a la) What are valence electrons ? Where are valence electrons situated in an atom?
) What is the number of valence electrons in the atoms of an element having atomic number 13 ? Name the
valence shell of this atom.
) What are isotopes ? Explain by giving an example,
(1) Give one similarity and one difference between a pair of isotopes.
Give the number of protons, neutrons and electrons per atom in the two isotopes of chlorine Pa and
pa
(a) What are radioactive isotopes ? Give two examples of radioactive isotopes.
(D) Give any two uses of radioactive isotopes.
(d) An element Z contains two naturally occurring isotopes 2 and 2. If the average atomic mass of this
element be 35.5 u, calculate the percentage of two isotopes.
2 (a) Define valency of an element. What valency will be shown by an element having atomic number 14?
(1) What is the relation between the valency of an element and the number of valence electrons in its atoms?
Explain with examples
Answers
Answer:
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Answer:
energy of atomic orbitals increases as the principal quantum number, n, increases. In any atom with two or more electrons, the repulsion between the electrons makes energies of subshells with different values of l differ so that the energy of the orbitals increases within a shell in the order s < p < d < f. Figure 1 depicts how these two trends in increasing energy relate. The 1s orbital at the bottom of the diagram is the orbital with electrons of lowest energy. The energy increases as we move up to the 2s and then 2p, 3s, and 3p orbitals, showing that the increasing n value has more influence on energy than the increasing l value for small atoms. However, this pattern does not hold for larger atoms. The 3d orbital is higher in energy than the 4s orbital. Such overlaps continue to occur frequently as we move up the chart.
A table entitled, “Subshell electron capacity,” is shown. Along the left side of the table, an upward pointing arrow labeled, “E,” is drawn. The table includes three columns. The first column is narrow and is labeled, “2.” The second is slightly wider and is labeled, “6.” The third is slightly wider yet and is labeled, “10.” The fourth is the widest and is labeled, “14.” The first column begins at the very bottom with a horizontal line segment labeled “1 s.” Evenly spaced line segments continue up to 7 s near the top of the column. In the second column, a horizontal dashed line segment labeled, “2 p,” appears at a level between the 2 s and 3 s levels. Similarly 3 p appears at a level between 3 s and 4 s, 4 p appears just below 5 s, 5 p appears just below 6 s, and 6 p appears just below 7 s. In the third column, a dashed line labeled, “3 d,” appears just below the level of 4 p. Similarly, 4 d appears just below 5 p and 5 d appears just below 6 p. Six d however appears above the levels of both 6 p and 7 s. The far right column entries begin with a dashed line labeled, “4 f,” positioned at a level just below 5 d. Similarly, a second dashed line segment appears just below the level of 6 d, which is labeled, “5 f.”
Figure 1. Generalized energy-level diagram for atomic orbitals in an atom with two or more electrons (not to scale).
Electrons in successive atoms on the periodic table tend to fill low-energy orbitals first. Thus, many students find it confusing that, for example, the 5p orbitals fill immediately after the 4d, and immediately before the 6s. The filling order is based on observed experimental results, and has been confirmed by theoretical calculations. As the principal quantum number, n, increases, the size of the orbital increases and the electrons spend more time farther from the nucleus. Thus, the attraction to the nucleus is weaker and the energy associated with the orbital is higher (less stabilized). But this is not the only effect we have to take into account. Within each shell, as the value of l increases, the electrons are less penetrating (meaning there is less electron density found close to the nucleus), in the order s > p > d > f. Electrons that are closer to the nucleus slightly repel electrons that are farther out, offsetting the more dominant electron–nucleus attractions slightly (recall that all electrons have −1 charges, but nuclei have +Z charges). This phenomenon is called shielding and will be discussed in more detail in the next section. Electrons in orbitals that experience more shielding are less stabilized and thus higher in energy. For small orbitals (1s through 3p), the increase in energy due to n is more significant than the increase due to l; however, for larger orbitals the two trends are comparable and cannot be simply predicted. We will discuss methods for remembering the observed order.