Chapter 7. Periodic Properties of the Elements.

7.1 Development of The Periodic Table.

The number of known elements in 1800 was 31, but had increased to 63 by 1865. In 1869 Mendeleev and Lothar Meyer published almost identical classification schemes. Both noticed that similar chemical properties recurred when elements were placed in order of increasing atomic weight. Mendeleev advanced his ideas more vigorously, and had the insight to leave blank spaces and predict the existence of yet undiscovered elements. Dimitri Mendeleev

(1834-1907)

MissingElement=eka-silicon

Group 5

Group 1Group 7

Group 4

Group 6

Group 3

Group 2

Group 2bGroup 1b

Group 4bGroup 5bGroup 6b

Thus, Mendeleev predicted the existence of what he called eka-silicon, which we today refer to as Germanium.

Mendeleev predictions. Observed props. of Ge

At. Wt. 72 72.59 g/molDensity: 5.5 5.35 g/ccSp. Heat: 0.305 0.309 J/g-KM.Pt. high 947 oCColor dark gray grayish white

Oxide: MO2 GeO2

Density oxide: 4.7 4.70

Chloride: XCl4 GeCl4

B.Pt. chloride: <100 84 oC

The shells and sub-shells being filled across the periodic table:

Appearance:

Li

Na

Be

Mg

Group I A Group II A

Because they have the same valence electron configuration, elements in the same group havesimilar chemistry:

Reactivity:

Sodium in water Potassium in water

All the elements in group 1 react violently with water:

2 M (s) + 2 H2O (l) 2 MOH (aq) + H2(g)

Atomic Numbers (Z):

Henry Moseley (1887-1915) developed the concept of atomic numbers. Each element produces X-rays of a distinct frequency, which increased with atomic number. Atomic number = number of protons in the nucleus. The concept of At. No. allowed correct placement of e.g. Ar (Z = 18) before K (Z = 19) even though At. Wt. Ar > K. A similar problem occurs with Co (At. Wt = 58.9) and Ni (58.7), even though Z = 27 for Co and 28 for Ni.

Henry Moseley(1887-1915)

The K x-ray spectrum of germanium:

Energy (keV) (kilo-electron-volts)

Relationship between the frequency of the K x-ray peak and atomic number of the element (Z)

Z

Al

Kr

Ar

K

determining Z from energyof x-ray peak now placesAr and K in the right order

The core electrons and the valence electrons:

Nucleus = 12+

coreElectrons= Ne 10 e-

valenceelectronsin 3s shellexperiencesmall netpositivecharge

the core electrons shieldthe valenceelectrons fromthe nucleus

e

e

The Mgatom

7.2 Effective Nuclear Charge.

In a many-electron atom, each electron is simultaneously attracted to the nucleus, and repelled by the other electrons. Generally, the valence electrons are screened from the nucleus by the core electrons (~S, where S is the screening constant = number of core electrons). The effective nuclear charge is the approximate charge experienced by the valence electrons. Thus, the effective nuclear charge Zeff, experienced by the valence electrons is given by:

Zeff = Z - S

Thus, for Na, Zeff = 11 – 10 ~ 1+. However, because electrons in higher shells are less good at shielding the nucleus, the charge experience by Na+ is actually about +2.5.

number ofcore electrons

Which would you expect to experience a greater Zeff, a 2p electron of an F atom or a 3s electron of an Na atom?

Zeff (F, 2p) = +9 – 2 = 7+

Zeff (Na, 3s) = +11 – 10 = 1+

much more tightly bound !

Which electron do you think would be easier to remove?

The Na 3s electron is easily removed => metals easily lose electrons

The F 2p electron is NOT easily removed => it is much harder to remove an electron from a halogen

7.3 Sizes of Atoms and Ions.

We can picture atoms as having size in various ways. The bonding atomic radius, or covalent radius is half the distance in homonuclear diatomic molecules, e.g. H2, Cl2. The C-C distance in diamond gives us the covalent radius of C as 1.54/2 = 0.77 Å. Bond lengths are given quite closely as the sum of covalent radii, e.g. C-Cl = 0.77 + 0.99 = 1.76 Å. The non-bonding atomic radius = van der Waals radius.

Atomic non-bonding (van der Waals) and bonding (covalent) radii of atoms:

iodine atom hydrogen atom

Non-bonding atomic radii of I and H atoms

2 x I-I bondedAtomic radius

I2 molecule H2 molecule HI molecule

2 x H-H bondedAtomic radius

The bonding atomic radius(‘covalent radius’) of I is ½ the I-I distance in I2. The H-Iinternuclear distance in HIis the sum of the bondingatomic radii of H and I.

