Chapter 81 Atomic Electronic Configurations and Chemical Periodicity Chapter 8.
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Chapter 8 1 Atomic Electronic Atomic Electronic Configurations and Configurations and Chemical Periodicity Chemical Periodicity Chapter Chapter 8 8
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Transcript of Chapter 81 Atomic Electronic Configurations and Chemical Periodicity Chapter 8.
Chapter 8 1
Atomic Electronic Atomic Electronic Configurations and Configurations and
Chemical PeriodicityChemical Periodicity
Chapter 8Chapter 8
Chapter 8 2
Electron SpinElectron Spin
Effective Nuclear ChargeEffective Nuclear Charge
Effective nuclear charge - The charge experienced by an electron in a many-electron atom.
Chapter 8 3
Electron SpinElectron Spin
Effective Nuclear ChargeEffective Nuclear Charge - Electrons are attracted to the nucleus, but repelled by
the electrons that screen it from the nuclear charge.- The nuclear charge experienced by an electron
depends on its distance from the nucleus and the number of core electrons.
- As the average number of screening electrons (S) increases, the effective nuclear charge (zeff) decreases.
- As the distance from the nucleus increases, S increases and zeff decreases.
zeff = z - S
Chapter 8 4
Electron SpinElectron SpinPauli Exclusion PrinciplePauli Exclusion Principle
Pauli’s Exclusion Principle - - no two electrons can have
the same set of 4 quantum numbers.- Therefore, two electrons in the same orbital must
have opposite spins.-One electron has a spin of +½, the other –½.
Chapter 8 5
Electron SpinElectron Spin
Magnetism Magnetism
Diamagnetic – Property of a substance that is repelled by a magnetic field.
Paramagnetic – Property of a substance that is attracted by a magnetic field.
Whether a substance is diamagnetic or paramagnetic depends on its electron configuration.
- Atoms without “paired” spins result in a paramagnetic substance.
Chapter 8 6
Electron AssignmentElectron Assignment
Orbitals and Quantum NumbersOrbitals and Quantum Numbers
- Orbitals can be ranked in terms of energy to yield an orbital energy diagram.
- As n increases, note that the spacing between energy levels becomes smaller.
- As shielding increases the energy of an orbital increases.
Chapter 8 7
Electron AssignmentElectron Assignment
Chapter 8 8
Electron ConfigurationElectron ConfigurationElectron configurations tells us in which orbitals the electrons for an element are located.
Three rules:- Electrons fill orbitals starting with lowest n and
moving upwards.- no two electrons can fill one orbital with the same
spin (Pauli Exclusion Principle).- for degenerate orbitals, electrons fill each orbital
singly before any orbital gets a second electron (Hund’s rule).
Chapter 8 9
Electron ConfigurationElectron Configuration
Chapter 8 10
- There is a shorthand way of writing electron configurations
- Write the core electrons corresponding to the filled Noble gas in square brackets.
- Write the valence electrons explicitly.
Example, P: 1s22s22p63s23p3
but Ne is 1s22s22p6
Therefore, P: [Ne]3s23p3.
Electron ConfigurationElectron Configuration
Chapter 8 11
Electron Configuration of IonsElectron Configuration of Ions
Cations:To form a cation, an electron from the highest
principle quantum number and angular momentum quantum number is removed. Mg: [ 1s22s22p63s2] Mg2+: [1s22s22p6] + 2e- Anions:
To form an anion, an electron is added to the highest principle quantum number and angular momentum quantum number. Cl: [1s22s22p63s23p5] + 1e- Cl-: [1s22s22p63s23p6]
Chapter 8 12
Atomic Properties and Periodic TrendsAtomic Properties and Periodic Trends
Chapter 8 13
Atomic Properties and Periodic TrendsAtomic Properties and Periodic TrendsElectron Shells in AtomsElectron Shells in Atoms- Elements in the same column have the same electron
configuration.
