Shells and Subshells The orbitals in an atom are arranged in shells and subshells. Shell: all orbitals with the same value of n Subshell: all orbitals with the same value of both n and l orbital n=3 3s 3p 3d n=3 3s 3p 3d n=3 3s 3p 3d

Subshell Energy For a hydrogen atom (or an ion containing only 1 electron) all orbitals within the same shell are degenerate. n=1 1s n=2 n=3 2s 2p 3s 3p 3d Energy

1s 3s 2p 3p 3d Energy 2s 4s 4p Subshell Energy For atoms with more than one electron, electron-electron repulsion causes different subshells within the same shell to have different energies. Within the same shell: s < p < d < f

Subshell Energy The relative energies of the various subshells can be predicted using the diagonal diagram: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f

Arrangement of Electrons The arrangement of the electrons in an atom can be depicted using three different but related methods: Orbital diagram Electron configuration 1s22s22p63s1 Electron configuration using core notation [Ne]3s1

Arrangement of Electrons Rules for populating orbitals with electrons: Pauli Exclusion Principle: Each electron in an atom must have a unique set of four quantum numbers n, l, ml, and ms. In order to put more than one electron in an orbital, electrons must have different values of ms. i.e. they must have different spins Maximum of 2 electrons per orbital

Arrangement of Electrons Rules for populating orbitals with electrons: Aufbau Principle: Electrons are placed in the lowest energy orbital available. Hund’s Rule: If more than one orbital in a subshell is available, electrons will fill empty orbitals in that subshell first. Keep electrons unpaired (in an orbital by itself) as long as an empty orbital with the same energy is available.

Orbital Diagrams Example: Draw an orbital diagram for each of the following atoms. Hydrogen: Helium: Lithium: Beryllium:

Orbital Diagrams Example: Draw an orbital diagram for each of the following atoms. Boron: Carbon: Nitrogen: Neon:

Orbital Diagrams The orbital diagram for Ne: The 2p subshell is completely filled. The outermost shell (n=2) contains an octet (8) of electrons. 1s 2s 2p

Orbital Diagrams All noble gases except helium have a similar octet of electrons in their outermost shell. This configuration is exceptionally stable. Responsible for the unreactive nature of the noble gases. Main group elements that ionize easily generally do so in a way that gives them the same octet of electrons. “n”s “n”p where n = period number

Orbital Diagrams Example: Draw an orbital diagram for each of the following atoms or ions. Iron: Bromine: Sodium ion:

Useful Information from the Periodic Table The period number of the element indicates the highest shell (value of n) that contains electrons for that atom. An element in the fourth period will have one or more electrons in the n=4 shell. The location of the atom in the periodic table indicates the subshell where the last e-’s are found. s block d block f block p block

Electron Configuration A short-hand notation (electron configuration) is commonly used instead of an orbital diagram. The electron configuration designates: Each subshell that contains electrons in order of increasing energy The number of electrons found in the subshell

Electron Configuration The orbital diagram for an oxygen atom: The electron configuration for an oxygen atom: 1s22s22p4 Notice: no commas between subshells! 1s 2s 2p

Electron Configuration Process for writing an electron configuration: Determine the number of electrons present Add electrons to each subshell in order of increasing energy until all electrons have been designated Use diagonal diagram Remember the maximum # of e- per subshell: _s2 _p6 _d10 _f14

Electron Configuration Example: Write the electron configuration for each of the following atoms: Titanium Lead

Electron Configuration Example: Write the electron configuration for each of the following ions: Oxide ion: Potassium ion:

Electron Configuration Using Core Notation Calcium atoms contain 20 electrons: The first 18 electrons are arranged exactly as the electrons present in argon. The argon core: [Ar] The last two electrons are referred to as valence electrons. Electrons located in the outermost shell that can be transferred to or shared with another atom during the formation of ions or covalent bonds Electrons over and above those found in the previous noble gas

Electron Configuration Using Core Notation Argon: 1s22s22p63s23p6 Calcium: 1s22s22p63s23p64s2 Electron configuration using core notation: [Ar]4s2 Valence e- [Ar]

