7.5 Periodic Trends | High School Chemistry

This paragraph introduces the topic of periodic trends, focusing on atomic radius, ionization energy, and electron affinity. It highlights that francium has the largest atomic radius and explains the trends in atomic radius both vertically and horizontally across the periodic table. The concept of effective nuclear charge is introduced to explain these trends, emphasizing how it affects the attraction of valence electrons to the nucleus. The paragraph also humorously relates the atomic radius to eating habits in France, suggesting that francium's large size is akin to overindulgence in French cuisine.

The paragraph delves deeper into the concept of effective nuclear charge, explaining how it is calculated and its role in determining atomic radius. It illustrates the difference between core and valence electrons and how the former shield the latter from the nucleus's positive charge. Examples using sodium, magnesium, and aluminum demonstrate how the effective nuclear charge increases from left to right across a period, leading to a decrease in atomic radius. The paragraph also discusses the application of atomic radius in determining bond lengths in covalent and ionic bonds.

This paragraph discusses ionic radii, differentiating between cations and anions, and how they compare to atomic radii. It explains that metals tend to form cations by losing electrons, leading to a significant decrease in size, while non-metals form anions by gaining electrons, resulting in an increase in size. The paragraph provides examples of potassium, calcium, sulfur, and chlorine to illustrate these changes. It also touches on the limitations of comparing different ions and the importance of isoelectronic series, where elements have the same electron configuration.

The paragraph introduces ionization energy as the energy required to remove an electron from an atom and electron affinity as the energy change associated with gaining an electron. It explains that ionization energy typically increases as you move up a group or across a period due to smaller atomic size and greater effective nuclear charge. Conversely, electron affinity tends to increase (become more negative) as you move left to right across the periodic table, with chlorine being the most electronegative element. Exceptions to these trends are noted, particularly for noble gases, which neither readily lose nor gain electrons.

This paragraph explores the exceptions to the general trends of ionization energy and electron affinity. It explains that elements with a filled s subshell (like beryllium) or a half-filled p subshell (like nitrogen) are more stable and thus have higher ionization energies. The paragraph also discusses how elements like oxygen, which are one electron past a half-filled shell, have lower ionization energies. Electron affinity is similarly affected, with filled or half-filled subshells leading to lower electron affinities. The paragraph emphasizes the importance of understanding these exceptions for high school chemistry.

The paragraph discusses the concept of successive ionization energy, explaining that it becomes increasingly difficult to remove additional electrons from an atom as the atom becomes smaller and more positively charged. This is illustrated with the example of boron, which has a significant jump in ionization energy after losing its third electron. The paragraph also addresses electron affinities, noting that while most elements gain one electron readily, gaining a second electron is generally unfavorable due to repulsive forces between like charges. Exceptions to this rule, such as oxygen gaining two electrons to achieve a noble gas configuration, are highlighted.

The final paragraph summarizes the key points covered in the lesson, including atomic radius, ionization energy, and electron affinity. It emphasizes the importance of understanding both the general trends and the exceptions, as these are often the focus of high school chemistry exams. The paragraph also provides advice on how to study, recommending the use of a study guide and practice problems to solidify understanding. The instructor encourages students to download the study guide and work through problems to become proficient in applying these concepts.