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Anomalies Cr: 4s13d5 Mo: 5s14d2 Cu: 4s13d10 Ag: 5s13d10 Au: 6s13d10 Categories of electrons Inner (core) electrons are those seen in the previous noble gas and any completed transition series. They fill all the lower energy levels of an atom. Outer electrons are those in the highest energy level (highest n value). They spend more of their time farthest from the nucleus. Valence electrons are those involved in forming compounds. Among the main-group elements, the valence electrons are the outer electrons. For the transition elements, all the (n -1)d electrons are counted among the valence electrons also, even though the elements Fe (Z = 26 through Zn (Z = 30) use only a few of them in bonding as we will see later towards the last weeks of the course. A schematic of the SCF method for obtaining the orbitals of a polyelectronic atom. The radial distribution of electron probability density for the sodium atom. The shaded area represents the 10 core electrons. The radial distributions of the 3s, 3p, and 3d orbitals are also shown. Note the difference in the penetration effects of an electron in these three orbitals. Radial probability distributions for the 3d and 4s orbitals. Note that the most probable distance of the electron from the nucleus for the 3d orbital is less than that for the 4s orbital. However, the 4s orbital allows more electron penetration close to the nucleus and thus is preferred over the 3d orbital. Periodic Trends • Ionization Energy • Electron Affinity • Atomic Radius Ionization Energy • Energy in kJ required for the complete removal of 1 mol of electrons from 1 mol of gaseous atoms or ion • Many-electron atoms can lose more than one electron. The first ionization energy (IE1) removes an outermost electron (highest sublevel) from the gaseous atom: Atom (g) → ion+ (g) + e ∆E = IE1 • The second ionization energy (IE2) removes a second electron: Ion+ (g) → ion2+ (g) + e ∆E = IE2 (IE2 > IE1) X(g) → X+(g) + e– Mg → Mg+ + e– I1 = 735 kJ/mol( 1st IE) Mg+ → Mg2+ + e– I2 = 1445 kJ/mol (2nd IE) Mg2+ → Mg3+ + e– I3 = 7730 kJ/mol *(3rd IE) *Core electrons are bound much more tightly than valence electrons. Ionization Energy • In general, as we go across a period from left to right, the first ionization energy increases. • Why? – Electrons added in the same principal quantum level do not completely shield the increasing nuclear charge caused by the added protons. – Electrons in the same principal quantum level are generally more strongly bound from left to right on the periodic table. Ionization Energy • In general, as we go down a group from top to bottom, the first ionization energy decreases. • Why? – The electrons being removed are, on average, farther from the nucleus. The Values of First Ionization Energy for the Elements in the First Six Periods Explain why the graph of ionization energy versus atomic number (across a row) is not linear. electron repulsions Where are the exceptions? some include from Be to B and N to O Electron Affinity The electron affinity (EA) is the energy change in kJ accompanying the addition of 1 mol of electrons to 1 mol of gaseous atoms or ions. The first EA refers to the formation of 1 mol of mono-valent (1 –) gaseous anions: Atom (g) + e → ion- (g) ∆E = EA1 In most cases energy is released when an e is added because it is attracted to the atom’s nuclear charge. The second EA2 must always be positive because energy must be absorbed to overcome electrostatic repulsions and add an e to a negative ion. a) elements in group 6A and especially in 7A have high IE and high EA. They lose e with difficulty but attract e strongly. In their ionic compounds they form negative ions. b) Elements in groups 1A and 2A have low IE and slightly negative EA. They lose e readily but attract e very weakly. In their ionic compounds they form positive ions. c) Noble gases (Group 8A) have high IE and slightly positive EA. They tend not to lose or gain e . Electron Affinity X(g) + e– → X–(g) • In general as we go across a period from left to right, the electron affinities become more negative. • In general electron affinity becomes more positive in going down a group. Defining covalent and metallic radii The metallic radius is ½ the distance between nuclei of adjacent atoms in a crystal of the element. The covalent radius is ½ the distance between bonded nuclei in a molecule of the element. C; in a covalent compound, the bond length and known covalent radii are used to determine other radii. Bond length of C – Cl : 177 pm Covalent radius of Cl: 100 pm ∴ covalent radius of C: 177 – 100 = 77 pm − − Trends among main-group elements 1. Changes in n: as n increases the probability that the outer electrons spend more time farther from the nucleus increases and the atoms are larger. 2. Changes in Zeff : as the Zeff - the positive charge “felt” by an e a) Down a group, n dominates. Atomic radius generally increases in a group from top to bottom. b) Across a period Zeff dominates. Atomic radius generally decreases in a period from left to right. Trends among transition metals As we move from left to right, size shrinks through the first two or three transition elements because of the increasing Zeff. But from then on, the size remains relatively constant because shielding by the inner d electrons counteracts the increasing Zeff. The Periodic Table – Final Thoughts 1. It is the number and type of valence electrons that primarily determine an atom’s chemistry. 2. Electron configurations can be determined from the organization of the periodic table. 3. Certain groups in the periodic table have special names. The trend in acid-base behavior of element oxides • As the elements become more metallic down a group, the oxides become more basic (blue) • As the elements become less metallic across a period, their oxides become more acidic (red) Sb2O5: weakly basic, SiO2, As2O3: weakly acidic Al2O3: amphoteric The Period 4 Crossover in Sublevel Energies For main-group, s-block metals remove all electrons with the highest n value , p-block metals remove np electrons before ns For transition metals (d-block) remove ns electrons before (n-1)d electrons For nonmetals add electrons to the p orbitals of the highest value − Magnetic Properties A species with unpaired electron exhibits paramagnetism, attracted by an external magnetic field. A species with all electrons paired exhibits diamagnetism, it is not attracted and, in fact, is slightly repelled by a magnetic field. Ionic radius Ionic vs. atomic radius The Alkali Metals • Li, Na, K, Rb, Cs, and Fr – Most chemically reactive of the metals React with nonmetals to form ionic solids – Going down group: Ionization energy decreases Atomic radius increases Density increases Melting and boiling points smoothly decrease