Question

Explain the general trend in ionization energy as you go from left to right along Periods \(1-5\) of the periodic table. (Chapter 6 )

Short answer

As we move from left to right along Periods 1-5 in the periodic table, ionization energy generally increases. This is due to the increase in effective nuclear charge as the number of protons increases, causing the electrons in the outer shell to experience a stronger attraction to the nucleus. Meanwhile, electron shielding remains constant and atomic radius decreases, further contributing to the increase in ionization energy. However, there are occasional exceptions to this trend due to electron configurations and electron shielding.

Step-by-step solution

Define ionization energy

Ionization energy is the energy required to remove an electron from an atom in the gaseous state. The first ionization energy is the energy required to remove the first electron, whereas the second ionization energy is the energy required to remove the second electron, and so on. In general, ionization energy increases as we move across a period and decreases as we move down a group in the periodic table.

Understand the factors affecting ionization energy

There are three main factors that affect ionization energy: atomic radius, effective nuclear charge, and electron shielding. 1. Atomic radius: This is the distance from the nucleus to the outermost electron shell. As the atomic radius decreases, the electrons are held more tightly to the nucleus, and ionization energy increases. 2. Effective nuclear charge: This is the net positive charge experienced by an electron in an atom. As the nuclear charge increases, the attractive force between the nucleus and the electrons increases, leading to a higher ionization energy. 3. Electron shielding: This occurs when inner electron shells shield outer electrons from the full attractive force of the nucleus. Greater shielding results in lower ionization energy as outer electrons are not as strongly attracted to the nucleus.

Explain the general trend across a period (left to right)

As we move from left to right across Periods 1-5, the number of protons in the nucleus increases, which results in a more positive nuclear charge. Additionally, the number of electrons in the same energy level also increases. However, the increase in positive nuclear charge outweighs the increase in electron-electron repulsion, causing the effective nuclear charge to increase. Consequently, electrons in the outer shell will experience a stronger attraction to the nucleus and will be harder to remove. Since the effective nuclear charge increases while electron shielding remains constant and atomic radius decreases, ionization energy generally increases from left to right across a period in the periodic table.

Exceptions to the general trend

Sometimes, there are exceptions to the general trend of ionization energy increasing from left to right across a period, due to electron configurations and electron shielding. For instance: 1. The ionization energy of oxygen is slightly lower than nitrogen; this is because oxygen's electron configuration has two half-filled p orbitals that create a slightly higher repulsion, making it easier to remove an electron compared to nitrogen. 2. The ionization energy of aluminum is lower than magnesium due to an extra electron in its 3p orbital, which experiences electron shielding, making it easier to remove an electron compared to magnesium with a fully filled 3s orbital. These exceptions are observed infrequently and the general pattern of increasing ionization energy across periods still holds.

Key concepts

Periodic Table Trends

The periodic table is not just a list of elements, but a map filled with trends that help us understand atomic behavior. One of the major trends is the ionization energy, which generally increases as we move from left to right across a period. This trend occurs because more protons are added to the nucleus, increasing the positive charge.
As the nuclear charge increases, electrons are pulled more strongly towards the nucleus. Even though new electrons are added to the same energy level, the increased nuclear attraction outweighs the added electron-electron repulsion.

• The effective nuclear charge increases.
• Atomic radius decreases.
• Electron shielding remains constant.

These factors combine to make it harder to remove an electron, thus increasing ionization energy.

Atomic Radius

Atomic radius is a key factor that influences ionization energy. It is essentially the size of an atom, measured from the center of the nucleus to the outermost electron shell. As you move across a period from left to right, the atomic radius decreases.
This shrinkage occurs because the additional protons in the nucleus create a stronger pull on the electrons, drawing them closer and effectively reducing the overall size of the atom. With a smaller atomic radius, the electrons are held more tightly to the nucleus, making it more difficult to remove an electron.
Thus, a smaller atomic radius is associated with higher ionization energy.

Effective Nuclear Charge

Effective nuclear charge ( Z_{eff} ) is a critical concept in understanding ionization energy. It refers to the net positive charge experienced by an electron in an atom. As you move from left to right across a period, the effective nuclear charge increases.
Why? Because each element has more protons boosting the nucleus' positive charge while electrons are added to the same shell. This increased nuclear charge means each electron feels a stronger attraction to the nucleus, pulling them in tighter.
This tightening of electron orbits results in a higher ionization energy since it takes more energy to remove an electron that is tightly bound to the nucleus.

Electron Shielding

Electron shielding, also known as screening, is what happens when inner shell electrons block or "shield" the outer electrons from the full attractive force of the nucleus. As a result, the outer electrons feel less of the nuclear charge.
As you move across a period, electron shielding remains constant because no new electron shells are being added; electrons are added to the same principal energy level.

• Hence, changes in shielding do not change the periodic trend of ionization energy.
• This consistency is why the effective nuclear charge has a more pronounced impact than shielding as you move across a period.

The lack of increased shielding across a period allows the full effect of increased nuclear charge to take hold, leading to a general increase in ionization energy across the period.