Periodic Properties for the DAT
/Learn key DAT concepts about periodic properties, plus practice questions and answers
Learn key DAT concepts about periodic properties
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Part 1: Introduction to periodic properties
Periodic properties refer to trends and characteristics that vary systematically across the rows and columns of the periodic table. The periodic table is organized by increasing atomic number, and as you progress across a period or down a group, various properties of the elements exhibit predictable patterns. These properties include atomic size, ionization energy, electron affinity, electronegativity, and metallic character. The study of these properties provides crucial insights into the fundamental nature of elements, making it a significant focus for students preparing for the DAT. While you study this guide, pay attention to the terms in bold. Also, test your understanding with practice questions and answers at the end of the guide.
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Part 2: Groups
Understanding trends in the periodic table is immensely beneficial on the DAT. During the exam, you'll be provided with a periodic table and asked to use it for comparing element properties based on your knowledge of periodic trends. One fundamental trend is observed between groups, which are the 18 vertical columns in the periodic table. Each group has elements with a varying number of valence electrons. For instance, group 2 contains elements with 2 valence electrons, while group 13 encompasses those with 3 valence electrons, and so forth. Elements within a group share similar atomic arrangements and properties.
Noble gases in group 18 are characterized by a complete outer shell and exhibit low reactivity as you move down the group. Conversely, alkali metals in group 1, with a single valence electron, are highly reactive. Alkaline earth metals in group 2 also exhibit a high reactivity. Generally, elements on the far left of the periodic table are more reactive than those in the center. Group 17 is the halogens, which are highly reactive diatomic elements. Chalcogens in group 16, such as oxygen, also exhibit distinct properties. All of these groups are representative elements. Representative elements, also known as main-group elements, are any elements that are found in groups 1, 2, and 13-18. The remaining groups, 3-12, are known as transition metals. These elements can form variable oxidation states, complex ions, and colorful compounds. The metallic character of elements decreases as you go up and right on the periodic table.
FIGURE 1: LABELED PERIODIC TABLE OF THE ELEMENTS
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Part 3: Periodic trends
The Periodic Table is thoughtfully arranged to catalog various trends, delineating alterations in the characteristics of each element. A comprehensive understanding of these periodic trends and the ability to adeptly apply them on Test Day are crucial. Electronegativity, signifying the nucleus’s attractive force for an electron, is influenced by the outermost electron shell’s stability. This outermost shell is known as the valence shell, and electrons in this shell are called valence electrons. Elements with nearly full valence shells exhibit heightened electronegativity compared to those with one or two electrons in the valence shell. Energetically, atoms with nearly full electron shells gain stability by attracting a last electron, while those with only a few valence electrons find stability by releasing them. Consequently, electronegativity ascends upwards and across to the right of the periodic table. For example, a halogen such as Fluorine (F) is highly electronegative, contrasting with an alkaline earth metal like barium (Ba), which is less electronegative. The noble gases are an exception, as they have nearly zero electronegativity due to their exceptional stability.
Ionization energy is the energy needed to remove an electron from the valence shell. It follows a similar trend as electronegativity. increasing upwards and to the right. Chlorine, for example, surpasses potassium in ionization energy due to its nearly full valence shell, making it energetically favorable to gain one last electron. Notably, the second ionization energy is considerably higher than the first. This means that removing a second electron demands much more energy than removing only one, as it is harder to ionize a charged particle compared to a neutral one.
Electron affinity denotes the energy released when an electron attracts another electron, and it is intricately linked with ionization energy and electronegativity. The electron affinity trend ascends upwards and extends across to the right of the periodic table. This phenomenon is rooted in the escalating attractive forces between the nucleus and valence electrons as the distance to the valence shell decreases. The increase in protons within the nucleus, while maintaining a constant distance, intensifies the electrostatic attraction.
Another pivotal trend is the atomic size, characterized by a decrease upwards and to the right of the periodic table. This diminution is primarily attributed to the concept of effective nuclear charge.
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