Presentation on theme: "Periodic Properties of the Elements"— Presentation transcript:
1Periodic 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
2Electron 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.
3Atomic SizeThe radius of an atom is found from the distance between nuclei in a molecule.
4Trends in Atomic SizeSize increases going down a column of the periodic table.Size decreases from left to right in a row.
5Trends in Atomic SizeIn a column, size increases with the addition of successive electron shells.Example: The alkali metalsCs6sRb5sK4sNa3sLi2s
6Effective Nuclear Charge The size of an orbital decreases with increasing Zeff for the valence electrons.2pLi+32s1s2p
7Effective Nuclear Charge The size of an orbital decreases with increasing Zeff for the valence electrons.Be+4
8Effective Nuclear Charge The size of an orbital decreases with increasing Zeff for the valence electrons.B+5
9Effective Nuclear Charge The size of an orbital decreases with increasing Zeff for the valence electrons.C+6
10Effective Nuclear Charge The size of an orbital decreases with increasing Zeff for the valence electrons.O+8
11Effective Nuclear Charge The size of an orbital decreases with increasing Zeff for the valence electrons.F+9
12Ionization 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
13Ionization 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:
14Trends 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.
15Trends 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.
16Electron 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
17Metals and NonmetalsMetals are characterized by low ionization energy; Nometals by high electron affinity.
18Group 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)
19Groups 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