Figure 11.20: (a) The hydrogen 1s orbital. (b) The size of the orbital is defined by a sphere that contains 90% of the total electron probability.

Figure 11.21: The first four principal energy levels in the hydrogen atom. Each level is assigned an integer, n.

Orbitals and Energy Levels (cont.)

The 1s orbital.

Atomic Sublevels & Orbitals

• Each type of sublevel has a different shape each and energy.

• Each sublevel contains one or more orbitals.

Atomic Sublevels & Orbitals (cont.)

Atomic Sublevels & Orbitals (cont.)

Atomic Sublevels & Orbitals (cont.)

Figure 11.24: The relative sizes of 1s and 2s orbitals of hydrogen.

Figure 11.25: The three 2porbitals: (a) 2px, (b) 2pz, and ( c)2py.

Figure 11.26: A diagram of principal energy levels 1 and2 showing the shapes of orbitals that compose the sublevels.

Figure 11.27: The relative sizes of thespherical 1s, 2s, and 3s orbitals of hydrogen.

Figure 11.28: The shapes and labels of the five 3d orbitals.

Ψ = fn(n, l, ml, ms)

magnetic quantum number ml

for a given value of lml = -l, …., 0, …. +l

orientation of the orbital in space

if l = 1 (p orbital), ml = -1, 0, or 1if l = 2 (d orbital), ml = -2, -1, 0, 1, or 2

Schrodinger Wave Equation

ml = -1 ml = 0 ml = 1

ml = -2 ml = -1 ml = 0 ml = 1 ml = 2

Ψ = fn(n, l, ml, ms)

spin quantum number ms

ms = +½ or -½

Schrodinger Wave Equation

ms = -½ms = +½

Existence (and energy) of electron in atom is described by its unique wave function Ψ.

Pauli exclusion principle - no two electrons in an atomcan have the same four quantum numbers.

Schrodinger Wave Equation

Ψ = fn(n, l, ml, ms)

Each seat is uniquely identified (E, R12, S8)Each seat can hold only one individual at a time

Pauli Exclusion Principle

• No orbital may have more than 2 electrons.

• Electrons in the same orbital must have opposite spins.

• s sublevel holds 2 electrons (1 orbital)

• p sublevel holds 6 electrons (3 orbitals)

• d sublevel holds 10 electrons (5 orbitals)

• f sublevel holds 14 electrons (7 orbitals)

• For a multiple-electron atom, build-up the energy levels, filling each orbital in succession by energy

• Degenerate orbitals: orbitals with the same energy

– e.g. Each p sublevel has 3 degenerate p orbitals

Orbitals, Sublevels & Electrons

Schrodinger Wave Equation

Ψ = fn(n, l, ml, ms)

Shell – electrons with the same value of n

Subshell – electrons with the same values of n and l

Orbital – electrons with the same values of n, l, and ml

How many electrons can an orbital hold?

If n, l, and ml are fixed, then ms = ½ or - ½

Ψ = (n, l, ml, ½) or Ψ = (n, l, ml, -½)An orbital can hold 2 electrons

Orbitals, Sublevels & Electrons (cont.)

Electron Configurations

• For a set of degenerate orbitals, fill each orbital half-way first before pairing

• Electron configurations show how many electrons are in each sublevel of an atom – describes where electrons are.

- 1s22s1 is the electron configuration for a ground state Li

- 1s22s22p3 is for nitrogen

Electron Configurations (cont.)

• Valence shell: highest energy level– Electrons in the valence shell are called valence

electrons.– Core electrons: electrons not in the valence

shell– Often use symbol of previous noble gas in

brackets to represent core electrons, giving[He]2s22p3 for nitrogen or [Ne]3s2 for

magnesium

Electron Configuration and the Periodic Table

• Elements in the same column on the periodic table have: – Similar chemical and physical properties

– Similar valence shell electron configurations• same numbers of valence electrons

• same orbital types

• different energy levels

The locations and electron configurations of hydrogen and helium.

The locations and electron configurations of lithium and beryllium.

Boron (Group 3), carbon (Group 4), and nitrogen (Group 5).

Oxygen (Group 6), fluorine (Group 7), and neon (Group 8)

Figure 11.29: The electron configurations in the sublevel last occupied for the first eighteen elements.

s1s2

d1 d2 d3 d4 d5 d6 d7 d8 d9 d10

p1 p2 p3 p4 p5 s2p6

f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14

1234567

Figure 11.30: Partial electron configurationsfor the elements potassium through krypton.

Figure 11.31: The orbitals being filled for elements in various parts of the periodic table.

Figure 11.32: A box diagram showing the order in which orbitals fill to produce the atoms

in the periodic table. Each box can hold two electrons.

Order of orbitals (filling) in multi-electron atom

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s

11

Figure 11.33: The positions of the elements considered in Example 11.3.

Figure 11.34: The periodic table with atomic symbols, atomic numbers, and partial electron configurations.

Metallic Character: Metals

• Metals– Malleable & ductile

– Shiny, lustrous

– Conduct heat and electricity

– Most oxides basic and ionic

– Form cations in solution

– Lose electrons in reactions - oxidized

Metallic Character: Metalloids

• Metalloids�Also known as semi-metals

�Show some metal and some nonmetal properties

Metallic Character: Nonmetals

• Nonmetals�Brittle in solid state

�Dull

�Electrical and thermal insulators

�Most oxides are acidic and molecular

�Form anions and polyatomic anions

�Gain electrons in reactions - reduced

Metallic Character

Metallic Character (cont.)

• Reactivity of metals increases to the left on the period and down in the column

– Follows ease of losing an electron

• Reactivity of nonmetals (excluding the noble gases) increases to the right on the period and up in the column

The Group 1 elements: the farther down a group and element appears, the more likely it is to lose an electron.

The Group 2 elements: the farther down a group and element appears, the more likely it is to lose an electron.

Trend in Ionization Energy

• Minimum energy needed to remove a valence electron from an atom– Gas state

• The lower the ionization energy, the easier it is to remove the electron.– Metals have low ionization energies

• Ionization energy decreases down the group.– Valence electron farther from nucleus

• Ionization energy increases across the period.– Left to right

Ionization energies tend to decrease in going from the top to the bottom of a group.

Ionization energies tend to increase from left to right across a given period on the periodic table.

Trend in Atomic Size

Trend in Atomic Size (cont.)