Electron Configuration SC1. Obtain, evaluate, and communicate information about the use of the modern atomic theory and periodic law to explain the characteristics of atoms and elements. g. Develop and use models, including electron configuration of atoms and ions, to predict an element’s chemical properties.

Electron Cloud Bohr’s model made the assumption that atoms are flat Electron location can be described with 1 coordinate, or quantum number n for energy level Schrödinger’s model expanded the atom into 3-dimensional space – the electron cloud This required more quantum numbers to describe the location of electrons l for subshell, ml for orbital, ms for spin

Electron Cloud Probable locations of electrons around the nucleus (based on their energy) can be predicted within the cloud in regions called subshells There are 4 kinds of subshells in use s (l=0), p (l=1), d (l=2), f (l=3)

Electron Cloud Variations on subshells are called orbitals s only has 1 orbital ml =0 p has 3 orbitals ml = -1, 0, 1 d has 5 orbitals ml = -2, -1, 0, 1, 2 f has 7 orbitals ml = -3, -2, -1, 0, 1, 2, 3

Shell-Subshell-Orbital

Quantum Numbers Quantum numbers are like the address for a particular electron in an atom. No two electrons, within an atom, can have the same set of 4 quantum numbers Carbon electron #1: n=1, l=0, ml=0, ms=+½ electron #2: n=1, l=0, ml=0, ms=-½ electron #3: n=2, l=0, ml=0, ms=+½ electron #4: n=2, l=0, ml=0, ms=-½ electron #5: n=2, l=1, ml=-1, ms=+½ electron #6: n=2, l=1, ml=0, ms=+½

Assigning Quantum Numbers Three rules: electrons fill subshells from lowest to highest energy (Aufbau) electrons fill orbitals singly before any orbital gets a second electron (Hund) no two electrons can fill one orbital with the same spin (Pauli)

Aufbau Principle 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f

Blocks and Subshells We can use the periodic table to predict which subshell is being filled by a particular element

Hund’s Rule Electrons fill orbitals INDIVIDUALLY before PAIRING UP!

Pauli Exclusion Principle Electrons PAIR UP with OPPOSITE spins ms= +½ or -½ 2 electrons per orbital Subshell maxima: s – 2 p – 6 d – 10 f – 14

Orbital Notation The orbital notation of an atom is a method of drawing the location of electrons. Aufbau is used to determine the order of filling, and arrows (either ↑ or ↓) are used to represent electrons.

Electron Configuration Notation The electron configuration notation of an atom is a shorthand method of writing the location of electrons. The energy level number is written first, followed by the subshell letter, and finally a superscript representing the number of electrons present.

Electron Configuration Notation There are a few exceptions to the “rules” If the e-config ends in d4, d9, f6 or f13 Chromium: Copper:

Noble Gas Notation This is a shorthand version of electron configuration notation You simply replace core electrons with a bracketed noble gas symbol Li = [He] 2s1 Si = [Ne] 3s23p2 Br = [Ar] 4s24p5 W = [Xe] 6s24f145d4

Valence Electrons When an atom undergoes a chemical reaction, only the outermost electrons are involved. These electrons, called the valence electrons, are of the highest energy and are farthest away from the nucleus. The valence electrons are typically the s and p electrons in the highest energy level represented.

Predicting Valence Electrons When using the IUPAC designations for group numbers for the representative elements, the ones digit corresponds to the number of valence electrons. The transition metals, which mostly have 2 (4s and 5s), will sometimes use d orbital electrons as valence electrons

Octet Rule Atoms which have a full set of eight valence electrons have increased stability The exceptions are elements 1-5. Why? The only atoms that have achieved this alone are called the “noble gases.” So how do the other atoms achieve this state? By sharing, gaining, or losing valence electrons!

Electron Dot Formulas An electron dot formula of an element shows the symbol of the element surrounded by its valence electrons. We use one dot for each valence electron. Consider phosphorus, P, which has 5 valence electrons. Here is the method for writing the electron dot formula.

Bonding When atoms share valence electrons, it is called covalent bonding. Nonmetals typically share with other nonmetals All atoms involved in the bond share enough valence electrons to achieve the stable octet. There are combinations that involve more than an octet (SF6 or PCl5)

Bonding When atoms lose or gain valence electrons, it is called ionic bonding Metals lose valence electrons and form cations, and nonmetals gain valence electrons and form anions The electrostatic attract between positive and negative cause the ionic bond to form

Predicting Ionic Charge Because there is a pattern of valence electrons, there is a pattern of ionic charges: Group 1 Group 2 Group 13 Group 14 Group 15 Group 16 Group 17

Ion Electron Configurations When we write the electron configuration of a positive ion, we remove one electron for each positive charge: Na → Na+ 1s2 2s2 2p6 3s1 → 1s2 2s2 2p6 When we write the electron configuration of a negative ion, we add one electron for each negative charge: O → O2- 1s2 2s2 2p4 → 1s2 2s2 2p6