PowerPoint to accompany Chapter 6 Periodic Properties of the Elements

PowerPoint to accompany

Chapter 6

Periodic Properties

of the Elements

Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia

Development of the Periodic Table

Elements in the same group generally have similar chemical properties.

Properties are not identical, however.


S8 rings stackto make

OrthorhombicCrystals

O2 gas

In a gas burette



Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped.

Figure 6.2



(which he called eka-silicon) as an element with an atomic weight between that of zinc and arsenic, but with chemical properties similar to those of silicon.

Mendeleev, for instance, predicted the discovery of germanium,

Table 6.1


Effective Nuclear Charge

In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons.

The nuclear charge that an electron experiences depends on both factors.


Effective Nuclear Charge

The effective nuclear charge, Zeff, is found this way:

Zeff = Z − S

where Z is the atomic number and S is a screening constant, usually close to the number of inner electrons.

Figure 6.3


Sizes of Atoms and Ions

The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.

Figure 6.4


Sizes of Atoms

Bonding atomic radius

tends to: decrease from left to

right across a row

due to increasing Zeff

increase from the top to the bottom of a column due to increasing value of n

Figure 6.5


Sizes of Ions

Ionic size depends upon: Nuclear charge Number of

electrons Orbitals in which

electrons reside

Figure 6.6


Sizes of Ions

Cations are smaller than their parent atoms: The outermost

electron is removed and repulsions are reduced.

Figure 6.6


Sizes of Ions

Anions are larger than their parent atoms: Electrons are

added and repulsions are increased.

Figure 6.6


Sizes of Ions

Ions increase in size as you go down a column Due to increasing value of n.

In an isoelectronic series, ions have the same number of electrons.

Ionic size decreases with an increasing nuclear charge.


Ionisation Energy

Amount of energy required to remove an electron from the ground state of a gaseous atom or ion:

First ionisation energy is that energy required to remove the first electron.

Second ionisation energy is that energy required to remove the second electron, etc.


Ionisation Energy It requires more energy to remove each successive

electron.

When all valence electrons have been removed, the ionisation energy takes a quantum leap.

Table 6.2



Trends in First Ionisation Energies

As one goes down a column, less energy is required to remove the first electron.

For atoms in the same group, Zeff is essentially the same, but the valence electrons are farther from the nucleus.

Figure 6.8


Trends in First Ionisation Energies

Generally, as one goes across a row, it gets harder to remove an electron:

As you go from left to right, Zeff increases. Figure 6.7


Electron Affinities

Energy change accompanying the addition of an electron to a gaseous atom:

Cl + e− Cl−

E = -349 kJ/mol


Trends in Electron Affinity

In general, electron affinity becomes more exothermic as you go from left to right across a row.

Figure 6.9


Metal, Nonmetals and Metalloids

Figure 6.10


Metals versus Nonmetals

Differences between metals and nonmetals tend to revolve around these properties.

Table 6.3


Metals versus Nonmetals

Metals tend to form cations. Nonmetals tend to form anions.

Figure 6.12


Metals

Tend to be lustrous, malleable, ductile, and good conductors of heat and electricity.

Figure 6.11


Metals Compounds formed between metals and nonmetals tend to be

ionic.

Metal oxides tend to be basic.

Figure 6.15


Reactivity of Metals Metal reactivity varies with the type of metal. Alkali metals are very reactive because it is

very easy to remove the 1s1 electron. All alkali metals react vigorously with water to produce metal hydroxide and hydrogen gas:

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

Figure 6.13


Reactivity of Metals Alkaline earth metals are less reactive than alkali metals. Mg reacts only with steam, Be does not react with water, but others react readily with water. Reactivity tends to increase as you move down groups.

Figure 6.14


Nonmetals Dull, brittle substances that are poor conductors

of heat and electricity. Tend to gain electrons in reactions with metals to

acquire noble gas configuration. Substances containing only nonmetals are

molecular compounds. Most nonmetal oxides are acidic.

Figure 6.16


Metalloids

Have some characteristics of metals and some of nonmetals.

For instance, silicon looks shiny, but is brittle and a fairly poor conductor.


Mn Oxides


Hg and Br


Documents

PowerPoint to accompany Chapter 6 Periodic Properties of the Elements