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General and Inorganic Chemistry I. Lecture 1 István Szalai Eötvös University István Szalai (Eötvös University) Lecture 1 1 / 40 Lecture 1 2 / 40 Outline 1 Periodic Table István Szalai (Eötvös University) Periodic Table Electron Configurations The electrons occupy the orbitals in the way that gives the lowest energy for the atom. Pauli Exclusion Principle: No two electrons in an atom may have identical sets of four quantum numbers. Hund’s Rule: Electrons occupy all the orbitals of a given subshell singly before pairing begins. These unpaired electrons have parallel spins. Electrons are assigned to orbitals in order of increasing value of (n + l). István Szalai (Eötvös University) Lecture 1 3 / 40 Periodic Table Electron Configurations Li 1s 2 2s 1 C 1s 2 2s 2 2p 2 Fe 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 Cu 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 Ce 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 1 4f 1 Valence shell: the electrons in the outer shell, those that were not present in the precding noble gas orbitals. István Szalai (Eötvös University) Lecture 1 4 / 40 Periodic Table Periodic Table István Szalai (Eötvös University) Lecture 1 5 / 40 Periodic Table The properties of the elements are periodic functions of their atomic number 2 Li + 2 H2 O → 2 Li+ + 2 OH− + H2 2 Na + 2 H2 O → 2 Na+ + 2 OH− + H2 2 K + 2 H2 O → 2 K+ + 2 OH− + H2 István Szalai (Eötvös University) Lecture 1 6 / 40 Periodic Table History of the periodic table About 330 B.C Aristotle proposed that everything is made up of a mixture of one or more of four ”roots”. The four elements were earth, water, air and fire. In 1661 Boyle defined an element as a substance that cannot be broken down into a simpler substance by a chemical reaction. Lavoisier published a list of elements in 1789, or substances that could not be broken down further, which included oxygen, nitrogen, hydrogen, phosphorus, mercury, zinc, and sulfur. Lavoisier’s descriptions only classified elements as metals or non-metals. In 1817, Johann Wolfgang Döbereiner began to formulate one of the earliest attempts to classify the elements. He found that some elements formed groups of three with related properties. He termed these groups ”triads”. (Cl-Br-I, Ca-Sr-Ba, S-Se-Te. . . ) István Szalai (Eötvös University) Lecture 1 7 / 40 Periodic Table History of the periodic table Alexandre-Emile Béguyer de Chancourtois, a French geologist, was the first person to notice the periodicity of the elements — similar elements seem to occur at regular intervals when they are ordered by their atomic weights. He devised an early form of periodic table, which he called the telluric helix. With the elements arranged in a spiral on a cylinder by order of increasing atomic weight, de Chancourtois saw that elements with similar properties lined up vertically. John Newlands was an English chemist who in 1865 classified the 56 elements that had been discovered at the time into 11 groups which were based on similar physical properties. István Szalai (Eötvös University) Lecture 1 8 / 40 Periodic Table Dmitri Mendeleev (1837-1907) Dmitri Mendeleev, a Siberian-born Russian chemist, was the first scientist to make a periodic table much like the one we use today. His table was published in 1869. It stated: The elements, if arranged according to their atomic weights, exhibit an apparent periodicity of properties. Elements which are similar as regards to their chemical properties have atomic weights which are either of nearly the same value (e.g., Pt, Ir, Os) or which increase regularly (e.g., K, Rb, Cs). István Szalai (Eötvös University) Lecture 1 9 / 40 Periodic Table Dmitri Mendeleev (1837-1907) The arrangement of the elements, or of groups of elements in the order of their atomic weights, corresponds to their so-called valencies, as well as, to some extent, to their distinctive chemical properties; as is apparent among other series in that of Li, Be, Ba, C, N, O, and Sn. We must expect the discovery of many yet unknown elements–for example, elements analogous to aluminium and silicon–whose atomic weight would be between 65 and 75. István Szalai (Eötvös University) Lecture 1 10 / 40 