Consider: Ne: 1s2 2s22p6
Ar: 1s2 2s22p6 3s23p6
Both elements have the same electron configuration:[Element]ns2np6
-Therefore, the elements in the periodic table should exhibit regular variations in there physical properties.
Chapter 8 14
- Atomic size varies consistently through the periodic table.
- As we move down a group, the atoms become larger.
- As we move across (left to right) a period, atoms become smaller.
- There are two factors at work:- principal quantum number, n- the effective nuclear charge, zeff
Atomic Properties and Periodic TrendsAtomic Properties and Periodic TrendsAtomic SizeAtomic Size
Chapter 8 15
- As the principle quantum number increases, the distance of the outermost electron from the nucleus becomes larger. Hence, the atomic radius increases.
- As we move across the periodic table, there is an increased attraction between the nucleus and the outermost electrons. This attraction causes the atomic radius to decrease.
Atomic Properties and Periodic TrendsAtomic Properties and Periodic TrendsAtomic SizeAtomic Size
Chapter 8 16
Atomic Properties and Periodic TrendsAtomic Properties and Periodic TrendsAtomic SizeAtomic Size
Chapter 8 17
First Ionization Energy - The first ionization energy, I1, is the amount of energy required to remove an electron from a gaseous atom:
Na(g) Na+(g) + e-
- The larger ionization energy, the more difficult it is to remove the electron.
- There is a sharp increase in ionization energy when a core electron is removed.
Atomic Properties and Periodic TrendsAtomic Properties and Periodic TrendsIonization EnergyIonization Energy
Chapter 8 18
Atomic Properties and Periodic TrendsAtomic Properties and Periodic TrendsIonization EnergyIonization Energy
Chapter 8 19
- Ionization energy decreases down a group. - As the atom gets bigger, it becomes easier to remove
an electron from the most spatially extended orbital.- Ionization energy generally increases across a period.
- Two exceptions: removing the first p electron and removing the fourth p electron have a lower energies.
- This indicates that half-filled and completely filled subshells are more stable.
Atomic Properties and Periodic TrendsAtomic Properties and Periodic TrendsIonization EnergyIonization Energy
Chapter 8 20
The s electrons are more effective at shielding than p electrons. Therefore, forming the s2p0 becomes more favorable.
When the p subshell has four electrons, one orbital has two electrons. When this electron is removed, the resulting s2p3 configuration is more stable than the starting s2p4 configuration (the final state has a much lower electron-electron repulsion).
Atomic Properties and Periodic TrendsAtomic Properties and Periodic TrendsIonization EnergyIonization Energy
Chapter 8 21
Atomic Properties and Periodic TrendsAtomic Properties and Periodic TrendsIonization EnergyIonization Energy
Chapter 8 22
Electron affinity – The energy required to add an electron to an atom in the gaseous state:
Cl(g) + e- Cl-(g)- Electron affinity can either be exothermic (as the
above example) or endothermic.
Atomic Properties and Periodic TrendsAtomic Properties and Periodic TrendsElectron AffinityElectron Affinity
Chapter 8 23
Atomic Properties and Periodic TrendsAtomic Properties and Periodic TrendsElectron AffinityElectron Affinity
Chapter 8 24
- Cations are smaller than the parent atom.- Anions are larger than the parent atom.
Atomic Properties and Periodic TrendsAtomic Properties and Periodic TrendsIon SizesIon Sizes
Chapter 8 25
- Since the representative (main group) elements have the same valence electron configuration, their chemical properties should be similar.
Example:2 Li(s) + Cl2(g) 2 LiCl(s)
2 Na(s) + Cl2(g) 2 NaCl(s)
2 K(s) + Cl2(g) 2 KCl(s)
Periodic Trends and Chemical PropertiesPeriodic Trends and Chemical Properties
Chapter 8 26
4, 20, 22, 32, 34, 44
HomeworkHomework
