Predict the relationship between \(\mathrm{IE}_{1}\) and atomic number by making a rough graph of IE \(_{1}\) vs. atomic number. DO NOT PROCEED TO THE NEXT PAGE UNTIL YOU HAVE COMPLETED THIS GRAPH.

The atomic number of an element is a fundamental concept in chemistry that represents the number of protons in the nucleus of an atom. It's given the symbol 'Z' and is crucial for defining the identity of an element. Each element on the periodic table has a unique atomic number, which also indirectly informs us about the number of electrons in a neutral atom.

Since electrons are negatively charged and protons are positively charged, an increase in the atomic number means a stronger positive charge in the nucleus. This results in a greater attraction to the negatively charged electrons, which ties directly into the concept of ionization energy. Higher atomic numbers generally correlate with higher ionization energies because more energy is required to remove an electron from an increasingly positively charged nucleus.

The periodic table is not just a list of elements; it's a matrix of information that conveys trends in physical and chemical properties. One such property is ionization energy (IE), the energy required to remove an electron from a gaseous atom or ion. IE trends are observed by moving across periods (from left to right) and down groups (from top to bottom).

• Across a Period: Ionization energy generally increases as you move from left to right across a period due to the increase in atomic number and consequently, a stronger nuclear charge attracting the electrons more firmly.
• Down a Group: As you move down a group, ionization energy decreases. This is because atoms have more electron shells, so the outer electrons are further from the nucleus and more shielded by inner electrons, decreasing the effective nuclear charge.

It's essential to recognize that while the pattern is consistent, it isn’t perfectly linear due to the underlying electron configurations of the elements.

Electron configuration describes how electrons are distributed in an atom's orbitals. The arrangement follows a set of defined rules, such as the Aufbau principle, Pauli exclusion principle, and Hund’s rule, to minimize the energy of the atom.

The subshells and orbitals in which the electrons reside have different energy levels. When electrons occupy these different orbitals, the ionization energy can vary. For example, elements with a half-full or full subshell have a particularly stable configuration, which impacts their ionization energies.

Electron configuration also helps explain anomalies in IE trends. For instance, elements in the transition metals have electrons entering an underlying 'd' subshell, which can lead to less pronounced increases in IE. Similarly, the slight drop in IE between groups 2 and 13 reflects the shift from a filled 's' orbital to the start of filling a 'p' orbital, which requires less energy to remove the first 'p' electron.

Thus, understanding electron configurations is crucial when predicting and explaining the ionization energy of elements and contributes to the nuanced understanding of the periodic trends.