Ionisation energy is a measure of how easily an atom loses its outer shell electron. If an atom loses this electron, it forms a positive ion. Some atoms form ions more readily than others and only with certain charges. To understand this better, let's recap ionisation energy. It is the energy required to remove one mole of outer shell electrons from one mole of a gaseous species. The lower the ionisation energy, the more readily an atom loses its electron to form a positive ion. Factors affecting ionisation energy are nuclear charge, distance from the nucleus, and shielding. The more protons an atom has, the stronger its nuclear charge, and the higher the ionisation energy. Electrons in shells with a higher quantum number are further from the nucleus, so the attraction is weaker, and the ionisation energy is lower. Inner shell electrons shield the outer shell electron, reducing the overall attraction it feels.
Ionisation energy is a topic that shows trends in the periodic table. These trends repeat across each period, which means that the pattern shown in period 2 will repeat in period 3. There are two important trends to remember. Firstly, ionisation energy increases across a period. Secondly, it decreases down a group.
One of the main reasons why ionisation energy increases across a period is due to an increase in nuclear charge. This means that there is an increase in the attraction between the nucleus and the outermost electron. For example, consider carbon and boron. Carbon has six protons, while boron has five. Because carbon's nuclear charge is larger, its first ionisation energy is higher than boron's. You can observe this by looking at the number of protons in the nucleus of each element. Carbon has one more proton than boron.
When it comes to the trend of ionisation energy decreasing down a group, there are several factors at play. Firstly, the distance between the nucleus and the outermost electron increases down a group. Additionally, the outer shell electron is shielded by more shells of inner electrons. This reduces the attraction between the electron and the nucleus, which negates the impact of increased nuclear charge. As a result, it becomes easier to remove an electron from the outer shell down a group.
The graph you mentioned does indeed follow the trend of increasing ionisation energy across period 3. However, there are some exceptions to this trend that can be explained by looking more closely at electron configuration.
For example, the first ionisation energy of magnesium is lower than that of aluminium, despite the fact that magnesium has a greater nuclear charge. This is because magnesium has a full outer shell (2 electrons in its valence shell), which makes it more stable and therefore easier to remove an electron. Similarly, the first ionisation energy of sulfur is lower than that of phosphorus, despite the fact that sulfur has a greater nuclear charge. This is because sulfur has one electron in its valence shell that is paired with another electron, which makes it more stable and therefore easier to remove an electron.
In general, the electron configuration of an element plays a significant role in determining its ionisation energy. Elements with full or half-filled valence shells are more stable and therefore have lower ionisation energies. Conversely, elements with partially filled valence shells are less stable and therefore have higher ionisation energies.
Elements in Group 6 in the periodic table have lower first ionisation energies than those in group 5, despite their increased nuclear charge. This can be explained by looking at their electron configuration.
For example, nitrogen (with seven electrons) has the structure while oxygen (with eight electrons) has the structure . Their electron configurations are shown below.
The electron configurations of nitrogen and oxygen. You’ll notice that nitrogen only has three electrons in the 2p subshell. According to Hund’s rule, electrons within a subshell will prefer to fill empty orbitals, and so there is a single electron located in each of 2p’s three orbitals. However, oxygen has four electrons in the 2p subshell. This means that one orbital must contain two electrons, which repel each other quite strongly. The electron-electron repulsion means that the outermost electron is easier to remove - it is already partially repelled by the other electron in its orbital. For a reminder on Hund’s rule, see Electron Configuration.
Group 3 elements have lower first ionisation energies than Group 2 elements, despite their increased nuclear charges. This is also due to their electron configuration. Boron's outermost electron is located in the 2p subshell, which is of a slightly higher energy level than the 2s subshell in which beryllium's outermost electron is located. The 2p subshell is also slightly further away from the nucleus than the 2s subshell, which reduces the attraction between the nucleus and the outermost electron and makes it easier to lose, thereby reducing the ionisation energy.
You are absolutely right. The successive ionisation energies of sodium show a large jump between the first and second ionisation energies, indicating that sodium tends to form ions with a charge of +1, unless supplied with a significant amount of energy. This is because the outermost electron in sodium is in the 3s subshell, and once it is removed, the electron configuration becomes that of a noble gas, which is energetically more stable. Similarly, the large jump between the third and fourth ionisation energies of aluminium indicates that it has a full outer shell, and therefore, belongs to Group 3. Overall, trends in ionisation energies show periodicity and can be used to determine an element's group based on its electron configuration.
What is the trend of ionisation energy?
Ionisation energy generally increases across a period. This is because nuclear charge increases. Ionisation energy also decreases down a group. This is because atomic radius increases and the outer electrons are shielded from the nucleus by inner electron shells.
What is the trend between successive ionisation energies?
Successive ionisation energies increase as you remove more and more electrons from a species. This is because you are removing negative electrons from an increasingly positively charged ion, which requires more energy.
How does ionisation energy work?
Ionisation energy is the energy needed to remove one mole of electrons from one mole of gaseous atoms under standard conditions. It takes energy to remove electrons because they are strongly attracted to the atom's positive nucleus, which is why values for ionisation energy are positive. Factors affecting ionisation energy include nuclear charge, atomic radius, and shielding from inner electron shells.
How do you determine ionisation energy?
Ionisation energy is determined by firing electrons at a gaseous sample. When the electrons have enough energy, they can knock off the least tightly-bound electron from each of the atoms in the sample. This generates a current which is proportional to the energy needed to ionise the atoms - their ionisation energy.
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