Atoms principal energy levels

Quantum mechanical model of the atom, principal energy level, sublevel, electron orbital, Pauli exclusion principle... 

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... 

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... 

Energy level diagram of the sodium atom. The energy levels are denoted by the values for the principal quantum number , the orbital quantum number/, and the spin quantum number s. Levels with 1 = 0 are not split for / = 1 two separate levels are drawn (s = 1/2) for/> 1 the splitting is too small to be shown in the figure. Wavelengths of a few special transitions are given in nanometers. 

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. 

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. 

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. 

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... 

Note that going across the periodic table, the atomic This is due to the fact that the principal energy level number) remains the same, but the number of electrons iii in the number of electrons causes an increase in the el... 

The coefficients show that the sodium atom has three principal energy levels because the largest of these coefficients is a 3. 

Name electron subshells and atomic orbitals with the lowercase roman letters s, p, d, and f. Write principal energy levels 1-7 on the line and to the left of the letter give the number of electrons in the orbital as a superscript to the right of the letter. Specify orbital axes with italic subscripts. 

The wave-mechanical model of the atom shows a more complex structure of the atom and the way electrons configure themselves in the principal energy levels. Principal energy levels are divided into sublevels, each with its own distinct set of orbitals. This more complex structure is outlined with the help of this diagram. The principal energy levels in the atom are numbered 1 through 7. 

Valence electrons play a huge role in bonding, as will be shown later. Valence electrons are the electrons that are in the outermost principal energy level (not to be confused with the outermost subshell). These electrons are important because they are the electrons that are lost, gained, or shared when forming chemical bonds. The valence electrons of an atom are the electrons that interact with the valence electrons of another atom to form these bonds. 

There are exceptions to the octet rule. Helium, for example, is incredibly stable with just two valence electrons in its outermost principal energy level. The same holds true for lithium and beryllium as well. This indicates that it isn t so much having an octet that stabilizes the atom, as it is the issue of having a full outermost principal energy level. 

One last exception to the octet rule lies in the bonding of the atom boron. Boron prefers six electrons in its outermost principal energy level. This allows compounds containing boron to make three bonds in a trigonal planar arrangement. Two examples are BH3 and BF3 as shown in Figure 3.13. 

Covalent bonds are formed when two nonmetal atoms share electrons in order to satisfy their need to have a full outermost principal energy level. Covalent bonds are not as strong as the bonds formed between ions. For example, it would take a high flame and a temperature of almost 800 degrees Celsius to break the bonds between the sodium and chlorine in sodium chloride. The covalent bonds found in methane can be broken instantly with the introduction of a lit match. 

This diagram showing how the valence electrons interact is called a Lewis structure. In this case both hydrogen atoms have satisfied their need to have a full outermost principal energy level. Because both hydrogen atoms have the same electronegativity, the atoms will share the electrons equally. This will be the case with any diatomic molecule, such as chlorine gas (see Figure 5.6). 

BECAUSE neutrons and protons are both located in the principal energy levels of the atom. 

T, F The nucleons (protons and neutrons) are located in the nucleus of an atom. An atom s electrons are located in the principal energy levels. 

A bond formed when two nonmetal atoms share electrons in order to satisfy their need to have a full outermost principal energy level. 

A negatively charged particle that orbits the nucleus of an atom in the principal energy levels. 

An atom will desire eight electrons in its outermost principal energy level to maximize its stability. 

Going across the periodic table the electronegativity increases because the principal energy level remain the same and the electrostatic attraction increases. The atoms also have a desire to have the most stable configuration which is that of the noble gas configuration. Going down the table the electronegativity decreases due to the increased distance from the nucleus. 

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