Periodic Trends in the Properties of Atoms

One of the most fundamental prindples of chemistry is the periodic law, which states that 

The chemical and physical properties of elements are a periodic function of atomic number. 

This is, of course, the principle behind the structure of the periodic table. Elements within a given vertical group resemble one another chemically because chemical properties repeat themselves at regular intervals of 2, 8,18, or 32 elements. 

In this section we will consider how the periodic table can be used to correlate properties on an atomic scale. In particular, we will see how atomic radius, ionic radius, ionization energy, and electronegativity vary horizontally and vertically in the periodic table. 

The decrease in atomic radius moving across the periodic table can be explained in a similar manner. Consider, for example, the third period, where electrons are being added to the third principal energy level. The added electrons should be relatively poor shields for each other because they are all at about the same distance from the nucleus. Only the ten core electrons in inner, filled levels (n = 1, n = 2) are expected to shield the outer electrons from the nucleus. This means that the charge felt by an outer electron, called the effective nuclear charge, should increase steadily with atomic number as we move across the period. As effective nuclear charge increases, the outermost electrons are pulled in more tightly, and atomic radius decreases. 

It is possible to explain these trends in terms of the electron configurations of the corresponding atoms. Consider first the increase in radius observed as we move down the table, let us say among the alkali metals (Group 1). All these elements have a single s electron outside a filled level or filled p sublevel. Electrons in these inner levels are much closer to the nucleus than the outer s electron and hence effectively shield it from the positive charge of the nucleus. To a first approximation, each inner electron cancels the charge of one pro- 

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Explain the structure of the periodic table and periodic trends in the properties of atoms... 

The first ionization energies and electronegativities of the typical elements show overall increases across the Periodic Table (Figures 8.9 and 4.4). By contrast, the single-bond covalent radii show a decrease. All three trends reveal the importance of the increase in the nuclear charge as successive protons are added to the nucleus of the atoms. Such an increase causes the outer electrons to be bound more and more tightly across the row. They are trends in the properties of atoms, and as such provide possible explanations of other trends in physical and chemical properties, which we examine in the following Sections. 

There are similar, but smaller, trends in the properties of elements in a column (a family) of the periodic table. Though the elements in a family display similar chemistry, there are important and interesting differences as well. Many of these differences are explainable in terms of atomic size. 

Niobium (formerly called columbium) and tantalum are Transition Metals having a considerable affinity for oxygen donor groups they are thus called oxophilic see Oxophilic Character). They occur as mixed-metal oxides such as columbites (Fe/Mn)(Nb/Ta)206 and pyrochlore NaCaNb206p. Their discovery in minerals extends back to the beginning of the nineteenth century, when they were believed to be identical and called tantalum. Rose showed that at least two different elements were involved in the minerals, and named the second one niobium. Their separation was resolved around 1866, especially by Marignac. These metals often display similar chemical behavior as a result of nearly identical atomic radii (1.47 A) due to the lanthanide contraction see Periodic Table Trends in the Properties of the Elements)... 

The concept of an atom s oxidation state see Oxidation Number) can provide fundamental information about the stmcture and reactivity of the compound in which the atom is found. In fact, it can be argued that oxidation states provided the basis for Medeleev s initial organization of the periodic table. For the main group elements, the relative stability of lower oxidation states within a given group increases as the atomic number increases. This trend in the periodic table see Periodic Table Trends in the Properties of the Elements) is generally attributable to the presence of an inert s pair see Inert Pair Effect) caused by relativistic effects see Relativistic Effects). 

Chapter 4, Atoms and Elements, introduces elements and atoms and the periodic table. The names and symbols of element 114, Herovium, FI, and 116, Livermorium, Lv, have been added to update the periodic table. Atomic numbers and mass number are determined for isotopes. Atomic mass is calculated using the masses of the naturally occurring isotopes and their abundances. Trends in the properties of elements are discussed, including atomic size, electron-dot symbols, ionization energy, and metallic character. 

We see that, no matter what type of bonding situation is considered, there is a trend in size moving downward in the periodic table. The alkaline earth atoms become larger in the sequence Be < Mg < Ca < Sr < Ba. These atomic sizes provide a basis for explaining trends in many properties of the alkaline earth elements and their compounds. 

Let s begin by surveying some of the key physical and chemical properties of the transition-metal elements and interpreting trends in those properties using the quantum theory of atomic structure developed in Chapter 5. We focus initially on the fourth-period elements, also called the first transition series (those from scandium through zinc in which the 3d shell is progressively filled). Then we discuss the periodic trends in the melting points and atomic radii of the second and third transition series elements. 

Many of the chemical properties of the elements can be understood IN TERMS OF THEIR ELECTRON CONFIGURATIONS. BECAUSE ELECTRONS FILL ATOMIC ORBITALS IN A FAIRLY REGULAR FASHION, IT IS NOT SURPRISING THAT ELEMENTS WITH SIMILAR ELECTRON CONFIGURATIONS, SUCH AS SODIUM AND POTASSIUM, BEHAVE SIMILARLY IN MANY RESPECTS AND THAT, IN GENERAL, THE PROPERTIES OF THE ELEMENTS EXHIBIT OBSERVABLE TRENDS. CHEMISTS IN THE NINETEENTH CENTURY RECOGNIZED PERIODIC TRENDS IN THE PHYSICAL AND CHEMICAL PROPERTIES OF ELEMENTS LONG BEFORE QUANTUM THEORY CAME ONTO THE... 

Next, we examine the periodic trends in physical properties such as the size of atoms and ions in terms of effective nuclear charge. (8.3)... 

Chemists in the nineteenth century recognized periodic trends in the physical and chemical properties of elements, long before quantum theory came onto the scene. Although these chemists were not aware of the existence of electrons and protons, their efforts to systematize the chemistry of the elements were remarkably successful. Their main sources of information were the atomic masses of the elements and other known physical and chemical properties. Modem quantum theory allows us to understand these periodic trends in terms of the ways in which the electrons are distributed among the atomic orbitals of an atom. 

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