© 2012 Pearson Education, Inc. Chapter 7 Periodic Properties of the Elements Subhash C Goel South GA State College Douglas, GA 31533

© 2012 Pearson Education, Inc.

Chapter 7

Periodic Propertiesof the Elements

Subhash C GoelSouth GA State College

Douglas, GA 31533

Copyright © Cengage Learning. All rights reserved. 8 | 2

You learned how the organization of the periodic table can be explained by the periodicity of the ground-state configurations of the elements. Now we will look at various aspects of the periodicity of the elements.

PeriodicProperties

of the Elements


Development of Periodic Table

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


Periodic Table

• Both Scientists noted:

Similar chemical and physical properties recur periodically when elements are arranged in order of increasing atomic weight.

• Mendeleev got more credit for advancing his ideas and stimulating new work.


Mendeleev’s periodic table generally organized elements by increasing atomic mass and with similar properties in columns. In some places, there were missing elements whose properties he predicted.

When gallium, scandium, and germanium were isolated and characterized, their properties were almost identical to those predicted by Mendeleev for eka-aluminum, eka-boron, and eka-silicon, respectively.


Periodic law states that when the elements are arranged by atomic number, their physical and chemical properties vary periodically.

We will look in more detail at three periodic properties:

atomic radius,

ionization energy, and

electron affinity.


Effective Nuclear Charge Effective nuclear charge is the positive charge that an electron experiences from the nucleus. It is equal to the nuclear charge, but is reduced by shielding or screening from any intervening electron distribution (inner shell electrons).


Effective nuclear charge increases across a period. Because the shell number (n) is the same across a period, each successive atom experiences a stronger nuclear charge. As a result, the atomic size decreases across a period.

PeriodicProperties

of the Elements


What Is the Size of an Atom?

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


PeriodicProperties

of the Elements


Sizes of Atoms

The bonding atomic radius tends to

— Decrease from left to right across a row (due to increasing Zeff).

— Increase from top to bottom of a column (due to the increasing value of n).

© 2012 Pearson Education, Inc.Chemistry, The Central Science, 12th EditionTheodore L. Brown; H. Eugene LeMay, Jr.; Bruce E. Bursten; Catherine J. Murphy; and Patrick Woodward

Sample Exercise 7.1 Bond Lengths in a Molecule

Solve

C—S bond length = bonding atomic radius of C + bonding atomic radius of S = 0.77Å + 1.02Å = 1.79Å

C—H bond length = 0.77Å + 0.37Å = 1.14ÅS—H bond length = 1.02Å = 0.37Å = 1.39Å

Solution

Natural gas used in home heating and cooking is odorless. Because natural gas leaks pose the danger of explosion or suffocation, various smelly substances are added to the gas to allow detection of a leak. One such substance is methyl mercaptan, CH3SH. Use Figure 7.6 to predict the lengths of the C—S, C—H, and S—H bonds in this molecule.

PeriodicProperties

of the Elements


Sizes of Ions

• Ionic size depends upon– The nuclear

charge.– The number of

electrons.– The orbitals in

which electrons reside.

PeriodicProperties

of the Elements


Sizes of Ions• Cations are smaller than their parent

atoms:

The outermost electron is removed and repulsions between electrons are reduced.

• Anions are larger than their parent atoms” Electrons are added and repulsions between electrons are increased.

• Ions increase in size as you go down a column:

This increase in size is due to the increasing value of n.


Sample Exercise 7.3 Atomic and Ionic Radii

Cations are smaller than their parent atoms, and so Ca2+ < Ca. Because Ca is below Mg in group 2A, Ca2+ is larger than Mg2+. Consequently, Ca > Ca2+ > Mg2+.

Solution

Arrange Mg2+, Ca2+, and Ca in order of decreasing radius..

PeriodicProperties

of the Elements


Sizes of Ions

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

• Ionic size decreases with an increasing nuclear charge.


Sample Exercise 7.4 Ionic Radii in an Isoelectronic Series

Arrange the ions K+, Cl–, Ca2+, and S2– in order of decreasing size.

Solution

This is an isoelectronic series, with all ions having 18 electrons. In such a series, size decreases as nuclear charge (atomic number) increases. The atomic numbers of the ions are S 16, Cl 17, K 19, Ca 20. Thus, the ions decrease in size in the order S2– > Cl– > K+ > Ca2+.