Electron Configuration Using Core Notation The electron configuration using core notation contains two components: the noble gas core valence electrons Fe: [Ar]4s23d6 C: [He]2s22p2

Electron Configuration Using Core Notation To write the electron configuration using core notation: Find the noble gas that comes before the atom and place its elemental symbol in [ ] Calculate the number of additional electrons Atomic # of atom – atomic # noble gas Determine the period number “n” of the atom and begin placing valence electrons in the “n”s subshell. Use diagonal diagram to determine order in which subsequent subshells are filled

Electron Configuration Using Core Notation Example: Write the electron configuration using core notation for each of the following atoms. Ni: Bi:

Electron Configuration Using Core Notation Example: Write the electron configuration using core notation for each of the following ions. Iodide ion: Magnesium ion:

Transition Metal Ions Transition metal ions form when electrons are lost from the parent atom in the following order: s electrons from outermost shell first d electrons from previous shell next Example: Ti: [Ar]4s23d2 Ti3+: [Ar]3d1

Anomalies Some irregularities occur when there are enough electrons to half-fill s and d orbitals on a given row.

Anomalies For instance, the core electron configuration for chromium is [Ar] 4s1 3d5 rather than the expected [Ar] 4s2 3d4. The core electron configuration for copper is [Ar]4s13d10 instead of [Ar]4s23d9.

Isoelectronic Series The following ions contain the same number of electrons (10) as Ne. These ions are isoelectronic with each other and neon. Having the same number of electrons These ions and neon form an isoelectronic series. A group of atoms and ions with the same number of electrons Nitride ion Oxide ion Fluoride ion Sodium ion Magnesium ion Aluminum ion

Isoelectronic Series Example: Which of the following atoms or ions in each group are isoelectronic? Fe2+, Co3+, Mn, Cr Se2-, Br, Kr, Sr2+

Periodic Properties of Elements Chemical and physical properties of the elements vary with their position in the periodic table. Atomic size Size of Atom vs. Ion Size of Ions in Isoelectronic series Ionization energy Electron affinity Metallic character

Periodic Properties--Atomic Size The relative size (radius) of an atom of an element can be predicted by its position in the periodic table. Trends Within a group (column), the atomic radius tends to increase from top to bottom Within a period (row), the atomic radius tends to decrease as we move from left to right

Periodic Properties--Atomic Size Periodic Table Increasing size Increasing size Lower “lefter” larger

Periodic Properties – Atom vs. Ion Size Trends to know: Cations (+) are smaller than their parent atoms. Electrons are removed from the outer shell. Anions (-) are larger than their parent atoms. Electron-electron repulsion causes the electrons to spread out more in space.

Periodic Properties – Ion Size Trends to know: For ions in the same group (same charge), size increases from top to bottom. Same trend as for the size of parent atoms I- is larger than F- For an isoelectronic series of ions, the size decreases with increasing atomic number. Na+ is smaller than O2-

Periodic Properties - Ionization Energy The ease with which an electron can be removed from an atom to form an ion is an important indicator of its chemical behavior. Ionization energy: the minimum energy required to remove an electron from the ground state of an isolated gaseous atom or ion. Formation of cation (+) or more positively charged cation Na (g)  Na+ (g) + e-

Periodic Properties - Ionization Energy As ionization energy increases it becomes harder to remove an electron/form a cation. Within each row, the ionization energy increases from left to right. Metals form cations more easily than nonmetals. Within each group, the ionization energy generally decreases from top to bottom. It’s easier to form K+ than Li+.

Periodic Properties – Electron Affinity The energy change that occurs when an electron is added to a gaseous atom is called the electron affinity. Cl (g) + e-  Cl- (g) The electron affinity becomes increasingly negative as the attraction between an atom and an electron increases more negative electron affinity = more likely to gain an electron and form an anion

Periodic Properties – Electron Affinity Trends: Halogens have the most negative electron affinities. Electron affinities become increasing negative moving from the left toward the halogens. Electron affinities do not change significantly within a group. Noble gases will not accept another electron.

Periodic Properties – Metallic Character Metals: shiny luster malleable and ductile good conductors of heat and electricity form cations Metallic character increases from top to bottom Increases from right to left