The Chemistry of the Elements PDF

CHM577: INORGANIC CHEMISTRY Chapter Outline The Chemistry of the Elements § To know the periodic classification of the elements §...

CHM577: INORGANIC CHEMISTRY Chapter Outline The Chemistry of the Elements § To know the periodic classification of the elements § Able to construct its electron configuration § Identify the periodic variation in physical properties: — Effective nuclear charge (Zeff) DR. AMALINA MOHD TAJUDDIN — Atomic & ionic radii — Ionization Energy (IE) — Electron Affinity (EA) 2 Introduction……. PERIODIC CLASSIFICATION OF THE ELEMENTS ns2np6 ns1 ns2np1 ns2np2 ns2np3 ns2np4 ns2np5 ns2 n = principal quantum number of the outermost subshell d5 d1 d10 Dmitri Mendeleev- Work on periodic classification of elements according to their properties.. -Most significant achievement in 19th century.. 4f 5f 4 Ground State Electron Configurations of the Elements 3 1  Classification of the Elements Continued……. Group 1A à 7A (known as representative elements/main group elements) Have incompletely filled s or p subshells Group 8A (except He) à Show completely filled p subshells. Example:- He 1s2 and ns2np6 for other noble gases (Ne: 1s2 2s2 2p6) Group 3B à 8B Known as d-block transition elements. Have incompletely filled d subshells. àThey will produce cations with these incompletely filled d subshells. Group 4F & 5F à Known as f block transition elements. Have incompletely filled f subshells Lanthanides Actinides 6 5 In the ground state of the atom, electrons will occupy the lowest energy orbitals first, and only fill Ground state electronic configurations the higher energy orbitals when no lower energy orbitals are left. (it is an electronic arrangement described for each atom) Hund’s first rule:- electrons occupy Fig. Order for filling all the orbitals of a given subshell energy sublevels The Aufbau Principle singly before pairing begins. with electrons. These unpaired electrons have Aufbau means ‘building up’ parallel spins. Used together with Hund’s rules and Pauli exclusion principle Pauli exclusion principle:- no two electrons in the same atom may have the same set of n, l, ml, ms quantum numbers 7 8 2  Electron Configurations of Cations and Anions of Representative Elements Condensed ground-state electron configurations in the first three periods Na [Ne]3s1 Na+ [Ne] Atoms lose electrons so that cation has a stable Ca [Ar]4s2 Ca2+ [Ar] noble-gas outer electron configuration. Al [Ne]3s23p1 Al3+ [Ne] H 1s1 H- 1s2 or [He] Atoms gain electrons so that anion has a F 1s22s22p5 F- 1s22s22p6 or [Ne] stable noble-gas outer electron configuration. O 1s22s22p4 O2- 1s22s22p6 or [Ne] N 1s22s22p3 N3- 1s22s22p6 or [Ne] 9 10 Cations and Anions of Representative Elements Isoelectronic… +1 +2 +3 -3 -2 -1 They have the same number of electrons and ground state electron configuration Examples:- Na+: [Ne] F-: 1s22s22p6 or [Ne] O2-: 1s22s22p6 or [Ne] Al3+: [Ne] N3-: 1s22s22p6 or [Ne] Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne Quiz : What neutral atom is isoelectronic with H- ? Answer: H-: 1s2 same electron configuration as He 11 12 3  Electron Configurations of Cations of Transition Metals ALL Periodic Table Trends When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns Influenced by three factors: orbital and then from the (n – 1)d orbitals. 1. Energy Level — Higher energy levels are further away from the Examples:- nucleus. Fe: [Ar]4s23d6 Mn: [Ar]4s23d5 Fe2+: [Ar]4s03d6 or [Ar]3d6 Mn2+: [Ar]4s03d5 or [Ar]3d5 ……1 2. Charge on nucleus (# protons) ….. 2 — More charge pulls electrons in closer. (+ and – attract Fe3+: [Ar]4s03d5 or [Ar]3d5 Mn2+: [Ar]4s23d3 each other) 3d orbital is more stable than the 4s orbital in transition metal ions.. 3. Shielding effect 13 14 Shielding Effect Atomic Size — The electron on the outermost energy level has to look through all the other energy levels to see the nucleus. } — Second electron has same Radius shielding, if it is in the same — Measure the Atomic Radius period Ø this is half the distance between the two nuclei of a diatomic molecule. 15 16 4  Periodic Variation in Physical Properties What do they influence? A. Effective nuclear charge (Zeff) ü It is the “positive charge” felt by an electron… ü Given by Zeff = Z – s, Z = actual nuclear charge, ØEnergy levels and Shielding have an effect s = shielding/screening constant on the GROUP (0 < s < Z) Zeff » Z – number of inner or core electrons Z Core Zeff Radius ØNuclear charge has an effect on a PERIOD Na 11 10 1 186 Mg 12 10 2 160 Al 13 10 3 143 17 Si 14 10 4 132 18 increasing Zeff Example: Li Be B C N O F Ne increasing Zeff 3 4 Core electron = 1s2 5 6 7 8 9 10 Which element’s outer shell or “valence” electrons is predicted to have the largest Effective nuclear charge? Cl, O, N or Ca? Cl: Zeff ≈ 17 - 10 = 7 O: Zeff ≈ 8 - 2 = 6 1. Core electron > closer to nucleus than valence electron, thus core e shield valence e> than valence e shield each other. 