Periodic Table and Properties of Elements Lecture - Kwame Nkrumah University PDF

Kwame Nkrumah University of Science & Technology, Kumasi, Ghana LECTURE THREE Periodic Table and Periodic Properties of Elements  The Periodic Table Mendeleev put forward the modern form of the periodic table. All elements are arranged on the periodic table...

Kwame Nkrumah University of Science & Technology, Kumasi, Ghana LECTURE THREE Periodic Table and Periodic Properties of Elements  The Periodic Table Mendeleev put forward the modern form of the periodic table. All elements are arranged on the periodic table according to their atomic numbers. The elements can be divided into groups, periods and blocks. Groups This is a vertical arrangement of the elements on the periodic table. Elements in a group have the same valence electrons or similar atomic configurations. www.knust.edu.gh 3Li= [He] 2s1, 11 Na = [Ne] 3s1, 19 K= [Ar] 4s1 11 Na → 1s2 2s2 2p6 3s1 19 K → 1s2 2s2 2p6 3s2 3p6 4s1 The elements have 1 valence electron in their s- orbital. Similarly, the atomic configuration has only one electron outside a noble gas configuration. www.knust.edu.gh Group elements have similar chemical properties. In most elements the number of valence electrons is the same as the group number, where the group members share similar physical and chemical properties. For example, they could be considered as “a family”. Example: Alkali metals consisting of Li, Na, K, Rb, Cs Halogen can also form a family: F, Cl, Br, I There is an isolated situation i.e. Group 4 or 14 are not family C, Si This is because elements in the same group have different n-numbers www.knust.edu.gh www.knust.edu.gh www.knust.edu.gh Periods This is the horizontal arrangement of elements on the periodic table. One element to the next as the atomic number increases by 1. There are 7 periods, i.e.. Length of the period is not uniform but varies. The shortest period has 2 elements and the longest period has 32 elements. www.knust.edu.gh The period number refers to the main energy level that the valence electron occupies. P1 (1s) - total number of elements 2 P2 (2s, 2p) - total number of elements is 8 (i.e.. 2+6= 8). P3 (3s, 3p) - total number of elements is 8 P4 (4s, 3d, 4p) - total number of elements is18 P5 (5s, 4d, 5p) - total number of elements is18 P6 (6s, 4f, 5d, 6p) - total number of elements is32 www.knust.edu.gh For clumsy arrangement, the elements with their valence electron filling the 4f- orbitals are placed in a separate row at the bottom of the periodic table. The row forms Lanthanides, where it starts from Lanthanum, La (57) - Lutetium, Lu (71). Similar situation is observed for valence electrons filling 5f orbitals which are placed below the Lanthanides. These are the Actinides, it starts from Ac (Actinium) 89 - Lr (Lawrencium) 103. www.knust.edu.gh Blocks These are groups of 5 B → 1s2 2s2 2p1 elements with their valence 2p2 6 C → 1s 2 2s 2 electrons occupying the N → [He] 2s2 2p3 same orbital. 7 8 O → [He] 2s2 2p4 Valence electrons in s- F → [He] 2s 2 2p5 orbitals for both G1 and G2 9 elements are therefore said N to be s-blocks elements. 10 e → [He] 2s2 2p6 Second period elements starting from Gp3 - Gp0 elements. i.e.. B, C, N, O, F, Ne www.knust.edu.gh www.knust.edu.gh From group 3 to group 0 elements, all the valence electrons are found in the p-orbital. The group 3 to group 0 elements are therefore called p-block elements. d- block elements (transition elements) valence electrons filled d-orbitals. 4th period Sc - Zn : 1st transition series 5th period Y - Cd; 2nd transition series. 6th period L - Hg: 3rd transition series F block elements, valence electrons filling the f-orbital i.e.. Lanthanides and Actinides. NB: the s and p blocks are the main group elements. www.knust.edu.gh  Periodic Properties These are physical and chemical properties of atoms which reoccur on the periodic table. Examples include: o Atomic/Ionic radius o Ionization energy o Electron affinity o Electronegativity www.knust.edu.gh  To investigate the above properties, the following terms must be understood. o Valence electrons: The outermost electrons that are involved in bond formation. o Core electrons: These are the innermost electrons. o Effective nuclear charge: The net attraction of protons in the nucleus for the valence electrons. www.knust.edu.gh o Shielding effect: The decreased attraction of protons for valence electrons by the core electrons. Across a period, there is an increase in effective nuclear charge. Why? Down a group, there is an increase in the shielding effect Why? www.knust.edu.gh  Atomic radius Illustration 1/2d=r The size of an atom is dependent on how strongly the protons in the nucleus attract the valence electrons. Across a period, screening effect due to electrons in the same shell almost remain constant but increase in the atomic number, even by 1 is very significant. Therefore, the greater the attraction for valence electrons, the smaller the size of the atom. NB: From Halogen to rare gases, there is an increase in atomic size due to repulsion among the electrons in the outer shell. www.knust.edu.gh  Ionic radius Cation: Li --- Li+ (He) Na ----- Na+ (Ne) They are isoelectronic with the noble gas. The cations are smaller in size than their parent atoms. Why? The reasons are that they have fewer electrons than protons, which leads to increase in effective nuclear charge. Cationic radius decreases across a period but increases down a group. Anion: The anions are larger in size than their parent atoms. Why? Because there are more electrons than protons and decrease in effective nuclear charge. Anion radius increases across the period and www.knust.edu.gh decreases down the group.  