Atomic+Structure+and+Orbitals

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=Bohr-Rutherford Model of the Atom = Experiments carried out by Ernest Rutherford showed that the atom was composed of a very small, massive, positively charged nucleus surrounded by negative electrons. Later the Danish scientist Niels Bohr modified the model and showed that electrons can only occupy orbits with certain allowable (“quantised”) energy levels. Atomic orbitals define regions of space in which there is a high probability of finding an electron. Each orbital has a particular shape and associated energy values. In the normal ground state of an atom, the electrons occupy orbitals with the lowest possible energies. On heating, the electrons can be excited to orbitals with higher energy - the ‘excited state’. As the electrons fall back to lower energy levels (orbitals) they will emit electromagnetic radiation, which may be in the region of visible light ie. it appears coloured. Each element has its own characteristic emission spectrum that can be used to identify that element - its “chemical fingerprint”. These spectra are used to identify the elemental composition of distant stars.

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The position and energy of any electron in an atom is identified using a set of four quantum numbers. 1) The first quantum number identifies which main **energy level** (“shell”) the electron is in. The first 2 electrons occupy the 1st energy level (1st row of the periodic table, H and He), then the next 8 electrons fill the 2nd energy level (2nd row of the periodic table Li to Ne). Following this electrons go into level 3 and then level 4 etc. 2) The second quantum number identifies the **sub-level** which are classified as either //s, p, d// (or //f//).  The first level has only 1 sub-level - **1//s//**  The second level (or row) has 2 sublevels - **2//s// and 2//p//.**  The 3rd level has 3 sub-levels **- 3//s//, 3//p// and 3//d//** 3) The third quantum number identifies the specific orbital (region of space) that the electron occupies. **Each orbital can accommodate a maximum of 2 electrons**. The number of orbitals in any particular sublevel is always the same
 * Energy levels, Sub-levels and Orbitals **
 * //s // sublevels only have **one** orbital and therefore have a maximum of 2 electrons
 * //p //<span style="font-family: Arial,sans-serif; font-size: 11pt;"> sublevels contain **3** orbitals and therefore contain a maximum of 3 x 2 = 6 electrons


 * //<span style="font-family: Arial,sans-serif; font-size: 11pt;">d //<span style="font-family: Arial,sans-serif; font-size: 11pt;"> sublevels contain **5** orbitals and therefore contain a maximum of 5 x 2 = 10 electrons


 * //<span style="font-family: Arial,sans-serif; font-size: 11pt;">f //<span style="font-family: Arial,sans-serif; font-size: 11pt;"> sublevels contain **7** orbitals - a maximum of 14 electrons. (Available from the 4th energy level)

<span style="font-family: Arial,sans-serif; font-size: 11pt;">4) The fourth quantum number uniquely identifies each of the two electrons in a specific orbital by defining its spin as either up or down.

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<span style="font-family: Arial,sans-serif; font-size: 11pt;">When placing electrons in orbitals they always fill up the sub-levels of lowest energy first. <span style="font-family: Arial,sans-serif; font-size: 11pt;">The order of energies of sublevels is 1//s// 2//s// 2//p// 3//s// 3//p// 4//s// 3//d// 4//p// and this is related to the periodic <span style="font-family: Arial,sans-serif; font-size: 11pt;">table as follows.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Aufbau Principle **



<span style="font-family: Arial,sans-serif; font-size: 11pt;"> “//s//” block “//d//” block “//p//” block <span style="font-family: Arial,sans-serif; font-size: 11pt;">Thus atoms which are in the first 2 groups (columns) of the periodic table have either 1 or 2 valence (outer shell) electrons in an //s// sub-shell. Atoms found in groups 3 to 12 will have 1 to 10 electrons in the //d// sublevel. These elements are called transition metals. Atoms in groups 13 to 18 will have from 1 to 6 valence electrons in a //p// sub-shell.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">The third level of the periodic table “should” contain a total of 18 elements (3s23//p//63//d//10) but actually only contains 8 elements. The “missing” 10 elements appear slightly later as the transition elements because electrons placed in the 4//s// orbital are at a lower energy level than those in the 3//d// sublevel and the //4s// orbital is therefore filled before any electrons go into the 3//d// orbitals.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">NOTE: **