Periodic Table Germanium Properties Atomic weight Density Oxide Chloride István Szalai (Eötvös University) Mendeleev prediction 72 5,5 g/cm3 EO2 ECl4 Lecture 1 Observed values 72,6 5,35 g/cm3 GeO2 GeCl4 11 / 40 Periodic Table Refinements to the periodic table In 1914 Henry Moseley found a relationship between an element’s X-ray wavelength and its atomic number (Z), and therefore resequenced the table by nuclear charge rather than atomic weight. Thus Moseley placed argon (Z=18) before potassium (Z=19) based on their X-ray wavelengths, despite the fact that argon has a greater atomic weight (39.9) than potassium (39.1). The new order agrees with the chemical properties of these elements, since argon is a noble gas and potassium an alkali metal. In 1945, Glenn T. Seaborg proposed a significant change to Mendeleev’s table: the actinide series. Seaborg’s actinide concept of heavy element electronic structure, predicting that the actinides form a transition series analogous to the rare earth series of lanthanide elements. István Szalai (Eötvös University) Lecture 1 12 / 40 Periodic Table Classification of Elements Name Representative Elements Alkali Metals Alkaline Earth Metals Boron group Carbon group Nitrogen group Oxygen group Halogens Noble Gases Transition Metals Lanthanides and Acthinides István Szalai (Eötvös University) Electron structure ns 1 ns 2 ns 2 np 1 ns 2 np 2 ns 2 np 3 ns 2 np 4 ns 2 np 1 ns 2 np 6 ns 2 (n − 1)d 1−10 ns 2 (n − 1)d 1 (n − 2)f 1−14 Lecture 1 13 / 40 Periodic Table Periodic Properties atomic radii, ionic radii ionization energy electron affinity electronegativity periodic chemical properties István Szalai (Eötvös University) Lecture 1 14 / 40 Periodic Table Atomic radii Covalent radius: the nominal radius of the atoms of an element when covalently bound to other atoms, as deduced the separation between the atomic nuclei in molecules. In principle, the distance between two atoms that are bound to each other in a molecule (the length of that covalent bond) should equal the sum of their covalent radii. Ionic radius: the nominal radius of the ions of an element in a specific ionization state, deduced from the spacing of atomic nuclei in crystalline salts that include that ion. In principle, the spacing between two adjacent oppositely charged ions (the length of the ionic bond between them) should equal the sum of their ionic radii. István Szalai (Eötvös University) Lecture 1 15 / 40 Periodic Table Atomic radii Metallic radius: the nominal radius of atoms of an element when joined to other atoms by metallic bonds. van der Waals radius: in principle, half the minimum distance between the nuclei of two atoms of the element that are not bound to the same molecule. István Szalai (Eötvös University) Lecture 1 16 / 40 Periodic Table Atomic radii István Szalai (Eötvös University) Lecture 1 17 / 40 Periodic Table Atomic radii István Szalai (Eötvös University) Lecture 1 18 / 40 Periodic Table Atomic radii Effective nuclear charge: Zeff = Z − S where S is the shielding effect of the inner electrons. l 0 1 2 3 ... 1,0 1,0 1,0 1,0 ni−1 0,85 0,85 1,0 1,0 ni 0,3 0,35 0,35 0,35 Na 1s 2 2s 2 2p 6 3s 1 Zeff = 11 − (2 × 1 + 8 × 0, 85) = 2, 2 Al 1s 2 2s 2 2p 6 3s 2 3p 1 Zeff = 13 − (2 × 1 + 8 × 0, 85 + 2 × 0, 3) = 3, 6 Na 1.86 × 10−10 m (186 pm) Al 1.43 × 10−10 m (143 pm) István Szalai (Eötvös University) Lecture 1 19 / 40 Periodic Table Ionic radii Isolectronic species (10e − ) Na+ Mg2+ radius (nm) 0,097 0,066 István Szalai (Eötvös University) Al3+ 0,051 Lecture 1 20 / 40 Periodic Table Ionization Energy First ionization energy: the minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom. Be(g) → Be+ (g) + e − 899 kJ/mol + 2+ − Be (g) → Be (g) + e 1757 kJ/mol IE1 < IE2 2+ 3+ − Be (g) → Be (g) + e 14849 kJ/mol IE1 < IE2 IE3 István Szalai (Eötvös University) Lecture 1 21 / 40 Periodic Table Ionization Energy István Szalai (Eötvös University) Lecture 1 22 / 40 Periodic Table Ionization Energy István Szalai (Eötvös University) Lecture 1 23 / 40 Periodic Table Electron Affinity The amount of energy absorbed when an electron is added to an isolated gaseous atom. Cl(g) + e − → Cl− (g) −349 kJ/mol − − O(g) + e → O (g) −141 kJ/mol − − 2− O (g) + e → O (g) +744 kJ/mol