9 -- Part 3: The Periodic Table and Some Atomic Properties
There are large zigzag breaks along the diagonal line.
Consider magnesium as an example.
The lower energy electron is 2 shell.
Si larger is not produced in ordinary chemical processes.
We don't see Na2+ or Al4+ in ordinary chemical processes.
Let's look at the exceptions to the regular trend.
The energy of Al is larger than that of Mg. To understand why, we have to consider the lost electrons.
The values of Ar are lower than expected and considerably lower than the values before the third row.
There are two explanations, one of which focuses elements into three sets, the other of which focuses elements into two sets.
The first explanation states that the key orbital is electron-electronpulsion repulsions.
The average unpaired 3p electron is closer together than the electrons in separate orbitals.
They are more easily removed than S, Cl, and repulsion.
The second explanation focuses on the strength of the attractions.
Unpaired electrons with parallel spins tend to avoid each other more, screen each other less, experience a higher effective nuclear charge, interact more strongly with the nucleus, and are harder to remove.
The value expected by the backward-extrapolation of the first ionization energies of S, Cl, and Ar is approximately 276 kJ mol-1.
Pairs of electrons screen each other to a greater extent, interact less strongly with the nucleus, and are easier to remove.
It is difficult to answer this question with complete certainty, but it is clear that the dip occurs once orbital sharing begins.
It's not surprising that it's hard to provide an unambiguous explanation for the observed dip.
The balance of electron-electron repulsions and electron-nucleus attractions is a delicate one.
As we move from P to S, a delicate balance of electron-electron repulsions and electron-nucleus attractions is needed.
Refer to the periodic table on the inside front cover and arrange the following in the expected order.
Ionization energies go down as atomic radii go up.
If we arrange these four atoms according to decreasing radius, we will likely have arranged them according to increasing ionization energy.
To the left and bottom of the table are the largest atoms.
The smallest atoms are to the right of the table.
The one that best fits the large-atom category is Sr.
The four atoms are not very close to the top of the table.
The two extremes are Sr with the lowest energy and Br with the highest.
The larger the tin atom, the lower the ionization energy.
The exceptions that occur when making comparisons between atoms of groups 2 and 13 as well as atoms of groups 15 and 16 are ignored by the generalization.
The periodic table can be found on the inside front cover.
The energy change for removing an electron is called irradiation energy.
The energy change is associated with the addition of an electron.
The definition states that the electron affinity of fluorine is a negative quantity.
The tendency for a neutral atom to gain an electron is reflected in the definition of electron affinity.
The X1g2 + e- indicates the tendency of an anion to lose an electron.
You should be aware that electron affinities are expressed in both ways in the chemical literature.
We have to consider the type of orbital in which the incoming electron has to be accommodated and the effect of the incoming electron on the electron-electron repulsions and electron-nucleus attractions.
C/eaH is a sign of net attraction between an atom and an electron.
X-(g) + e- is a process.
Period 3 tends to gain an electron.
Let's start with the variation of C/eaH.
Ignoring for the moment the atoms of groups 2, 15, and 18, we see from Figures 9-15 and 9-16 that as we move from left to right across a period, the addition of an electron becomes more favorable, with C/eaH taking on increasingly negative values.
The variation of C/eaH within a group, again ignoring groups 2, 15, and 18.
The value of C/eaH is the most negative for the atom of the third period and less negative as we move down the group.
The most negative C/eaH value is found in group 17 C/eaH becomes less negative because the incoming electron is accommodated in larger and larger orbitals and is less attracted to the nucleus.
The atom of the second row has a lower electron affinity than the third row.
The second row atoms are more compact than the third row atoms.
When an electron is added to a second row atom, it is likely that it encounters strong repulsive forces from other electrons in the atom and is not as tightly bound as we might otherwise expect.
C/eaH is positive for some atoms.
There is no tendency for these atoms to gain an electron.
The added electron would have to enter an orbital of the next subshell in one of the other cases.
The nitrogen atom shows little tendency to gain an electron.
The positive value shows that N-([He]2s22p4) is slightly more stable than N([He]2s22p3).
The increase in electron-electron repulsions or the decrease in electron-nucleus attractions may be why N is less stable than N.
Positive electron affinities are encountered when we consider the gain of a second electron by a nonmetal atom.
The electron is approaching a negative ion, not a neutral atom.
