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Atomic Structure and Periodicity. Schrodinger Wave Equation. Equation for probability of a single electron being found along a single axis (x-axis). Erwin Schrodinger. Pauli Exclusion Principle. No two electrons in an atom can have the same four quantum numbers. Wolfgang Pauli. - PowerPoint PPT Presentation
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Atomic Structure and Periodicity
Schrodinger Wave Equation
22
2 28dh EVm dx
Equation for probability of a single electron being found along a single axis (x-axis)Erwin Schrodinger
Pauli Exclusion Principle
No two electrons in an atom can have the same four quantum numbers.
Wolfgang Pauli
Heisenberg Uncertainty Principle
You can find out where the electron is, but not where it is going.
OR…You can find out where the electron is going, but not where it is!
“One cannot simultaneously determine both the position and momentum of an electron.”
WernerHeisenberg
Quantum NumbersEach electron in an atom has a unique set of 4 quantum numbers which describe it.
Principal quantum number Angular momentum quantum number
Magnetic quantum number Spin quantum number
Principal Quantum NumberGenerally symbolized by n, it denotes the shell (energy level) in which the electron is located.
Tell the energyNumber of electrons that can fit in a shell:
2n2
N= 1, 2, 3, ….
Angular Momentum Quantum NumberThe angular momentum quantum number,
generally symbolized by l, denotes the orbital (subshell) in which the electron is located.
n2= the number of orbitals on an energy level
All orbitals with the same value forN are said to be degenerate.
Magnetic Quantum NumberThe magnetic quantum number, generally symbolized by m, denotes the orientation of the electron’s orbital with respect to the three axes in space.
Assigning the Numbers The three quantum numbers (n, l, and
m) are integers. The principal quantum number (n)
cannot be zero. n must be 1, 2, 3, etc. The angular momentum quantum
number (l ) can be any integer between 0 and n - 1.
For n = 3, l can be either 0, 1, or 2. The magnetic quantum number (ml) can
be any integer between -l and +l. For l = 2, m can be either -2, -1, 0, +1,
+2.
Principle, angular momentum, and magnetic quantum numbers: n, l, and ml
Spin Quantum NumberSpin quantum number denotes the behavior (direction of spin) of an electron within a magnetic field.
Possibilities for electron spin:
12
12
Magnetic Properties• Although an electron behaves like a tiny
magnet, two electrons that are opposite in spin cancel each other. Only atoms with unpaired electrons exhibit magnetic susceptibility – A paramagnetic substance is one that is
weakly attracted by a magnetic field, usually the result of unpaired electrons.
– A diamagnetic substance is not attracted by a magnetic field generally because it has only paired electrons.
Half of the distance between nucli in covalently bonded diatomic molecule
"covalent atomic radii"
Periodic Trends in Atomic Radius Radius decreases across a period
Increased effective nuclear charge dueto decreased shielding
Radius increases down a group Addition of principal quantum levels
Determination of Atomic Radius
Bond Radius
Table of Atomic Radii
Zeff constantn ↑r ↑
n constant
Zeff ↑r ↓
I.E. ↓
I.E. ↑
II/
Tend to be positive values endothermic process A + energy = A+ + e-
Ionization Energy: the energy required to remove an electron from an atom
John A. SchreifelsChemistry 211
Chapter 8-18
HIGHER IONIZATION ENERGIES
big jump in I.E. when core electrons start to be removed:
electrons from a lower main shell start to get removed. II/
Increases for successive electrons taken from the same atom Tends to increase across a period
Electrons in the same quantum level do not shield as effectively as electrons in inner levels
Irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove Tends to decrease down a group
Outer electrons are farther from thenucleus
Ionization Energy: the energy required to remove an electron from an atom
Trends in Electron Affinity
• The first occurs between Groups IA and IIA.– Added electron must
go in p-orbital, not s-orbital.
– Electron is farther from nucleus and feels repulsion from s-electrons.
Trends in Electron Affinity
• The second occurs between Groups IVA and VA.– Group VA has no empty
orbitals.– Extra electron must go
into occupied orbital, creating repulsion.
A + e- + energy → A-
A + e- → A- + energy
Trends in Electron Affinity
.
Can be an endothermic or exothermic process
Affinity tends to decrease as you go down in a period
Electrons farther from the nucleusexperience less nuclear attraction
Some irregularities due to repulsive forces in the relatively small p orbitals
Electron Affinity - the energy change associated with the addition of an electron
Affinity tends to increase across a period
Affinity tends to decrease as you go down in a period
Electrons farther from the nucleusexperience less nuclear attraction
Some irregularities due to repulsive forces in the relatively small p orbitals
Electron Affinity - the energy change associated with the addition of an electron
Table of Electron Affinities
A measure of the ability of an atom in a chemicalcompound to attract electrons
Electronegativities tend to increase across a period Electronegativities tend to decrease down a group or remain the same
Electronegativity
ElectronegativityWhat is it?Electronegativity is the power
of an atom to attract electron densityin a covalent bond
ElectronegativityPauling’s electronegativity scale
The higher the value, the more electronegative the element
Fluorine is the most electronegative elementIt has an electronegativity value of 4.0
ElectronegativityPauling’s electronegativity scale
Electronegativity
Periodic Table of Electronegativities
Cations Positively charged ions
Smaller than the corresponding
atomAnions Negatively charged ions Larger than the corresponding
atom
Ionic Radii
Sizes of ions• Ions are atoms that have either gained or lost
electrons (so that the # of electrons is not equal to the # of protons)
• The size of an atom can change dramatically if it becomes an ion
• E.g. when sodium loses its outer electron to become Na+ it becomes much smaller. Why?
• Na+ is smaller than Na because it has lost its 3s electron. Its valence shell is now 2s22p6 (it has a smaller value of n)
• Changing n values is one explanation for the size of ions. The other is …
Sizes of ions: electron repulsion• Valence electrons push each other away
9
+
• When an atom becomes a –ve ion (adds an electron to its valence shell) the repulsion between valence electrons increases without changing ENC
• Thus, F– is larger than F
Table of Ion Sizes
Summary of Periodic Trends electronegativity
electronegatvity
Ionic radius
Ioni
c ra
dius