2. Moving across the period, core e remains constant , but Z increases. N: Zeff ≈ 7 - 2 = 5 3. The added e will be valence e, and due to valence e does not shield each other, thus, moving across the period, > Zeff will be felt by valence e. Ca: Zeff ≈ 20 - 18 = 2 4. Moving the group, Zeff. As n increases, large shells increases, thus valence e are added to these large shells. Thus, electrostatic attraction between nucleus & valence e decreases. 19 20 5  Atomic Size - Group trends #1. Atomic Size - Period Trends H — Decreases across a period from left to right — Increases down a given group. Li — Electrons are in the same energy level and unshielded — the number of electrons and towards attraction by protons. filled electron shells Na ¡ protons are being added to the nucleus thus creates increases a "higher effective nuclear charge." — electrons are found further — stronger force of attraction pulling the electrons from the nucleus K closer to the nucleus — Therefore, the atomic radii increase. Rb 21 22 Na Mg Al Si P S Cl Ar Atomic Radii Ionic Radii — Anions (negative ions) are larger than their respective atoms. WHY? — Electron-electron repulsion forces them to spread further apart. — The protons cannot pull the extra electrons as tightly toward the nucleus. — Cations (positive ions) are smaller than their respective atoms. WHY? — There is less electron-electron repulsion. — Protons outnumber electrons; the protons can pull the fewer 23 24 electrons toward the nucleus more tightly. 6  Ion Group Trends Ion Period Trends — Each step down a group — Across the period from left to right, the nuclear charge is adding an energy Li1+ increases - so they get smaller. level. Na1+ — Notice the energy level changes between anions and K1+ cations. — Ions therefore get bigger 1+ as you go down, Rb N3- because of the additional Li1+ O2- F1- energy level Cs 1+ B3+ Be2+ C4+ 25 26 QUIZ: Arrange the following atoms in order of Size of Isoelectronic ions increasing atomic radius P, Si, N — Iso- means “the same” — Isoelectronic ions have the same no. of electrons STRATEGY: From left to right across period—Decreases — Al3+ Mg2+ Na1+ Ne F1- O2- and N3- Moving up to down the group - Increases — all have 10 electrons — all have the same configuration: 1s22s22p6 ANSWER: N Xp1 N F — Example : 4Be > 5B — The energy of an electron in an Xp orbital is greater than Xs orbital. H C O — Less energy to remove the first electron in a p orbital than it is Be to remove one from a filled s orbital. 2) Xp3 > Xp4 B — Example : 7N > 8O Li — After the separate degenerate orbitals have been filled with single electrons, the fourth electron must be paired. The Na electron-electron repulsion makes it easier to remove the outermost, paired electron. (Hund's Rule) 33 Atomic number 34 Second and Higher Ionization Energies Ionization Energy (IE) Symbol First Second Third — Definition: Second Ionization Energy is the energy required to remove a second outermost electron from H 1312 a ground state atom. He 2731 5247 Li 520 7297 11,810 — Subsequent ionization energies increase greatly once Be 900 1757 14,840 an ion has reached the state like that of a noble gas. B 800 2430 3569 C 1086 2352 4619 — For elements that reach a filled or half-filled orbital by N 1402 2857 4577 removing 2 electrons, 2nd IE is lower than expected. O 1314 3391 5301 (True for s2 ) F 1681 3375 6045 Ne 2080 3963 6276 35 36 9  Electron Affinity (EA) What factors determine IE Electron affinity is the negative of the energy change — The greater the nuclear charge, the greater IE. that occurs when an electron is accepted by an atom in the gaseous state to form an anion. — Greater distance from nucleus decreases IE X (g) + e- X-(g) — Filled and half-filled orbitals have lower energy, so achieving them is easier, lower IE. F (g) + e- F-(g) DH = -328 kJ/mol EA = +328 kJ/mol O (g) + e- O-(g) DH = -141 kJ/mol EA = +141 kJ/mol — Shielding effect 37 38 Electron Affinity Exceptions on electron affinity trends — Definition: The energy given off when a neutral atom — Nonmetals elements in the first period have lower electron in the gas phase gains an extra electron to form a affinities than the elements below them in their respective negatively charged ion. groups. H — Elements with electron configurations of Xs2, Xp3, and Xp6 — 1) down a group, electron affinity decreases. Li have electron affinities less than zero because they are Na unusually stable. e.g. Be, N, Ne K — WHY? - Electron affinities are all much smaller than — 2) across a period, electron affinity increases. ionization energies. — Xs2 < 0: Stable, diamagnetic atom with no unpaired electrons. Li Be B C N O F — Xp3 < 0: Stable atom with 3 unpaired p-orbital electrons each occupying its own subshell. — Xp6 < 0: Stable atom with filled valence (outermost) shell. 39 40 10  Group 7-the highest ē affinity. Group 8-the lowest (zero or –ve) value…Y???? ē affinity of Group 2A< 1A and 5A