Question

Explain why a graph of ionization energy versus atomic num ber (across a row) is not linear. Where are the exceptions? Why are there exceptions?

Short answer

The graph of ionization energy versus atomic number across a row is not linear because of variations in electron shielding, subshell types, and nuclear charge. Exceptions to the trend occur when transitioning between subshells (e.g., s to p, p to d, and d to f), as seen between groups 2 and 13, groups 5 and 6, and the lanthanides and actinides. These exceptions are due to poorer electron shielding in higher-energy subshells, making it easier to remove electrons and causing a decrease in ionization energy.

Step-by-step solution

Understand ionization energy and atomic number

Ionization energy is the energy required to remove an electron from a neutral atom in the gaseous state. Atomic number is the number of protons in the nucleus of an atom. As we move across a row (period) in the periodic table, the atomic number increases, meaning that the nucleus becomes more positively charged.

Consider the trend of ionization energy across a row

As we move across a row in the periodic table, ionization energy generally increases. This is because the increase in positive charge in the nucleus makes it more difficult to remove an electron from the atom. However, this trend is not linear, and there are exceptions to this pattern due to variations in electron shielding and subshell types.

Identify the exceptions

The exceptions to this trend occur at the transition from one subshell type to another (i.e., from s to p, p to d, and d to f subshells). These are found between groups 2 and 13, groups 5 and 6, and between the lanthanides and actinides. These elements show a lower ionization energy than expected based on the general trend.

Explain why exceptions occur

The exceptions to the ionization energy trend occur due to the differing electron shielding and subshell types within a row. When an electron occupies a higher-energy subshell, it is further from the nucleus, making it easier to remove. Additionally, electrons in the same subshell do not shield each other effectively, so the ionization energy tends to decrease when transitioning between subshells. In summary, the graph of ionization energy versus atomic number is not linear due to variations in electron shielding, subshell types, and nuclear charge. Exceptions to the trend occur when transitioning between subshells, where ionization energy tends to be lower due to poorer electron shielding.

Key concepts

Atomic Number

The atomic number, which can be thought of as an atom's identification label, is fundamental to understanding trends on the periodic table. It's not just a simple count of protons within an atom's nucleus, it also plays a pivotal role in the atom's chemical behavior.

As the atomic number increases, so does the nuclear charge, meaning that each proton added exerts stronger forces of attraction on the electrons. This can make the atoms smaller and pull the electrons closer to the nucleus, increasing the energy required to remove an electron. However, while this hints at a consistent upward trend in ionization energy, the relationship is influenced by multiple factors, creating a landscape with peaks and troughs rather than a smooth climb.

Electron Shielding

Electron shielding is akin to a game of tug-of-war between the positive charge of the nucleus and the negative charge of electrons. It's the effect where inner electrons repel outer electrons, reducing the effective nuclear charge felt by those outermost electrons.

The strength of this 'buffering' action of inner electrons impacts how tightly the outer electrons are held and therefore the ionization energy required to remove them. As you might expect, if there is more electron shielding, the outer electron will experience a weaker pull from the nucleus, thus lowering the ionization energy. This concept is crucial to understand why ionization energy does not simply rise with increasing atomic number but instead shows a more complex pattern.

Periodic Table Exceptions

While trends can be powerful predictors of chemical behavior, the periodic table is also dotted with exceptions that break the expected patterns. These deviations often occur due to subtle differences in electron configurations which can have outsized effects on an atom's chemistry.

For instance, when moving from group 2 elements (with a filled s subshell) to group 13 (where a p subshell starts filling), there's a notable dip in the ionization energy trend. This is because the p orbital electrons are less effectively shielded by the s orbital electrons, making them easier to remove. Similarly, within transition metals and between the lanthanides and actinides, inner d or f electrons don't shield outer electrons as effectively, leading to lower than expected ionization energies.

Subshell Electron Configuration

For an atom's electrons, location is everything. Their energy and the ease with which they can be removed (aka ionization energy) are influenced by which subshell they occupy. Subshells are like the neighborhood within an atom where electrons reside: they can be s, p, d, or f.

Electrons in s subshells are closer to the nucleus and are held more tightly, thus requiring more energy to be removed, while p, d, and f electrons are progressively further away, often leading to lower ionization energies. But there's a twist: electron configurations don't always fill neatly. For example, the first electron in a p subshell will have a lower ionization energy than the last electron in an s subshell, despite being in the same period. This nuance in electron occupancy explains why the graph of ionization energy versus atomic number zigs and zags, rather than follows a straight path.