Presentation on theme: "Periodic Properties of the Elements"— Presentation transcript:
1 Periodic Properties of the Elements Early versions of the Periodic table were constructed by Mendeleev and Meyer.We now know that the periodic properties are due to the electronic structure of atoms.Electronic structure explains the observed trends inAtomic sizeIonization EnergyElectron AffinityAlZnGaCdInSn
2 Electron ShellsElectrons in successive shells (n = 1,2,3...) produce maxima in the electron density at increasing distances from the nucleus.The inner, completed electron shells are called core shells.The outermost electron shell is called the valence shell. It may include subshells of inner shells.Example: The valence shell of 33As includes 4s, 3d, and 4p subshells.
3 Atomic SizeThe radius of an atom is found from the distance between nuclei in a molecule.
4 Trends in Atomic SizeSize increases going down a column of the periodic table.Size decreases from left to right in a row.
5 Trends in Atomic SizeIn a column, size increases with the addition of successive electron shells.Example: The alkali metalsCs6sRb5sK4sNa3sLi2s
6 Effective Nuclear Charge The size of an orbital decreases with increasing Zeff for the valence electrons.2pLi+32s1s2p
7 Effective Nuclear Charge The size of an orbital decreases with increasing Zeff for the valence electrons.Be+4
8 Effective Nuclear Charge The size of an orbital decreases with increasing Zeff for the valence electrons.B+5
9 Effective Nuclear Charge The size of an orbital decreases with increasing Zeff for the valence electrons.C+6
10 Effective Nuclear Charge The size of an orbital decreases with increasing Zeff for the valence electrons.O+8
11 Effective Nuclear Charge The size of an orbital decreases with increasing Zeff for the valence electrons.F+9
12 Ionization EnergyThe energy required to remove an electron from an atom in its ground level.Example: HydrogenH(g) H+(g) + eI = EE-RH1s= 1312 kJ/mol
13 Ionization EnergyEnergies can be defined for successive ionizations. For Mg:Mg(g) Mg+(g) + e I1 = 738 kJMg+(g) Mg+2(g) + e I2 = 1450 kJMg+2(g) Mg+3(g) + e I3 = 7730 kJFor Al:
14 Trends in Ionization Energy Ionization energy increases with Zeff for the valence electrons.I1 decreases going down a column of the periodic table.I1 increases from left to right in a row.Alkali metals have the lowest ionization energies in a period. Rare gases have the highest.
15 Trends in Ionization Energy Arrange the following with increasing ionization energy:C, K, Mg, Na, Ne, SiK < Na < Mg < Si < C < NeWhy is the ionization energy of N greater than O?The 2p subshell of N has a slightly higher energy due to electron repulsion.
16 Electron AffinityThe change in energy when an electron is added to a gaseous atom.Cl(g) + e Cl(g) E = -349 kJ/molA large, negative E indicates strong attraction between the atom and the added electron.A positive E indicates the addition of an electron is unfavorable.Ne(g) + e Ne(g) E = 40 kJ/mol
17 Metals and NonmetalsMetals are characterized by low ionization energy; Nometals by high electron affinity.
18 Group 1A: Alkali Metals Alkali metals react to lose electrons. These metals are very reactive with reactivity increasing down the group.Reaction with H2O:Na(s) + H2O(l) NaOH(aq) + H2(g)
19 Groups 2A and 7AThe alkaline earth metals have higher ion-ization energies than alkali metals and are less reactive.Ca + H2O Ca(OH)2 + H2The halogens react to gain electrons. Reactivity decreases going down the group.Cl2 + 2 Br 2 Cl + Br2