Chapter 5. Periodicity and the Electronic Structure of Atoms

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1 Chapter 5 Periodicity and the Electronic Structure of Atoms

spinning charged particles generate a magnetic field if there is an even number of electrons, about half the

2 Electron Spin experiments by Stern and Gerlach showed a beam of silver atoms is split in two by a magnetic field the experiment reveals that the electrons spin on their axis as they spin, they generate a magnetic field spinning charged particles generate a magnetic field if there is an even number of electrons, about half the atoms will have a net magnetic field pointing North and the other half will have a net magnetic field pointing South 2

orbital paired Electrons have spin, which gives rise to a tiny magnetic field and to a spin quantum

3 Electron Spin and the Pauli Exclusion Principle spin quantum number describes how the electron spins on its axis clockwise or counterclockwise spin up or spin down spins must cancel in an orbital paired Electrons have spin, which gives rise to a tiny magnetic field and to a spin quantum number (m s ). Pauli Exclusion Principle: No two electrons in an atom can have the same four quantum numbers.

4 Spin Quantum Number, m s spin quantum number describes how the electron spins on its axis clockwise or counterclockwise spin up or spin down spins must cancel in an orbital m s

4 4 Spin Quantum Number, m s spin quantum number describes how the electron spins on its axis clockwise or counterclockwise spin up or spin down spins must cancel in an orbital m s can have values of ±½ Two electrons in the same orbital are referred as Spin paired A single electron in an orbital is referred to as unpaired. An orbital without any electrons is empty

5 Pauli Exclusion Principle no two electrons in an atom may have the same set of 4 quantum numbers therefore no orbital may have more than 2 electrons, and they must have with opposite spins Complete set of four quantum number (n, l, ml, ms) for the two electrons in a 3s orbital by supplying the spin quantum number, ms. 3s (3, 0, 0, ) abd (3,0,0, ) knowing the number orbitals in a sublevel allows us to determine the maximum number of electrons in the sublevel s sublevel has 1 orbital, therefore it can hold 2 electrons p sublevel has 3 orbitals, therefore it can hold 6 electrons d sublevel has 5 orbitals, therefore it can hold 10 electrons f sublevel has 7 orbitals, therefore it can hold 14 electrons 5

6 Electron Configurations of Multielectron Atoms Electron Configuration: A description of which orbitals are occupied by electrons Ground-State Electron Configuration: The lowest-energy (most stable) configuration A set of rules has been developed so that electrons can be placed into an orbital energy diagram in a ground state configuration Always place electrons into the lowest energy available orbital first. Aufbau Principle ( building up ): A guide for determining the filling order of orbitals Lower-energy orbitals fill before higher-energy orbitals. An orbital can hold only two electrons, which must have opposite spins (Pauli exclusion principle). If two or more degenerate orbitals are available, follow Hund s rule. Hund s Rule: If two or more orbitals with the same energy are available, one electron goes into each until all are half-full. The electrons in the half-filled orbitals all have the same value of their spin quantum number. Degenerate Orbitals: Orbitals that have the same energy level for example, the three p orbitals in a given subshell

7 Electron Configurations of Multielectron Atoms

8 Electron Configurations of Multielectron Atoms Electron Configuration H: 1s 1 1 electron s orbital (l = 0) n = 1 Na O P

9 Electron Configurations of Multielectron Atoms Ground states of elements in orbital energy filling diagram

10 Electron Configurations of Multielectron Atoms Electron Configuration Shorthand Configuration Na: 1s 2 2s 2 2p 6 3s 1 Ne configuration

11 Electron Configurations of Multielectron Atoms Electron Configuration Na: 1s 2 2s 2 2p 6 3s 1 Shorthand Configuration [Ne] 3s 1 P: 1s 2 2s 2 2p 6 3s 2 3p 3 Ne configuration

12 Anomalous Electron Configurations Expected Configuration Actual Configuration Cr: [Ar] 4s 2 3d 4 [Ar] 4s 1 3d 5 Cu: [Ar] 4s 2 3d 9 [Ar] 4s 1 3d 10

13 Electron Configurations and the Periodic Table

Electron Configurations and the Periodic Table Valence Shell: Outermost shell and any electrons on the outermost shells are called valence electrons Valence electrons are important because they

14 Electron Configurations and the Periodic Table Valence Shell: Outermost shell and any electrons on the outermost shells are called valence electrons Valence electrons are important because they re most energetic (least stable) electrons in an atoms. - They are involved in chemical reaction (e.g transferring electrons) or forming covalent bonds Li: 2s 1 Na: 3s 1 Cl: 3s 2 3p 5 Br: 4s 2 4p 5

This periodic table shows the patterns of valence electron configurations in the periodic table. Remember, the elements are grouped and displayed in a repetitive pattern, why?