PeriodicProperties

of the Elements


Ionization Energy

• The ionization energy is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion.– The first ionization energy is that energy

required to remove the first electron.– The second ionization energy is that

energy required to remove the second electron, etc.

PeriodicProperties

of the Elements


Ionization Energy• It requires more energy to remove each

successive electron.• When all valence electrons have been removed,

the ionization energy takes a quantum leap.


PeriodicProperties

of the Elements


Trends in First Ionization 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.

• Generally, as one goes across a row, it gets harder to remove an electron.– As you go from left to right, Zeff increases.

PeriodicProperties

of the Elements


Trends in First Ionization EnergiesHowever, there are two apparent discontinuities in this trend.

• The first occurs between Groups IIA and IIIA (Be and B)• In this case the electron is removed from a p orbital rather

than an s orbital.– The electron removed is farther from the nucleus.– There is also a small amount of repulsion by the s

electrons.

• The second discontinuity occurs between Groups VA and VIA (N and O).– The electron removed comes from a doubly occupied

orbital.– Repulsion from the other electron in the orbital aids

in its removal.

PeriodicProperties

of the Elements


Electronic Configuration of Ions

• Li (1s22s1)→ Li+ (1s2) + e-

• Fe([Ar]3d64s2)→ Fe2+ ([Ar]3d6) + 2e-

• Fe2+ ([Ar]3d6) → Fe3+ ([Ar]3d5) + e-

• Sn([Kr]4d105s25p2 → Sn2+([Kr]4d105s2)

• Sn2+([Kr]4d105s2)→Sn4+([Kr]4d10)

• F(1s22s22p5) + e- → F-(1s22s22p6)


Sample Exercise 7.7 Electron Configurations of Ions

(a) Calcium (atomic number 20) has the electron configuration [Ar]4s2.

To form a 2+ ion, the two outer electrons must be removed, giving an ion that is isoelectronic with Ar: Ca2:[Ar] (b) Cobalt (atomic number 27) has the electron configuration [Ar]3d74s2.

To form a 3+ ion, three electrons must be removed. As discussed in the text, the 4s electrons are removed before the 3d electrons.

Co3+:[Ar]3d6 (c) Sulfur (atomic number 16) has the electron configuration [Ne]3s23p4.

To form a 2– ion, two electrons must be added. There is room for two additional electrons in the 3p orbitals.

S2–:[Ne]3s23p6 = [Ar]

Solution

Write the electron configuration for (a) Ca2+, (b) Co3+, and (c) S2–.

PeriodicProperties

of the Elements


Electron Affinity

Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom:

Cl + e− Cl−

PeriodicProperties

of the Elements


Trends in Electron Affinity

PeriodicProperties

of the Elements


Trends in Electron Affinity

In general, electron affinity becomes more

exothermic as you go from left to right across a row.

There are again, however, two discontinuities in this trend.

PeriodicProperties

of the Elements


Trends in Electron Affinity• The first occurs between Groups IA and IIA.

– The added electron must go in a p orbital, not an s orbital.– The electron is farther from the nucleus and feels

repulsion from the s electrons.

• The second discontinuity occurs between Groups IVA and VA.– Group VA has no empty orbitals.– The extra electron must go into an already occupied

orbital, creating repulsion.

PeriodicProperties

of the Elements


Properties of Metal, Nonmetals,and Metalloids

PeriodicProperties

of the Elements


Metals versus Nonmetals

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

PeriodicProperties

of the Elements


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

PeriodicProperties

of the Elements


Metals• Compounds formed between metals and

nonmetals tend to be ionic.• Metal oxides tend to be basic. Metals have

low IE


Sample Exercise 7.8 Metal Oxides

(a) Because scandium oxide is the oxide of a metal, we expect it to be an ionic solid. Indeed it is, with the very high melting point of 2485ºC.(b) In compounds, scandium has a 3+ charge, Sc3+, and the oxide ion is O2–. Consequently, the formula of scandium oxide is Sc2O3. Metal oxides tend to be basic and, therefore, to react with acids to form a salt plus water. In this case the salt is scandium nitrate, Sc(NO3)3:

Sc2O3(s) + 6 HNO3(aq) → 2 Sc(NO3)3(aq) + 3 H2O(l)

Solution

(a) Would you expect scandium oxide to be a solid, liquid, or gas at room temperature?(b) Write the balanced chemical equation for the reaction of scandium oxide with nitric acid.