<span style="font-family: Arial,sans-serif; font-size: 11pt;">The electron configuration for any atom is a statement giving the number of electrons at each sublevel. Each sub-level is filled, starting with the one of lowest energy, before proceeding to the <span style="font-family: Arial,sans-serif; font-size: 11pt;">next sublevel. <span style="font-family: Arial,sans-serif; font-size: 11pt;">eg Li (atomic number 3) 1//s//2 2//s//1 (total 3 electrons) <span style="font-family: Arial,sans-serif; font-size: 11pt;"> F (at no 9) 1//s//2 2//s//2 2//p//5 (total 9 electrons) <span style="font-family: Arial,sans-serif; font-size: 11pt;"> K (at no 19) 1//s//2 2//s//2 2//p//6 3//s//2 3//p//6 4//s//1 (19 electrons) <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Fe (at no 26) 1//s//2 2//s//2 2//p//6 3//s//2 3//p//6 3//d//6 4//s//2 (26 electrons) <span style="font-family: Arial,sans-serif; font-size: 11pt;">In Fe (like all of the transition metals) the 4//s// orbital (which is filled first) is full but the 3//d// <span style="font-family: Arial,sans-serif; font-size: 11pt;">sublevel is only partially filled.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Electron configuration **
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">NOTE: **

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Because all orbitals within the same sub-level have the same energy, they are filled so thatas many electrons remain unpaired as possible as this reduces the repulsion resulting from electrons occupying the same region of space. <span style="font-family: Arial,sans-serif; font-size: 11pt;">In the same way that atoms react to try and achieve completely full energy levels, a half-filled sub-level is also a particularly stable (or low energy) state. This explains the unusual electron configurations for the transition metals Cr and Cu, where in each case an electron is promoted from <span style="font-family: Arial,sans-serif; font-size: 11pt;">the slightly lower 4//s// sub-level to the 3//d// sub-level to achieve half-filled and full sub-levels respectively. <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Cr 1//s//2 2//s//2 2//p//6 3//s//2 3//p//63//d//54//s//1 **OR** [Ar] 3//d//54//s//1 <span style="font-family: Arial,sans-serif; font-size: 11pt;">Cu 1//s//2 2//s//2 2//p//6 3//s//2 3//p//63//d//104//s//1 **OR** [Ar]3//d//104//s//1 <span style="font-family: Arial,sans-serif; font-size: 11pt;">This stability also explains the small “kinks” in the graph of ionisation energy vs atomic number. <span style="font-family: Arial,sans-serif; font-size: 11pt;">e.g. the 1st IE of sulphur is lower than the 1st IE of phosphorus despite sulfur’s higher nuclear charge. This is because the resulting S+ ion has a stable half filled //p// sub-level whereas the ionistaion of phosphorus requires the loss of an electron from the half filled sub-level.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Stability of half-filled sub-levels **

=<span style="font-family: Arial,sans-serif; font-size: 11pt;">Electron Configurations of Monatomic Ions = <span style="font-family: Arial,sans-serif; font-size: 11pt;">a) When non-metal atoms (found in Groups 13 – 17) gain electrons and form anions the electrons are added to the outermost, //p// sub-level. The atom will form a negative ion having the same electron configuration as the inert gas at the end of the same row of the periodic table. <span style="font-family: Arial,sans-serif; font-size: 11pt;">This means P3–, S2–, Cl– and Ar will all have the same electron arrangement of 1//s//2 2//s//2 2//p//6 3//s//2 3//p//6. They are all **isoelectronic**.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">b) When metal atoms lose electrons they are removed from the sub-level furthest from the nucleus. This means Na loses its 3//s//1 electron to form Na+ which has the same electron configuration as Ne.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Atoms in Groups 1 and 2 only lose the electrons from their outer s orbital and form either M+ or M2+ ions. For the atoms in the first transition series, both the 3//d// and 4//s// electrons are part of the valence shell and can be removed to form transition metal cations. <span style="font-family: Arial,sans-serif; font-size: 11pt;">Because of this they can form species with a variety of oxidation states. <span style="font-family: Arial,sans-serif; font-size: 11pt;">e.g. Fe 1//s//2 2//s//2 2//p//6 3//s//2 3//p//6 3//d//6 4//s//2 Fe2+ 1//s//2 2//s//2 2//p//6 3//s//2 3//p//6 3//d//6 **NOT** [Ar] 3//d//4 4//s//2
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">NOTES: **
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">When forming a cation the metal atoms only lose electrons from their valence shell.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Although the 4//s// sublevel is filled before the 3//d//, when forming a cation the transition metals always lose **the 4//s// electrons** **before losing electrons from the 3//d// sublevel.** This is because on average, the 3//d// sub-level is closer to the nucleus than the 4//s// sub-level despite it having a slightly higher energy level.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Fe2+ can be further oxidised to form the Fe3+ ion because of the xtra stability arising from its half filled 3//d//5 sub-level

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