Chapter 7 Periodic Properties of the Elements
periodic table the most significant tool that chemist use for organizing and remembering chemical facts 7.1 Development of the periodic table discovery of elements 9 42 11 6 18 17 11 1860 D. Mendeleev & L. Meyer similar chemical and physical properties 1915 H. Moseley concept of atomic number atomic number = number of protons = number of electrons
Mendeleev s periodic table predict unknown elements Sc Ga Ge Tc
7.2 Effective nuclear charge the attraction force between electrons and nucleus the magnitude of the net nuclear charge acting on the electron the average distance between the nucleus and the electron effective nuclear charge Z eff = Z S S: average number of electrons between the nucleus and the electron in question effective nuclear charge experienced by the outer electrons is determined primarily by the difference between the charge on the nucleus and the charge of the core electrons Na roughly estimated Z eff = +1 detailed calculated Z eff = +2.5
effective nuclear charge increases as moving across the period (more significant) slightly increases as moving down the family 7.3 Sizes of atoms and ions nonbonding atomic radius Ar atoms collide with one another (van der Waals radius) bonding atomic radius Cl 2 molecule (covalent radius) determination of the bonding atomic radius distance between two I nuclei in I 2 : 2.66 Å bonding atomic radius of I = 1.33 Å distance between two adjacent C nuclei in diamond: 1.54 Å bonding atomic radius of C = 0.77 Å
bonding atomic radius ex. 7.1 estimate the length of C S, C H, S H bonds in methyl mercaptan CH 3 SH C S: 0.77 + 1.02 = 1.79 Å (1.82) C H: 0.77 + 0.37 = 1.14 Å (1.10) S H: 1.02 + 0.37= 1.39 Å (1.33) periodic trends in atomic radii 1. within each group the atomic radius tends to increase as proceeding from top to bottom 2. within each period the atomic radius tends to decreases as moving from left to right
trends in the sizes of ions cations are smaller than their parent atoms anions are larger than their parent atoms for ions of same charge, size increases as going down a group isoelectronic (same number of electrons) series O 2-, F -, Na +, Mg 2+, Al 3+ increasing nuclear charge 1.40 1.33 0.97 0.66 0.51 Å decreasing ionic radius
7.4 Ionization energy the minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion first ionization energy I 1 Na (g) Na + (g) + e- second ionization energy I 2 Na + (g) Na 2+ (g) + e- I 1 < I 2 < I 3 <.. sharp increase when an inner-shell electron is removed periodic trends in I 1 generally increases with increasing atomic number in each row O (slight irregularity) S generally decreases with B Se Al Ga increasing atomic number in each group
the representative elements show a larger range of values of I 1 than do the transition-metal elements electron configurations of ions Li (1s 2 2s 1 ) Li + (1s 2 ) Fe ([Ar]3d 6 4s 2 ) Fe 2+ ([Ar]3d 6 ) Fe 3+ ([Ar]3d 5 ) F (1s 2 2s 2 2p 5 ) F - (1s 2 2s 2 2p 6 )
7.5 Electron affinities energy change when an electron is added to a gaseous atom Cl (g) + e - Cl - (g) E = -349 kj/mol Ar (g) + e - Ar - (g) E > 0
7.6 Metals, nonmetals, and metalloids nonmetals metals metalloids metal character solids at room temperature Cr (mp =1900 o C) except Hg liquid (mp = -39 o C) Cs (mp = 28.4 o C) Ga (mp = 29.8 o C)
metals tend to have low ionization energy tend to form positive ions relatively easily compounds of metals with nonmetals tend to be ionic substance metal oxides and halides are ionic solids 2Ni (s) + O 2(g) 2NiO (s) most metal oxides are basic metal oxide + water metal hydroxide O 2- (aq) + H 2 O (l) 2OH (aq) metal oxide + acid salt + water NiO (s) + 2HCl (aq) NiCl 2(aq) + H 2 O (l)
nonmetals nonmetals vary greatly in appearance gas: H 2, N 2, O 2, F 2, Cl 2 liquid: Br 2 solid: I 2, C(diamond, graphite), S 8, P 4 nonmetals tend to gain electrons when they react with metals 2Al (s) + 3Br 2(l) 2AlBr 3(s) compounds composed entirely of nonmetals are molecular substances most nonmetal oxides are acidic nonmetal oxide + water acid P 4 O 10(s) + 6H 2 O (l) 4H 3 PO 4(aq) CO 2(g) + H 2 O (l) H 2 CO 3(aq) nonmetal oxide + base salt + water CO 2(g) + NaOH (aq) Na 2 CO 3(aq) + H 2 O (l)
7.7 Group trends for the active metals Group 1A alkali metals soft metallic solids low densities low melting points very reactive, readily losing an electron to form cation with a +1 charge combine directly with most nonmetals 2M (s) + H 2(g) 2MH (s) hydride 2M (s) + S (s) M 2 S (s) react vigorously with water 2M (s) + 2H 2 O (l) 2MOH (aq) + H 2(s) oxides: Li 2 O peroxide: Na 2 O 2 superoxide: KO 2 flame test Li red Na yellow K blue Group 2A alkaline earth metals less reactive than alkali metals Mg (s) + H 2 O (g) Ca (s) + 2H 2 O (l) MgO (s) + H 2(s) 2Ca(OH) 2(aq) + H 2(s) Mg Ca essential for living organism flame test Ca brick red Sr crimson red Ba green
7.8 Group trends for selected nonmetals hydrogen H 2 H + (aq) hydronium ion H - hydride ion Group 6A oxygen group O 2 O 3 ozone allotropes O 2- oxide O 2-2 peroxide O 2- superoxide S 8 other allotropes S 2- sulfide S (s) + O 2(g) SO 2(g) Group 7A halogens F 2 pale yellow gas Cl 2 yellow-green gas Br 2 reddish brown liquid I 2 grayish black solid form halide ion X - 2F 2(g) + 2H 2 O (l) 4HF (aq) + O 2(g) 2F 2(g) + SiO 2(s) 4SiF 4(g) + O 2(g) Cl 2(g) + H 2 O (l) HCl (aq) + HOCl (aq) X 2(g) + H 2(g) HX (g) Group 8A noble gases monoatomic inert gas 1962 Neil Bartlett XeF 2 XeF 4 XeF 6 KrF 2 HArF only stable at low temperature