Ionic radii:

The ionic radii are the radii of the ions. For cations the ionic radius is smaller than the atomic radius, because there are fewer electrons around the nucleus, while for anions the ionic radii are larger than the atomic radii because there are now more electrons to repel each other:

Lithium atom lithium cationradius = 1.34 Å ionic radius =

0.68 Å

fluorine atom fluorine anionradius = 0.71 Å ionic radius =

1.33 Å

Li+

F-

Li

F

The sizes of atoms (gray) cations (pink) and anions (blue)(textbook p. 263):

Periodic trends in Atomic and ionic radii.

1) Atomic and ionic radius increases as we move down a group.

2) Atomic and ionic radius tends to decrease along a row from left to right

Variation in atomic radius with Z (atomic number)

Atomic number (Z)

He

NeAr

KrXe Rn

CsRb

K

Na

Li

Periodic trends in ionic radii.

Ionic radii are based on the distance between ions in ionic compounds. In general, cations are smaller than the atoms from which they are derived, but anions are larger.

O2- F- Na+ Mg2+

Al3+

1.40 1.33 0.97 0.66 0.51Å

The sizes of atoms (gray) cations (pink) and anions (blue):

Ionization energy.

The ionization energy is the minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion. E.g. First (I1)and second (I2) ionization energies of Na+ are given by ΔH for the processes:

Na(g) → Na+(g) + e-

Na+(g) → Na2+(g) + e-

Trends in 1st Ionization Energies

1st ionization energy

Variations in successive ionization energies:

The value of In increases as n increases for any one element. It increases sharply as we start to remove inner-shell electrons. Core (inner-shell) electrons are much harder to remove that valence electrons:

Why does the ionization energy increase sharply after removalof the 5th electron for Nitrogen?

What would a similar graph look like for beryllium (Be)?

Nitrogen Beryllium

core electrons are much harder to remove than valence electrons !

Periodic Trends in First Ionization Energies:

If we look at how I1 varies across the periodic table, we find that:

1. I1 increases across a row with increasing Z.

2. I1 decreases down a group, e.g. Li>Na>K>Rb>Cs, or He>Ne>Ar>Kr>Xe.

3. The representative elements (‘main group elements’) show a larger variation in I1 than do transition elements. The f-block elements show only a small variation in I1.

Variation in first ionization energy with Z along the periodic table:

Atomic number (Z)

He

ArKr

XeRn

Ne

CsRbKNaLi

Noble gases occur at the peaks

Alkali metals occurat the troughs

Electron Configuration of Ions:

When electrons are lost from an atom to form a cation, they are always lost from the shell having the highest principal quantum number, i.e. the highest energy.

e.g. Li(1s22s1) → Li+ (1s2)

loseselectron

Electronic configurations of transition metals (the d-block)

For transition metal ions one finds that in e.g. Fe2+, the two electrons lost are lost from the 4s2 level, even though in the neutral atom this level is lower in energy than the 3d level.

Fe ([Ar]3d64s2) → Fe2+([Ar]3d6).

Addition of electrons:

When an electron is added, it goes into the empty or partially filled orbital having the lowest value of n.

F (1s22s22p5) → F-(1s22s22p6)

7.5 Electron affinities:

This is the energy for adding electrons to the atom.

Cl(g) + e- → Cl-(g) ΔE = -349 kJ/mol

The halogens have the most negative electron affinities, followed by the chalcogens. Group 2 has very small EA’s, because they have a filled s orbital. Similarly, group 5 has small EA’s because they have a half-filled p subshell, which is particularly stable. For the noble gases, EA’s are all positive, corresponding to the lack of chemistry, except for Xe.

Electron affinities (kJ/mol) for some elements:

7.6 Metals, non-metals, and metalloids.

Metals are lustrous (shiny) and malleable (can be pounded into thin sheets) and ductile (can be drawn into wires), and conduct heat and electricity. All are solids at room temperature except Hg, and Cs and Ga melt at slightly above room temperature. Compounds of metals with non-metals tend to be ionic substances. Most metal oxides are basic, i.e. dissolve in water to give basic solutions.

Na2O + H2O = 2 NaOH

Non-metals.

These tend not to be lustrous, and usually are poor conductors of heat and electricity. Generally low-melting, although diamond melts at 3570 oC. Because of their electron affinities, non-metals tend to gain electrons when they react with metals, e.g.

2 Mg(s) + O2(g) → 2 MgO(s)

Compounds composed entirely of non-metals tend to be molecular substances.

Most non-metal oxides dissolve in water to form acids:

SO3(s) + H2O(l) → H2SO4(aq)

CO2(g) + H2O(l) → H2CO3(aq)

acid + base = salt + water

H2SO4 + CaO = CaSO4 + H2O

7.7. Trends for the active metals.

Students read 7.7 through 7.9 on own time.