1. Chapter 7 Periodic Properties of the Elements
CHEMISTRY The Central Science 9th Edition
Chapter 7 Periodic Properties of the Elements
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Chapter 7

2. 7.1: Development of the Periodic Table
Dimitri Mendeleev and Lothar Meyer arranged the elements in order of increasing atomic weight
Certain elements were missing from this scheme:
Mendeleev noted:
As properly belonged underneath P and not Si
A missing element underneath Si
He predicted a number of properties for this element
In 1886 Ge was discovered
Properties of Ge match Mendeleev’s predictions
Modern periodic table: arrange elements in order of increasing atomic number
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3. Effective Nuclear Charge
Effective nuclear charge is the charge experienced by an electron on a many-electron atom
The effective nuclear charge is not the same as the charge on the nucleus
Due to the effect of the inner electrons
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4. 7.2: Effective Nuclear Charge
Electrons are attracted to the nucleus, but repelled by the electrons that screen them from the nuclear charge
The nuclear charge experienced by an electron depends on
its distance from the nucleus
the number of core electrons
As the average number of screening electrons (S) increases, the effective nuclear charge (Zeff) decreases
OR…as the distance from the nucleus increases, S increases and Zeff decreases
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5. The ns orbitals all have the same shape, but have different sizes and different numbers of nodes (places where the probability function = 0)
Zeff = Z – S or
Protons – Electrons
The radial electron density is the probability of finding an electron at a given distance
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6. The period trend is of greater significance than the group trend
Zeff increases across a period
Added electrons don’t increase shielding
Zeff increases down a group
Added electrons are in increasingly higher energy levels but they are less able to shield the outer electrons from the nucleus
The period trend is of greater significance than the group trend
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7. 7.3: Sizes of Atoms and Ions
Consider a simple diatomic molecule:
The distance between the two nuclei is called the bond distance
Half of the bond distance is called the covalent radius of the atom
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8. Periodic Trends in Atomic Radii
The properties of elements vary periodically
Atomic size varies consistently through the periodic table
Atomic radius increases down a group (more noticeable)
Atomic radius decreases across a period
There are two factors at work:
Principal quantum number
Electrons fill higher energy levels
Effective nuclear charge
Electrons fill the same energy level, no increase in shielding
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9. Text, P. 263

10. Trends in the Sizes of Ions
Ion size is the distance between ions in an ionic compound
Cations vacate the most spatially extended orbital and are smaller than the parent atom
Anions add electrons to the most spatially extended orbital and are larger than the parent atom
Remember: electrons are gained or lost to achieve a noble gas configuration
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11. Text
P. 266

12. O2- > F- > Na+ > Mg2+ > Al3+
For ions of the same charge, ion size increases down a group
Electrons in higher energy levels
All the members of an isoelectronic series have the same number of electrons (O2-, F-, Na+, Mg2+, Al3+)
As nuclear charge increases in an isoelectronic series, the ions become smaller:
O2- > F- > Na+ > Mg2+ > Al3+
Added electrons are in the same energy level, so no increase in shielding
Removed electrons result in an increasingly positive nucleus which pulls electrons toward it
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13. 7.4: Ionization Energy
The first ionization energy, I1, is the amount of energy required to remove an electron from a gaseous atom:
Na(g)  Na+(g) + e-
The second ionization energy, I2, is the energy required to remove an electron from a gaseous ion:
Na+(g)  Na2+(g) + e-
The larger ionization energy, the more difficult it is to remove the electron
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Chapter 7

15. Variations in Successive Ionization Energies
There is a sharp increase in ionization energy when a core electron is removed:
Text, 268
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16. Think: does it want to lose an electron???
Periodic Trends in Ionization Energies
Think: does it want to lose an electron???
Ionization energy decreases down a group
The outermost electron is more readily removed as electrons fill higher energy levels
The s electrons are more effective at shielding than p electrons
Irregularity in group 2: sharp increase in I1
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17. Prentice Hall © 2003
Chapter 7
Text, P. 270

18. Ionization energy generally increases across a period
Across a period, Zeff increases
It becomes more difficult to remove an electron
Irregularity in group 15: sharp increase in I1 (half-filled sublevel)
Transition metals and f-block elements show slow variations in I1 due to shielding
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19. Text, P. 270

21. Fe ([Ar]3d6 4s2)  Fe3+ ([Ar]3d5)
Electron Configuration of Ions
Cations: electrons removed from orbital with highest principle quantum number, n, first:
Li (1s2 2s1)  Li+ (1s2)
Fe ([Ar]3d6 4s2)  Fe3+ ([Ar]3d5)
Anions: electrons added to the orbital with highest n:
F (1s2 2s2 2p5)  F- (1s2 2s2 2p6)
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22. Think: does it want to gain an electron???
7.5: Electron Affinities
Think: does it want to gain an electron???
Electron affinity is the energy change when a gaseous atom gains an electron to form a gaseous ion:
Cl(g) + e-  Cl-(g)
Electron affinity can either be exothermic (above) or endothermic:
Ar(g) + e-  Ar-(g)
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23. EA becomes more negative across a period
A greater attraction between the atom and the added electron
Groups 2, 15 and 18
EA has no real group trend
Text, P. 272

