amount of shielding depends on the number of electrons between the nucleus and the valence electrons. The strength with which the nucleus pulls on the valence electrons can pull the valence shell in tighter when the attraction is strong and not so tight when the attraction is weakened.

The more shielding that occurs, the further the valence shell can spread out. When we compare an atom of sodium to one of cesium, we notice that the number of protons increases as well as the number of energy levels occupied by electrons. There are also many more electrons between the outer electron and the nucleus, thereby shielding the attraction of the nucleus.

The outermost electron, 6 s 1 , therefore, is held very loosely. In other words, because of shielding, the nucleus has less control over this 6 s 1 electron than it does over a 3 s 1 electron. The result of all of this is that the atom's size will be larger. Table What is true for the Group 1 metals is true for all of the groups, or families, across the periodic table. As you move downward in the periodic table through a family group, the size of the atoms increases.

For instance, the atoms that are the largest in the halogen family are bromine and iodine since astatine is radioactive and only exists for short periods of time, we won't include it in the discussion. As noted earlier for the main group metals, the outermost energy level in the electron configuration is indicated by the period number. Understanding the importance of education with e-learning transforming.

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Apna phone number register karein. Published by Katy Edger Modified over 6 years ago. What is ionization energy? Note: Figure 16 on p. The picture below shows the ionization of sodium. Both reactions require some amount of ionization energy. Why is measuring the size of an atom difficult? The electron cloud has no clear edge. Also, the size of the electron cloud can change based on the chemical and physical environment. Notes: Figure 19 on p. What can you tell about an atom that has high electronegativity?

Notes: Figure 22 and 23 show the trends for electronegativity. In the picture below, the Oxygen O is more electronegative pulls harder on electrons than Hydrogen H. What periodic trends exist for ionization energy?

Ionization energy increases as you move across a period from left to right. Notes: Figure 17 and 18 show the ionization energy trends. Compare the chart below to Figure Do they show the same trend?

What periodic trends exist for electronegativity? Electronegativity increases as you move across a period from left to right.

Notes: Figure 22 and 23 show the ionization energy trends. Compare to the ionization energy trend. Explain why the noble gases have high ionization energies. This is a stable state and it takes a lot of energy to make an atom unstable.