PeriodicProperties

of the Elements


Nonmetals

• Nonmetals are dull, brittle substances that are poor conductors of heat and electricity.

• They tend to gain electrons in reactions with metals to acquire a noble-gas configuration.

PeriodicProperties

of the Elements


Nonmetals

• Substances containing only nonmetals are molecular compounds.

• Most nonmetal oxides are acidic.


Sample Exercise 7.9 Nonmetal Oxides

Write the balanced chemical equation for the reaction of solid selenium dioxide, SeO2(s), with (a) water, (b) aqueous sodium hydroxide.

(a) The reaction between selenium dioxide and water is like that between carbon dioxide and water (Equation 7.13):

SeO2(s) + H2O(l) → H2SeO3(aq)(It does not matter that SeO2 is a solid and CO2 is a gas under ambient conditions; the point is that both are water-soluble nonmetal oxides.)(b) The reaction with sodium hydroxide is like the reaction in Equation 7.15:

SeO2(s) + 2 NaOH(aq) → Na2SeO3(aq) + H2O(l)

Solution

PeriodicProperties

of the Elements


Metalloids

• Metalloids have some characteristics of metals and some of nonmetals.

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

PeriodicProperties

of the Elements


Group Trends

PeriodicProperties

of the Elements


Alkali Metals

• Alkali metals are soft, metallic solids.

• The name comes from the Arabic word for ashes.

PeriodicProperties

of the Elements


Alkali Metals• They are found only in compounds in nature,

not in their elemental forms.• They have low densities and melting points.• They also have low ionization energies.

PeriodicProperties

of the Elements


Alkali Metals

Their reactions with water are famously exothermic.

PeriodicProperties

of the Elements


Alkali Metals• Alkali metals (except Li) react with oxygen to form

peroxides.• K, Rb, and Cs also form superoxides:

K + O2 KO2

• They produce bright colors when placed in a flame.


Sample Exercise 7.10 Reactions of an Alkali Metal

Write a balanced equation for the reaction of cesium metal with (a) Cl2(g), (b) H2O(l), (c) H2(g).

(a) The reaction between Cs and Cl2 is a simple combination reaction between a metal and a nonmetal, forming the ionic compound CsCl: 2 Cs(s) + Cl2(g) → 2 CsCl(s)

(b) 2Cs(s) + 2 H2O(l) → 2 CsOH(aq) + H2(g)

(c) 2 Cs(s) + H2(g) → 2 CsH(s)

All three reactions are redox reactions where cesium forms a Cs+ ion. The Cl–, OH–, and H– are all ions,

which means the products have 1:1 stoichiometry with Cs+.

Solution

PeriodicProperties

of the Elements


Alkaline Earth Metals

• Alkaline earth metals have higher densities and melting points than alkali metals.

• Their ionization energies are low, but not as low as those of alkali metals.

PeriodicProperties

of the Elements


Alkaline Earth Metals

• Beryllium does not react with water, and magnesium reacts only with steam, but the other alkaline earth metals react readily with water.

• Reactivity tends to increase as you go down the group.

PeriodicProperties

of the Elements


Group 6A

• Oxygen, sulfur, and selenium are nonmetals.• Tellurium is a metalloid.• The radioactive polonium is a metal.

PeriodicProperties

of the Elements


Sulfur

• Sulfur is a weaker oxidizer than oxygen.

• The most stable allotrope is S8, a ringed molecule.

PeriodicProperties

of the Elements


Group VIIA: Halogens

• The halogens are prototypical nonmetals.• The name comes from the Greek words halos

and gennao: “salt formers.”

PeriodicProperties

of the Elements


Group VIIA: Halogens

• They have large, negative electron affinities.– Therefore, they tend to oxidize

other elements easily.

• They react directly with metals to form metal halides.

• Chlorine is added to water supplies to serve as a disinfectant.

PeriodicProperties

of the Elements


Group VIIIA: Noble Gases

• The noble gases have astronomical ionization energies.

• Their electron affinities are positive.– Therefore, they are relatively unreactive.

• They are found as monatomic gases.

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© 2012 Pearson Education, Inc. Chapter 7 Periodic Properties of the Elements Subhash C Goel South GA State College Douglas, GA 31533