There is a repulsive force between the electron and the ion.
Oxygen has a negative first electron affinity and a positive second electron affinity.
The ion O2 can be found in ionic compounds, where formation of the ion is accompanied by other favorable processes.
The behavior of atoms and ion in a magnetic field is an important property.
An electron's spin causes it to generate a magnetic field.
A diamagnetic species is repelled by a magnetic field.
The unpaired electrons have a magnetic moment that causes the atom or ion to be attracted to a magnetic field outside.
The stronger this attraction is, the more unpaired electrons are present.
When a manganese atom loses two electrons, it becomes a paramagnetic ion, and the strength of its paramagnetism corresponds to five unpaired electrons.
The ion has a paramagnetism when a third electron is lost.
We need to determine if there are any unpaired electrons to determine if an atom or ion is paramagnetic.
The species is paramagnetic if there is at least one unpaired electron.
Outside the Ne core, the Na atom has a single 3s electron.
The electron is notpaired.
The other electrons in the Ne core must be pairs as well.
There are all electrons with 11s22s22p63s23p62.
We don't have to work out the exact configuration of Ag.
At least one of the electrons must be unpaired because the atom has 47 electrons.
The species is paramagnetic if the total number of electrons is odd.
The species may be paramagnetic if the total number of electrons is even.
Only 18 of the known atoms have diamagnetic states in their ground electronic states.
One of the 18 atoms must be isoelectronic for an ion to be diamagnetic.
The average distribution of electronic charge about the nucleus is spherical for an isolated atom.
This is not the case for an atom in the vicinity of another atom, molecule, or ion.
When an atom is placed in the electric field between two parallel plates, the position of the heavier nucleus is left essentially unchanged but the electron cloud is distorted.
The atom is said to be different.
Positive and negative charges are displaced from each other.
The magnitude of the displacement depends on how easy it is for the electron cloud to be distorted.
The atom is thought to be different.
Units of volume are used to express it.
The polarizability of an atom depends on how diffuse or spread out its electron cloud is, and in general, polarizability increases with the size of the atom.
polarizability decreases from left to right across a period and increases from top to bottom within a group.
The quantum mechanical calculations on atoms show that the polarizability of atoms is more important than the tightly bound molecule.
The result is not completely unexpected because, as we have already learned, the valence electrons are farther from the nucleus and experience a smaller effective nuclear charge than the phase electrons.
The electrons experience a greater shift in their positions.
There are two types of atomic radius, metallic and covalent.
The first ionization energy is referred to as Ionization energies.
We will use our knowledge of the variation of ion ization energy, electron affinity, and polarizability to explain bonding in the next few chapters.
The Al atom and Al3+ ion have very different polarizabilities.
Mercury should be a solid with a melting point over 300 C. It is a liquid at room temperature.
For a discussion of why mercury is a liquid, go to the Focus On feature for Chapter 9, The Periodic Law and Mercury, on the MasteringChemistry site.
Similar electron configurations can be found across the periodic table.
The periodic table has three classes of elements.
The periodic law is similar to that of atomic radii.
The size of the atom affects polarizability.
Interpretation of trends is more difficult than before.
An interesting relationship is observed for some of the isoelectronic atoms when they are compared.
The corresponding graph for the series is linear.
There is a graph shown.
There are differences in the numerical coefficients.
The electron configuration of the second-row atoms is 1s22s1, which is a single electron 12s12 beyond the helium core.
The electron configuration for the third row is 1s22s22p63s1, which is a single electron 13s12 beyond the neon core.
The inner-core electrons screen the single electron from the nucleus.
We can use the expression for the energy levels for hydrogen-like atoms or hydrogen-like ion as an approximation, because these species are all reminiscent of the Bohr atom.
We should be able to use equation 8.15 to derive equations for the energy required to remove the electron from the valence shells of a hydrogen-like species.
We can compare the equations for the two straight-line graphs once we have them.
We are on the correct path.
We have to take into account that we are considering oneelectron systems with a nucleus shielded by a closed shell.
The second-row series 1Li, Be+, B2+, C3+, N4+, O5+, and F6+2 all have the same electron configuration.
If we assume that the closed shell 1s2 screens the outer electron 2s1, the value of Zeff will be Z - 2.
The third-row series has been looked at.
The number of electrons screening the valence-shell electron is the reason for the difference in intercepts.