Why are ionization energies always positive

As you move closer to the nucleus 2nd and 3rd ionization energies , it becomes harder more energy to remove them because they are held more tightly by the protons. The more positive the ionization energy and the more negative the electron affinity, the greater the electronegativity of an element. Ionization is defined as the amount of energy require to remove an electron from an atom.

Electrons are attracted to the nucleus, so the energy required to remove these electrons is always positive. Electron affinity however, is defined as the change in the atom's energy when an electron is added. Most electrons attract extract extra electrons and therefore the atom gives off energy when the electron is gained and becomes negative.

Some atoms though, do not attract any further electrons and need a positive energy given to the atom in order to add an electron. Francium in group 1 has the lowest first ionization energy. No fire is not an ionization energy. Lithium has an ionization energy of 5. He has the highest 1st ionization energy. Li has the highest 2nd ionization energy.

Be has the highest ionization energy. Oxygen's ionization energy is Tins ionization energy is Neon's ionization energy is Sulfer's ionization energy is Argon's ionization energy is Indium's ionization energy is Mercury's ionization energy is Helium has the highest first ionization energy and francium has the lowest first ionization energy.

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Data taken from John Emsley, The Elements , 3rd edition. Oxford: Clarendon Press, The ionization energy of an atom is the amount of energy that is required to remove an electron from a mole of atoms in the gas phase:. Ionization energies are always positive numbers, because energy must be supplied an endothermic energy change to separate electrons from atoms.

The second ionization energy is always larger than the first ionization energy, because it requires even more energy to remove an electron from a cation than it is from a neutral atom. The first ionization energy varies in a predictable way across the periodic table. The ionization energy decreases from top to bottom in groups , and increases from left to right across a period. Thus, helium has the largest first ionization energy, while francium has one of the lowest.

There are some "fluctuations" in these general trends. For instance, the first ionization decreases from beryllium to boron, and from magnesium to aluminum, as electrons from the p-block start to come into play. In the case of boron, which has an electron configuration of 1s 2 2s 2 2p 1 , the 2s electrons shield the higher-energy 2p electron from the nucleus, making it slightly easier to remove. A similar effect occurs in aluminum, which has an electron configuration of 1s 2 2s 2 2p 6 3s 2 3p 1.

Even though oxygen is to the right of nitrogen in period 2, its first ionization energy is slightly lower than that of nitrogen. Nitrogen has an electron configuration of 1s 2 2s 2 2p 3 , which puts one electron in each p orbital, making it a half-filled set of orbitals:. Half-filled sets of p orbitals are slightly more stable than those with 2 or 4 electrons, which makes it slightly harder to ionize a nitrogen atom. Oxygen has an electron configuration of 1s 2 2s 2 2p 4 , which puts another electron in one p orbital; since this is one electron away from being half-filled, it is slightly easier to remove this additional electron:.

Uub n. H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn The same pattern can be seen in the ionization energies of aluminum. The first ionization energy of aluminum is smaller than magnesium. The second ionization energy of aluminum is larger than the first, and the third ionization energy is even larger. Predict the group in the periodic table in which an element with the following ionization energies would most likely be found.

Click here to check your answer to Practice Problem 5. Use the trends in the ionization energies of the elements to explain the following observations. Click here to check your answer to Practice Problem 6.

Ionization energies measure the tendency of a neutral atom to resist the loss of electrons. It takes a considerable amount of energy, for example, to remove an electron from a neutral fluorine atom to form a positively charged ion. The electron affinity of an element is the energy given off when a neutral atom in the gas phase gains an extra electron to form a negatively charged ion.

A fluorine atom in the gas phase, for example, gives off energy when it gains an electron to form a fluoride ion. Electron affinities are more difficult to measure than ionization energies and are usually known to fewer significant figures. The electron affinities of the main group elements are shown in the figure below.

Several patterns can be found in these data. Electron affinities generally become smaller as we go down a column of the periodic table for two reasons. First, the electron being added to the atom is placed in larger orbitals, where it spends less time near the nucleus of the atom. Second, the number of electrons on an atom increases as we go down a column, so the force of repulsion between the electron being added and the electrons already present on a neutral atom becomes larger.

Electron affinity data are complicated by the fact that the repulsion between the electron being added to the atom and the electrons already present on the atom depends on the volume of the atom. As a result, these elements have a smaller electron affinity than the elements below them in these columns as shown in the figure below. From that point on, however, the electron affinities decrease as we continue down these columns.

At first glance, there appears to be no pattern in electron affinity across a row of the periodic table, as shown in the figure below. When these data are listed along with the electron configurations of these elements, however, they make sense. These data can be explained by noting that electron affinities are much smaller than ionization energies. As a result, elements such as helium, beryllium, nitrogen, and neon, which have unusually stable electron configurations, have such small affinities for extra electrons that no energy is given off when a neutral atom of these elements picks up an electron.

These configurations are so stable that it actually takes energy to force one of these elements to pick up an extra electron to form a negative ion.