24. 7.6: Metals, Nonmetals, and Metalloids
Text, P. 274

25. Metallic character refers to the properties of metals
shiny or lustrous
malleable and ductile
oxides form basic ionic solids
form cations in aqueous solution
Metallic character increases down a group
Metallic character decreases across a period
Metals have low ionization energies
Most neutral metals are oxidized rather than reduced
form cations
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26. Text, P. 275
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27. Metal oxide + water  metal hydroxide
Most metal oxides are basic:
Metal oxide + water  metal hydroxide
Nonmetals
Nonmetals are more diverse in their behavior than metals
Properties are opposite those of metals
Compounds composed only of nonmetals are molecular
When nonmetals react with metals, nonmetals tend to gain electrons:
metal + nonmetal  salt
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28. Most nonmetal oxides are acidic: nonmetal oxide + water  acid
Nonmetal oxide + base  salt + water
The greater the nonmetallic character of the central atom, the stronger the acidic character of the oxide
Metalloids
Metalloids have properties that are intermediate between metals and nonmetals
Si has a metallic luster but it is brittle
Semiconductors
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29. 7.7: Group Trends for the Active Metals
Group 1A: The Alkali Metals
Alkali metals are all soft
Chemistry dominated by the loss of their single s electron:
M  M+ + e-
Reactivity increases down the group
Alkali metals react with water to form MOH and hydrogen gas:
2M(s) + 2H2O(l)  2MOH(aq) + H2(g)
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30. Alkali metals produce different oxides when reacting with O2:
4Li(s) + O2(g)  2Li2O(s) (oxide)
2Na(s) + O2(g)  Na2O2(s) (peroxide)
K(s) + O2(g)  KO2(s) (superoxide)
(Rb and Cs, too)
Alkali metals: Flame tests
The s electron is excited by the flame and emits energy when it returns to the ground state
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31. Na line (589 nm): 3p  3s transition Li line: 2p  2s transition
K line: 4p  4s transition
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32. Text, P. 278
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33. M(s) + 2H2O(l)  M(OH)2(aq) + H2(g)
Group 2A: The Alkaline Earth Metals
Alkaline earth metals are harder and more dense than the alkali metals
The chemistry is dominated by the loss of two s electrons:
M  M2+ + 2e-
Mg(s) + Cl2(g)  MgCl2(s)
Increasing activity down the group:
Be does not react with water
Mg will only react with steam
Ca, Sr and Ba:
M(s) + 2H2O(l)  M(OH)2(aq) + H2(g)
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34. Group 2A: The Alkaline Earth Metals
Text, P. 281
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35. 7.8: Group Trends for Selected Nonmetals
Hydrogen
Colorless diatomic gas, H2
High IE because there are no inner electrons
It can either gain another electron to form the hydride ion, H-, or lose its electron to become H+:
2Na(s) + H2(g)  2NaH(s)
2H2(g) + O2(g)  2H2O(g)
H+ is a proton
The aqueous chemistry of hydrogen is dominated by H+(aq)
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36. Group 16: The Oxygen Group
Metallic character increases down a group (O2 is a gas, Te is a metalloid, Po is a metal)
Two important forms of oxygen: O2 and ozone, O3
Ozone can be prepared from oxygen:
3O2(g)  2O3(g) H = kJ.
Ozone is pungent and toxic
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37. Group 16: The Oxygen Group
Text, P. 283
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38. Oxygen (O2) is a potent oxidizing agent
O2- ion has a noble gas configuration
There are two oxidation states for oxygen: 2- and 1-
Sulfur is another important member of this group
Most common form of sulfur is yellow S8
Sulfur tends to form S2- in compounds (sulfides)
Allotropes: different forms of the same element in the same state
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39. Fluorine is one of the most reactive substances known
Group 17: The Halogens
The chemistry of the halogens is dominated by gaining an electron to form an anion:
X2 + 2e-  2X-
Fluorine is one of the most reactive substances known
All halogens consists of diatomic molecules, X2
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40. Group 17: The Halogens
Text, P. 285
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41. 2NaCl(aq) + 2H2O(l)  2NaOH(aq) + H2(g) + Cl2(g)
Chlorine is the most industrially useful halogen
Produced by the electrolysis of brine (NaCl):
2NaCl(aq) + 2H2O(l)  2NaOH(aq) + H2(g) + Cl2(g)
Hydrogen compounds of the halogens are all strong acids with the exception of HF
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42. These are all nonmetals and monatomic
Group 18: The Noble Gases
These are all nonmetals and monatomic
Unreactive because they have completely filled s and p sub-shells
In 1962 the first compound of the noble gases was prepared: XeF2, XeF4, and XeF6
Xe has a low enough IE to react with substances that readily remove electrons
To date the only other noble gas compounds known are KrF2 and HArF
“Argon hydrofluoride”
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43. Group 18: The Noble Gases
Text, P. 286
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