Ionization energy (Ei) This is the energy required to remove the most loosely bound electrons from a gaseous atom. The determining factors include:  Atomic radius  Shielding effect  Stability of electronic configuration  Penetrative effect (effective nuclear charge) www.knust.edu.gh Down a group, there is a decrease in ionization energy. Why? Because there is increase in atomic radius, increase in screening effect and decrease in penetrative effect. Therefore, less energy is required to remove valence electrons. Across the period, there is increase in ionization energy (Ei). Why? Because there is decrease in atomic radius, screening effect remain almost the same and increase in penetrative effect, therefore high energy is required to remove valence electrons. www.knust.edu.gh Discrepancy: (Best explained using stability of electronic Be and 5B 4 Be= 1s2 2s2 > 5B= 1s2 2s2 2p1 4 2p electrons are at higher energy level than 2s electrons. The higher the energy level the less stable the electron, thus making the removal of electrons from 2p much easier. Why? The s-electrons are also closer to the nucleus than the p- electrons therefore the 2s electrons have a greater penetrative power (effective nuclear charge) than the 2p electrons. In totality, there is greater stability with an outermost orbital www.knust.edu.gh which is fully filled as in Be. 7 N and 8O 7 N = 1s2, 2s2, 2p3 > 8O = 1s2, 2s2, 2p4 2p orbitals in N is half filled implying high stability as single electrons in p- orbital are spread out reducing repulsion. There is lower ionization energy of oxygen due to instability resulting from repulsion between the two electrons which become paired up in the p – orbital. To remove 2nd, 3rd electrons in an atom, the input of 2nd and 3rd ionization energy will be increasing. Where a noble gas electronic configuration is involved, an unusual high ionization energy is required because of stability of the configuration. www.knust.edu.gh Electron affinity Energy (Eea) It is the energy change when an electron is added to a gaseous atom. Examples: Cl (g) + e- ------- Cl-(g) Eea= - 348 KJ/mol O (g) + e- ------- O-(g) Eea= - 141 KJ/mol O-(g) + e- -------O2- (g) Eea= www.knust.edu.gh +798 KJ/mol The 2nd electron in the oxygen was added against the repulsive force of negatively charged oxygen. Therefore, the need for energy input leading to positive electron affinity value. Down the group, there is a decrease in electron affinity. Why? Because there is decrease attraction for added electrons, increase atomic size and decrease effective nuclear charge. www.knust.edu.gh The trend is clear with the more electronegative element. Across the same period, there is increase in electron affinity. The reason is the opposite. Discrepancy: Unexpected lower negative electron affinity of fluorine (F) which is caused by its very small size which created very strong repulsive field for any incoming electron. www.knust.edu.gh Factors that affect Electron affinity include:  Atomic radius  Stability of electronic configuration  Charge carried by anion For unexpected values, this may be due to more stable electronic configuration and repulsive forces of electrons. Discrepancy on period 2 B + e- -----B- Eea = -27 KJ/mol B = 1s2 2s2 2p1 www.knust.edu.gh C + e- -----C- Eea = -122 KJ/mol C = 1s2 2s2 2p2 N + e- -----N- Eea = 0 KJ /mol N = 1s2 2s2 2p3 www.knust.edu.gh  Electronegativity This is the ability of an electron in a molecule to attract bonded/shared electrons to itself. An atom of a molecule which tends to attract more electrons to itself is described as more electronegative atom. The main determinant of electronegativity is atomic radius. The smaller the atomic radius, the higher its electronegativity. Let’s discuss NB: For gaseous isolated atoms, the electron affinity is measured directly. However, for electronegativity, the atoms in the molecule are measured indirectly by a comparative scale. www.knust.edu.gh The scale gives a quantitative measure of an atom’s ability to attract electrons. There is no unit for electronegativity. Values are negative and given accordingly to Pauling’s scale. On that scale, the highest value is 4.0 which is fluorine, while the lowest value is francium, 0.7. www.knust.edu.gh Most electronegative elements are reactive non- metals (Halogens). The least electronegative elements are reactive metals (Alkali). NB: Standard enthalpy of formation is the energy change when one mole of a compound is formed from elements in their standard states. www.knust.edu.gh