Generally, nonmetals have more positive EA than metals. Atoms whose anions are more stable than neutral atoms have a greater EA. István Szalai (Eötvös University) Lecture 1 24 / 40 Periodic Table Electron Affinity István Szalai (Eötvös University) Lecture 1 25 / 40 Periodic Table Electron Affinity István Szalai (Eötvös University) Lecture 1 26 / 40 Periodic Table Electronegativity The electronegativity of an element is a measure of the relative tendency of an atom to attract electrons to itself when it is chemically combined with another atom. Pauling proposed the concept of electronegativity in 1932 as an explanation of the fact that the covalent bond between two different atoms (A–B) is stronger than would be expected by taking the average of the strengths of the A–A and B–B bonds. István Szalai (Eötvös University) Lecture 1 27 / 40 Periodic Table Electronegativity To calculate Pauling electronegativity for an element, it is necessary to have data on the dissociation energies of at least two types of covalent bond formed by that element. p ∆ = EAB − (EAA × √ EBB ) ENA − ENB = 0.102 ∆ ENF = 4.0 István Szalai (Eötvös University) Lecture 1 28 / 40 Periodic Table Electronegativity Mulliken proposed that the arithmetic mean of the first ionization energy and the electron affinity should be a measure of the tendency of an atom to attract electrons. EN = István Szalai (Eötvös University) IE + EA 130 Lecture 1 29 / 40 Periodic Table Electronegativity Allred and Rochow considered that electronegativity should be related to the charge experienced by an electron on the ”surface” of an atom: the higher the charge per unit area of atomic surface, the greater the tendency of that atom to attract electrons. EN = 0, 359 István Szalai (Eötvös University) Zeff + 0, 744 r2 Lecture 1 30 / 40 Periodic Table Electronegativity István Szalai (Eötvös University) Lecture 1 31 / 40 Periodic Table Chemical properties: hydrides Ionic hydride → covalent hydrides LiH NaH KH BeH2 MgH2 CaH2 B2 H6 (AlH3 )x Ga2 H6 CH4 SiH4 GeH4 NH3 PH3 AsH3 H2 O H2 S H2 Se HF HCl HBr LiH + H2 O → Li+ + OH− + H2 HCl + H2 O → H3 O+ + Cl− István Szalai (Eötvös University) Lecture 1 32 / 40 Periodic Table Chemical properties: oxides Metal oxides → Nonmetal oxides Li2 O Na2 O2 KO2 BeO MgO CaO B2 O2 Al2 O3 Ga2 O3 CO2 SiO2 GeO2 N2 O5 P4 O10 As2 O5 SO3 SeO3 OF2 Cl2 O7 Br2 O7 CaO + H2 O → Ca2+ + 2 OH− SO3 + 3 H2 O → 2 H3 O+ + SO2− 4 István Szalai (Eötvös University) Lecture 1 33 / 40 Periodic Table Ionic Bonding 2 Na(s) + Cl2 (g) → 2 NaCl(s) István Szalai (Eötvös University) Lecture 1 34 / 40 Periodic Table Ionic Bonding 2 Na(s) + Cl2 (g) → 2 NaCl(s) Na [Ne] Cl [Ne] ↑ 3s ↑↓ 3s ↑↓ ↑↓ ↑ 3p → Na+ [Ne] → Cl− [Ne] ↑↓ 3s ↑↓ ↑↓ ↑↓ 3p Q2 α EB = 4π0 r Electrostatic interaction (Coulomb force) α(NaCl) = 1.7475 István Szalai (Eötvös University) Lecture 1 35 / 40 Periodic Table Ionic Bonding István Szalai (Eötvös University) Lecture 1 36 / 40 Periodic Table Types of Ions Noble gas configuration s 2 : H− , Li+ , Be2+ s 2 p 6 : pl. Na+ , Ca2+ , Sc3+ , Cl− , O2− . . . d 10 s 2 configuration Sn ([Kr]5s 2 4d 10 5p 2 ) → Sn2+ ([Kr]5s 2 4d 10 ) + 2e− Tl+ ,Pb2+ , Bi3+ , . . . István Szalai (Eötvös University) Lecture 1 37 / 40 Periodic Table Types of Ions Ions of transition metals The first transition series is the result of the 3d orbitals being filled after the 4s orbital. However, once the electrons are established in their orbitals, the energy order changes - and in all the chemistry of the transition elements, the 4s orbital behaves as the outermost, highest energy orbital. The reversed order of the 3d and 4s orbitals only applies to building the atom up in the first place. In all other respects, the 4s electrons are always the electrons you need to think about first. István Szalai (Eötvös University) Lecture 1 38 / 40 Periodic Table Types of Ions Ions of transition metals d 10 configuration Zn ([Ar]4s 2 3d 10 ) → Zn2+ ([Ar]3d 10 ) + 2e− Cu+ , Ag+ , Cd2+ , Tl3+ , . . . [Ar]3d 1 : [Ar]3d 2 : [Ar]3d 3 : [Ar]3d 4 : [Ar]3d 5 : [Ar]3d 6 : [Ar]3d 7 : [Ar]3d 8 : [Ar]3d 9 : Ti3+ V3+ Cr3+ Mn3+ Mn2+ , Fe3+ Fe2+ , Co3+ Co2+ Ni2+ Cu2+ István Szalai (Eötvös University) Lecture 1 39 / 40 Periodic Table Types of Ions István Szalai (Eötvös University) Lecture 1 40 / 40