From their positions in the periodic table, arrange the atoms in each of the following series in order of increasing electronegativity: (a) \(\mathrm{As}, \mathrm{H}, \mathrm{N}, \mathrm{P}, \mathrm{Sb}\) (b) Cl, H, P, S, Si (c) \(\mathrm{Br}, \mathrm{Cl}, \mathrm{Ge}, \mathrm{H}, \mathrm{Sr}\) (d) \(\mathrm{Ca}, \mathrm{H}, \mathrm{K}, \mathrm{N}, \mathrm{Si}\) (e) Cl, Cs, Ge, H, Srv

The periodic table is a tabular arrangement of elements, organized on the basis of their atomic numbers, electron configurations, and recurring chemical properties. Elements are presented in order of increasing atomic number in rows called periods. Each column is known as a group and contains elements with similar properties due to their similar valence electron configurations. The layout of the periodic table highlights the regular trend in properties, including electronegativity, which is pivotal for understanding chemistry.

The organization of the periodic table is such that elements with the highest electronegativities are found on the right side, excluding the noble gases, and those with lower electronegativities are found to the left. Recognizing the structure and layout of the periodic table thus allows students to predict and understand the trends of various chemical properties.

Chemical properties of elements are characteristics that can be measured or observed only when atoms undergo a chemical reaction and change into different substances. These properties include reactivity with other chemicals, acid-base properties, combustion, and, notably, electronegativity. Understanding these properties can give insight into how an element will behave in different chemical circumstances.

For example, chemical reactivity is influenced by an element’s electronegativity. Elements with high electronegativity, such as fluorine, are more likely to attract electrons and form negative ions, whereas those with low electronegativity, such as metals, tend to lose electrons and form positive ions. This knowledge is invaluable when trying to grasp complex chemical reactions and predict product formation.

Atomic electronegativity is a measure of an atom's ability to attract and hold onto the electrons within a chemical bond. It varies across the periodic table, affected by both the atomic number and the distance of the valence electrons from the charged nucleus. Generally, electronegativity increases from left to right within a period and decreases from top to bottom within a group. This is due to the increase in positive charge of the nucleus moving across a period, which draws electrons more strongly, and the increase in atomic radii moving down a group, which makes it harder for the nucleus to attract electrons from a bigger distance.

Different scales, such as the Pauling scale, assign values to the relative electronegativities of different elements, assisting students in learning to predict the type of bond that will form between atoms and the polarity of molecules.

Electron affinity refers to the amount of energy released when an electron is added to a neutral atom to form a negative ion. This property is related to electronegativity since elements with higher electronegativity typically have higher electron affinity. Elements with high electron affinities tend to be non-metals, as they have a greater tendency to gain electrons, becoming anions. In contrast, metals usually have lower electron affinities, as they are more likely to lose electrons and form cations.

To understand this concept, students should know that electron affinity can vary greatly among different elements, and it can also provide insight into the relative reactivity of elements, especially halogens which have very high electron affinities. This characteristic is crucial in predicting how atoms will interact during chemical reactions.