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Development of the Periodic Table
Elements in the same group generally have similar chemical properties.
Properties are not identical, however.
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
S8 rings stackto make
OrthorhombicCrystals
O2 gas
In a gas burette
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Development of the Periodic Table
Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped.
Figure 6.2
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Development of the Periodic Table
(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
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
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.
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
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
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Sizes of Atoms and Ions
The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.
Figure 6.4
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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
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Sizes of Ions
Ionic size depends upon: Nuclear charge Number of
electrons Orbitals in which
electrons reside
Figure 6.6
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Sizes of Ions
Cations are smaller than their parent atoms: The outermost
electron is removed and repulsions are reduced.
Figure 6.6
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Sizes of Ions
Anions are larger than their parent atoms: Electrons are
added and repulsions are increased.
Figure 6.6
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
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.
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
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.
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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
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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
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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
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Electron Affinities
Energy change accompanying the addition of an electron to a gaseous atom:
Cl + e− Cl−
E = -349 kJ/mol
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Trends in Electron Affinity
In general, electron affinity becomes more exothermic as you go from left to right across a row.
Figure 6.9
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Metal, Nonmetals and Metalloids
Figure 6.10
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Metals versus Nonmetals
Differences between metals and nonmetals tend to revolve around these properties.
Table 6.3
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Metals versus Nonmetals
Metals tend to form cations. Nonmetals tend to form anions.
Figure 6.12
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Metals
Tend to be lustrous, malleable, ductile, and good conductors of heat and electricity.
Figure 6.11
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Metals Compounds formed between metals and nonmetals tend to be
ionic.
Metal oxides tend to be basic.
Figure 6.15
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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
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
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
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
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
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Metalloids
Have some characteristics of metals and some of nonmetals.
For instance, silicon looks shiny, but is brittle and a fairly poor conductor.