Chapter 7 Periodic Properties of the Elements PDF

7.2 Effective Nuclear Charge  The attractive force between an electron and the nucleus increases as the nuclear charge increases and decreases as the electron moves farther from the nucleus.  In a many-electron atom, in addition to the attraction of each electron to the nucleus, each elect...

7.2 Effective Nuclear Charge  The attractive force between an electron and the nucleus increases as the nuclear charge increases and decreases as the electron moves farther from the nucleus.  In a many-electron atom, in addition to the attraction of each electron to the nucleus, each electron experiences the repulsion due to other electrons.  Each electron in a many-electron atom is screened from the nucleus by the other electrons.  The partially screened nuclear charge is called the effective nuclear charge, Zeff. 3 7.2 Effective Nuclear Charge  The amount of screening of the actual nuclear charge (Z) is defined by using a screening constant (S), appositive number, such that 4 7.2 Effective Nuclear Charge  The notion of effective nuclear charge also explains an important effect: For a many-electron atom, the energies of orbitals with the same n value increase with increasing value.  The effective nuclear charge increases from left to right across any period of the periodic table.  The effective nuclear charge increases slightly as we go down a column because the more diffuse core electron cloud is less able to screen the valence electrons from the nuclear charge. 5 7.3 Sizes of Atoms and Ions  According to the quantum-mechanical model, atoms do not have sharply defined boundaries at which the electron distribution becomes zero.  The radius of an atom is called the nonbonding atomic radius or the van der Waals radius.  The bonding atomic radius (also known as the covalent radius) for any atom in a molecule is equal to half of the bond distance d. 6 7.3 Sizes of Atoms and Ions 7 Sample Exercise 7.1 Natural gas used in home heating and cooking is odorless. Because natural gas leaks pose the danger of explosion or suffocation, various smelly substances are added to the gas to allow detection of a leak. One such substance is methyl mercaptan, CH3SH. Use Figure 7.7 to predict the lengths of the C―S, C―H, and S―H bonds in this molecule. Bond length Bond length Bond length 8 7.3 Sizes of Atoms and Ions Periodic Trends in Atomic Radii  Within each group, bonding atomic radius tends to increase from top to bottom: This trend results primarily from the increase in the principal quantum number (n) of the outer electrons.  Within each period, bonding atomic radius tends to decrease from left to right: The major factor influencing this trend is the increase in effective nuclear charge Zeff across a period. The increasing effective nuclear charge steadily draws the valence electrons closer to the nucleus, causing the bonding atomic radius to decrease. 9 Sample Exercise 7.2 Referring to the periodic table, arrange (as much as possible) the atoms B, C, Al, and Si in order of increasing size. C Ca2+ > Mg2+ 13 Sample Exercise 7.3: Practice Exercise 1 Arrange the following atoms and ions in order of increasing ionic radius: F, S2–, Cl, and Se2–.  F < Cl < S2– < Se2– 14 7.3 Sizes of Atoms and Ions  An isoelectronic series is a group of ions all containing the same number of electrons.  In any isoelectronic series, because the number of electrons remains constant, ionic radius decreases with increasing nuclear charge as the electrons are more strongly attracted to the nucleus: 15 Sample Exercise 7.4 Arrange the ions K+, Cl–, Ca2+, and S2– in order of decreasing size.  S2– > Cl– > K+ > Ca2+ 16 Sample Exercise 7.4: Practice Exercise 1 Arrange the following ions in order of increasing ionic radius: Br–, Rb+ , Se2–, Sr2+, Te2–.  Sr2+ < Rb+ < Br– < Se2– < Te2– 17 7.4 Ionization Energy  The ionization energy of an atom or ion is the minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion.  The first ionization energy, I1, is the energy needed to remove the first electron from a neutral atom.  The second ionization energy, I2, is the energy needed to remove the second electron. 18 7.4 Ionization Energy Variations in Successive Ionization Energies  The greater the ionization energy, the more difficult it is to remove an electron.  Thus, the ionization energies for a given element increase as successive electrons are removed: I1 < I2 < I3, and so forth. 19 7.4 Ionization Energy Periodic Trends in First Ionization Energies  I1 generally increases as we move left to right across a period.  I1 generally decreases as we move down any column in the periodic table. 20 7.4 Ionization Energy  The s- and p-block elements show a larger range of I1 values than do the transition-metal elements.  Generally, the ionization energies of the transition metals increase slowly from left to right in a period.  As we move across a period, there is both an increase in effective nuclear charge and a decrease in atomic radius, causing the ionization energy to increase.  As we move down a column, the atomic radius increases, while the effective nuclear charge increases only gradually; the increase in radius dominates, so the attraction between the nucleus and the electron decreases, causing the ionization energy to decrease. 21 7.4 Ionization Energy  The irregularities in a given period are subtle but still readily explained.  The decrease in ionization energy from beryllium ([He]2s2) to boron ([He]2s22p1), occurs because the third valence electron of B must occupy the 2p subshell, which is empty for Be.  The slight decrease in ionization energy when moving from nitrogen ([He]2s22p3) to oxygen ([He]2s22p4) is the result of the repulsion of paired electrons in the p4 configuration. 22 7.4 Ionization Energy Electron Configurations of Ions  When electrons are removed from an atom to form a cation, they are always removed first from the occupied orbitals having the largest principal quantum number, n.  Thus, in forming ions, transition metals lose the valence-shells electrons first, then as many d electrons as required to reach the charge of the ion. 23 7.4 Ionization Energy  If there is more than one occupied subshell for a given value of n, the electrons are first removed from the orbital with the highest value of.  Electrons added to an atom to form an anion are added to the empty or partially filled orbital having the lowest value of n. 24 Sample Exercise 7.6 Referring to the periodic table, arrange the atoms Ne, Na, P, Ar, K in order of increasing first ionization energy.  K < Na < P < Ar < Ne 25 Sample Exercise 7.7 Write the electron configurations for a. Ca2+ 20 20 b. Co3+ 27 27 c. S2– 16 16 26 7.5 Electron Affinity  All ionization energies for atoms are positive: Energy must be absorbed to remove an electron.  The energy change that occurs when an electron is added to a gaseous atom is called the electron affinity because it measures the attraction, or affinity, of the atom for the added electron.  For most atoms, energy is released when an electron is added. 27 7.5 Electron Affinity  The greater the attraction between an atom and an added electron, the more negative the atom’s electron affinity.  For some elements, such as the noble gases, the electron affinity has a positive value, meaning that the anion is higher in energy than are the separated atom and electron:  The fact that the electron affinity is positive means that an electron will not attach itself to an Ar atom; in other words, the Ar– ion is unstable and does not form. 28