Atomic Structure and the Periodic Table

Atomic Structure and the Periodic Table

Bohr’s single quantum number (n) was expanded to a total of four quantum numbers n, l, ml, and ms • These four quantized values describe an electron in an atom (quantized values are restricted to certain discrete values) • These values add order to our description of the electron in the atom

Table 1: Summary of Quantum Numbers

Table 2: Values and Letters for the Secondary Quantum Numbers

Table 3: Comparing Orbits and Orbitals

Table 4: Energy Levels, Orbitals, and Shells • The first two quantum numbers (n and l) describe electrons with different energies under NORMAL circumstances • The last two quantum numbers (ml and ms) describe electrons with different energies under SPECIAL conditions (e.g. magnetic field)

Moving forward, we will be focusing on the electrons position in space (not energy), the language will change • Main (principal) energy level = shell • Energy sublevel = subshell • WHY? • Its easier! • 1s orbital can be communicated as n=1, l=0 • 2p orbital can be communicated as n=2, l=1

Table 5: Classification of Energy Sublevels (subshells)

Energy-Level Diagrams • Figure 1: Energy-level diagrams show the relative energies of electrons in various orbitals under normal conditions (each orbital can hold a maximum of 2 e-)

The energy of an electron increases with an increasing value of principal quantum number, n • For a given number of n, the subshells increase in energy, in order, s<p<d<f • When creating energy-level diagrams, an electron in an orbital is represented by drawing an arrow, pointed up or down in a specific circle, but two arrows in a circle MUST be in opposite directions Figure 2: Energy-level diagrams for (a) hydrogen (b) helium

Pauli Exclusion Principle – no two electrons in an atom can have the same four quantum numbers; no two electrons in the same atomic orbital can have the same spin, only two electrons with opposite spins can occupy any one orbital • What order do we fill the orbitals? • Aufbau Principle – each electron is added to the lowest energy orbital available in an atom or ion • An energy sublevel must be filled before moving onto the next higher sublevel

Figure 3: In this aufbau diagram, start at the bottom (1s) and add electrons in the order shown by the diagonal arrows. You work your way from the bottom left corner to the top right corner.

Figure 4: Classification of elements by the sublevels that are being filled

Hund’s Rule – one electron occupies each of the several orbitals at the same energy before a second electron can occupy the same orbital SEATWORK • Read pp. 189 – 199 - Drawing energy-level diagrams for atoms, anions, cations • Practice p. 191 UC # 3, 4

Complications • If you’re thinking this is too easy to be true, you’re right! • There are a few complications as the atoms get larger • As the energy level gets farther from the nucleus, the distance between energy levels decreases • As a matter of fact, it is believed that the energy levels actually overlap

Complications • Therefore, some energy levels start filling orbitals before the previous energy level is finished filling its subshell • The first time this is encountered is with potassium, in which the 4s starts to fill before the 3d

There’s More… • The second complication has to do with a variation of Hund’s Rule that takes into account the minimizing of the electron-electron repulsion • It states, the most stable arrangement of electrons is the arrangement with the maximum number of unpaired electrons. • So, when the transition metals’ orbitals are filling with electrons, at d4 and d9, an electron from the s JUMPS up into the d5 and d10

Why are some electrons promoted? • Overall energy state of the atom is lower after the promotion of the electrons • Half-filled and filled subshells are more stable (lower energy) than unfilled subshells

Electron Configuration • A method of communicating the location and number of electrons in electron energy levels (presents same information as energy-level diagrams BUT much more concise) Figure 5: Example of electron configuration

Writing Electron Configurations • The electron configuration below represents a boron atom in its ground state. • The superscripts indicate the number of electrons occupying each sublevel.

Electron Configuration Shorthand • Writing out electron configurations can become awkward as the atoms increase in the number of electrons • The shorthand involves using the abbreviation of the last noble gas (placed in square brackets) to indicate that all the orbitals to that point are full. Then the configuration is continued as usual.

Nitrogen 1s22s22p3 [He] 2s22p3 Chromium 1s22s22p63s23p64s23d4 [Ar] 4s23d4 Strontium 1s22s22p63s23p64s23d104p65s2 [Kr] 5s2

Learning Checkpoint • Read pp. 192 – 193 • Understand FULL electron configuration and Shorthand (NOBLE GAS CORE) electron configurations • Add the summary for “Procedure for Writing an Electron Configuration” on p. 193 • Complete “Electron Configuration” worksheet • Practice Questions p. 194 UC # 6, 8, 9, 10 • Section 3.6 Questions p. 197 UC # 2, 3, 4, 5, 6, 7, 8, 9 10, 11, 12, 13, 14

Atomic Structure and the Periodic Table