Principal energy levels

First Quantum Number, n Principal Energy Levels... 

Each principal energy level includes one or more sublevels. The sublevels are denoted by the second quantum number, . As we will see later, the general shape of the electron cloud associated with an electron is determined by . Larger values of produce more complex shapes. The quantum numbers n and are related can take on any integral value starting with 0 and going up to a maximum of (n — 1). That is,... 

If n = 1, there is only one possible value of —namely 0. This means that, in the first principal level, there is only one sublevel, for which = 0. If n = 2, two values of are possible, 0 and 1. In other words, there are two sublevels ( = 0 and = 1) within the second principal energy level. In the same way,... 

From Figure 6.8 it is possible to predict the electron configurations of atoms of elements with atomic numbers 1 through 36. Because an s sublevel can hold only two electrons, the Is is filled at helium (Is2). With lithium (Z = 3), the third electron has to enter a new sublevel This is the 2s, the lowest sublevel of the second principal energy level. Lithium has one electron in this sublevel (ls s1)- With beryllium (Z = 4), the 2s sublevel is filled (ls22s2). The next six elements fill the 2p sublevel. Their electron configurations are... 

After argon, an overlap of principal energy levels occurs. The next electron enters the lowest sublevel of the fourth principal level (4s) instead of the highest sublevel of the third principal level (3d). Potassium (Z = 19) has one electron in the 4s sublevel calcium (Z = 20) fills it with two electrons ... 

The atoms of elements in a group of the periodic table have the same distribution of electrons in the outermost principal energy level... 

To understand how position in the periodic table relates to the filling of sublevels, consider the metals in the first two groups. Atoms of the Group 1 elements all have one s electron in the outermost principal energy level (Table 6.4). In each Group 2 atom, there are two s electrons in the outermost level. A similar relationship applies to the elements in any group ... 

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. 

This idea is readily extended to simple molecules of compounds formed by nonmetal atoms. An example is the HF molecule. You will recall that a fluorine atom has the electron configuration ls22s22p5. ft has seven electrons in its outermost principal energy level (n = 2). These are referred to as valence electrons, in contrast to the core electrons filling the principal level, n = 1. If the valence electrons are shown as dots around the symbol of the element, the fluorine atom can be represented as... 

A the same number of orbitals B the same number of valence electrons C atomic numbers that are multiples of each other D the same principal energy levels... 

A the principal energy level increases and the first ionization energy increases... 

D the principal energy level decreases and the first ionization energy decreases This question covers NSCS B1 and B6. This question tests the material that was covered in the textbook on page 168. 

The valence electron for the cesium atom is in the 6s orbital. In assigning quantum numbers, n = principal energy level = 6. The quantum number l represents the angular momentum (type of orbital) with s orbitals = 0, p orbitals = 1, d orbitals = 2, and so forth. In this case, l = 0. The quantum number m is known as the magnetic quantum number and describes the orientation of the orbital in space. For, v orbitals (as in this case), mt always equals 0. For p orbitals, mt can take on the values of -1, 0, and +1. For d orbitals, can take on the values -2, -1, 0, +1, and +2. The quantum number ms is known as the electron spin quantum number and can take only two values, +1/2 and -1/2, depending on the spin of the electron. 

Regardless of its name, the second quantum number refers to the energy sublevels within each principal energy level. The name that this hook uses for the second quantum number is orbital-shape quantum number (i), to help you remember that the value of 1 determines orbital shape. (You will see examples of orbital shapes near the end of this section.)... 

There is, however, a system that chemists use alongside electron configurations to help them plot and keep track of electrons in their orbitals. An orbital diagram uses a box for each orbital in any given principal energy level. (Some chemists use a circle or a line instead of a box.) An empty box represents an orbital in which there are no electrons (an unoccupied orbital). A box that has an upward-pointing arrow represents an orbital with an electron that spins in one direction. A box with a downward-pointing arrow represents an orbital with an electron that spins in the opposite direction. You can... 

Which of the following is the correct orbital diagram for the third and fourth principal energy levels of a vanadium atom (Z = 23) Justify your answer. 

Chemical bonding involves the interaction of valence electrons—the electrons that occupy the outermost principal energy level of an atom. 

Explain how to determine the maximum number of electrons in any principal energy level. 

When a principal energy level (shell) receives its full complement of electrons (e.g., inert noble gases in group 18 (VIII)), a new row begins, which is the start of new period. 

E) Atoms get bigger as you go down groups. The reason is that principal energy levels of electrons are being added. Leaving the noble gases out, atoms get smaller as you go across a period. 

A) Positive ions are smaller than their neutral atoms and negative ions are larger than their neutral atoms. Mg is the only ion from the choices with only two principal energy levels of electrons. . . so it is the smallest. 

The transition must be n- cr [Problem 12.4(6)]. On going from Cl to Br to 1 the n electrons (a) are found in higher principal energy levels (the principal quantum numbers are 3, 4, 5, respectively), (b) are further away from the attractive force of the nucleus, and (c) are more easily excited. Hence absorption occurs at progressively higher A , , since less energy is required. 

While the theory of Bohr was a major step forward, and it helped to rmderstand the observed hydrogen spectrum, it left many other observations in the dark. New light was shed on the subject of atomic structure and the line spectra by Arnold Sormnerfeld (1868-1951) (27). He elaborated the basic theory of Bohr by observing that the orbits eould also be elliptical, and that for each principal energy level, there eotrld be a specific number of elliptical orbits of different... 

The inner shells of cere elecirofiS are often abbreviated since no nei bonding takes place in them. The symbols used, K, L. M. elc refer lo the older system of designating the principal energy levels, n = I (K). n = 2 (L), etc. Thus Na2 = KK LL oj, ... 

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