15 This periodic table shows the patterns of valence electron configurations in the periodic table. Remember, the elements are grouped and displayed in a repetitive pattern, why? When arranged in the order of the periodic table, the valence electron configurations of the elements generally repeat in a pattern, so do the properties of the elements repeat in pattern as a result. Now let s fill in the general valence electrons for the groups headed by H, Be, B, C, N, O, F, Ne, Ti, Zn

16 Anomalous Electron Configurations Elements past argon: Rather than continue filling the third shell by populating the 3d orbitals, the next two electrons in potassium and calcium go into the 4s subshell. Only then does filling of the 3d subshell occur to give the first transition metal series from the scandium through zinc. Expected Configuration Actual Configuration Cr: [Ar] 4s 2 3d 4 [Ar] 4s 1 3d 5 Cu: [Ar] 4s 2 3d 9 [Ar] 4s 1 3d 10

Orbital Energy Levels in Multielectron Atoms One of the many periodic properties of the elements that can be explained by electron configurations in size, or atomic radius.

17 Orbital Energy Levels in Multielectron Atoms One of the many periodic properties of the elements that can be explained by electron configurations in size, or atomic radius. Positive charge (e.g protons) is attracted to negative charg (e.g electrons). The greater the positive charge the stronger the attraction. By contrast, negative charge is repelled by other negative charge.

Orbital Energy Levels in Multielectron Atoms Effective Nuclear Charge (Z eff ): The nuclear charge actually felt by an electron Z eff = Z actual Electron shielding Outer electrons are attracted to

18 Orbital Energy Levels in Multielectron Atoms Effective Nuclear Charge (Z eff ): The nuclear charge actually felt by an electron Z eff = Z actual Electron shielding Outer electrons are attracted to the positive charge (protons) in the nucleus, but repelled by the negative charge of the electrons that are nearer to the nucleus (i.e., electrons in the lower energy shell). As result, the effective or net positive charge felt by an outer electron is less than the full positive charge of nucleus. The outer electrons are said to be shielded from some of the positive charge in the nucleus by the inner electrons.

19 Effective Nuclear Charge in a multi-electron system, electrons are simultaneously attracted to the nucleus and repelled by each other outer electrons are shielded from full strength of nucleus screening effect effective nuclear charge is net positive charge that is attracting a particular electron Z is nuclear charge, S is electrons in lower energy levels electrons in same energy level contribute to screening, but very little effective nuclear charge on sublevels trend, s > p > d > f Z effective = Z - S 19

Electron Configurations and Periodic Properties: Atomic Radii The increase in radius going down in a group of the periodic table occurs because

20 Electron Configurations and Periodic Properties: Atomic Radii The increase in radius going down in a group of the periodic table occurs because successively larger valence-shell orbitals are occupied. Larger shells are occupied, the atomic radii are also larger. Column Radius Row Radius

Electron Configurations and Periodic Properties: Atomic Radii Atomic size decreases as you go across a row from left to right because of an increase in effective nuclear charge for the valence-shell

21 Electron Configurations and Periodic Properties: Atomic Radii Atomic size decreases as you go across a row from left to right because of an increase in effective nuclear charge for the valence-shell electrons. Going across the period, each additional electron adds to the same shell which at approximately the same distance from the nucleus, they are relatively ineffective at shielding one another. Z eff for the valence-shell electrons increases across the period, drawing all the valence-shell electrons closer to the nucleus and progressively shrinking the atomic radii

elements and compounds Atomic Radius Increases down group valence shell farther from nucleus effective nuclear charge fairly close Atomic

22 Trend in Atomic Radius Main Group Different methods for measuring the radius of an atom, and they give slightly different trends van der Waals radius = nonbonding covalent radius = bonding radius atomic radius is an average radius of an atom based on measuring large numbers of elements and compounds Atomic Radius Increases down group valence shell farther from nucleus effective nuclear charge fairly close Atomic Radius Decreases across period (left to right) adding electrons to same valence shell effective nuclear charge increases valence shell held closer 22

Chapter 8. Periodic Properties of the Element

Chapter 8. Periodic Properties of the Element Chapter 8 Periodic Properties of the Element Mendeleev (1834 1907) Ordered elements by atomic mass Saw a repeating pattern of properties Periodic law when the elements are arranged in order of increasing