Chapter 7 – Periodic Properties of the Elements
7.1 Development of the Periodic Table
 Majority of elements discovered 1735-1843
 Elements arranged to reflect trends in chemical and physical properties
 Periodic table arose from periodic patterns in electronic configurations
o Valence orbitals
o Trends in rows and columns
 Mendeleev and Meyer
 Modern periodic table
7.2 Effective Nuclear Charge
Effective nuclear charge (Zeff) 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 because of the effect of the inner electrons.
 The electron is attracted to the nucleus, but repelled by electrons that shield or screen it from the full nuclear charge.
 This shielding is called the screening effect.
 The nuclear charge experienced by an electron depends on its distance from the nucleus and the number of electrons
in the spherical volume out to the electron in question.
 As the average number of screening electrons (S) increases, the effective nuclear charge (Zeff) decreases.
Zeff = Z – S

As the distance from the nucleus increases, S increases and Zeff decreases.

S is called the screening constant which represents the portion of the nuclear charge that is screened from the
valence electron by other electrons in the atom.

The value of S is usually close to the number of core electrons in an atom.
7.3 Sizes of Atoms and Ions
Atomic Sizes
 Consider a collection of Ar atoms in the gas phase
o When undergoing collisions, atoms ricochet apart
o Apparent radii determined by closest distances separating the nuclei during collisions
o Nonbonding radius
 Consider a diatomic molecule
o Bond distance
o Bonding atomic radius
 Shorter than nonbonding radius
Prac. Prob. 7.1
 Atomic size varies consistently throughout periodic table
o Becomes larger down a group
o Becomes smaller across a period
Fig. 7.6
o Two factors at work
 Principal quantum number
 Zeff
Prac. Prob. 7.2
Periodic Trends in Ionic Radii





Ionic size is important:
in predicting lattice energy.
in determining the way in which ions pack in a solid.
Just as atomic size is periodic, ionic size is also periodic.
In general:

Cations are smaller than their parent atoms.

Electrons have been removed from the most spatially extended orbital.

The effective nuclear charge has increased.

Therefore, the cation is smaller than the parent atom.

Anions are larger than their parent atoms.



Electrons have been added to the most spatially extended orbital.

This means total electron-electron repulsion has increased.

Therefore, anions are larger than their parent atoms.
For ions with the same charge, ionic size increases down a group.
isoelectronic series •
As nuclear charge increases in an isoelectronic series the ions become smaller:
O2– > F– > Na+ > Mg2+ > Al3+
Prac. Prob. 7.3, 7.4
7.4 Ionization Energy
 Ionization energy
 First ionization energy, I1
 Second ionization energy, I2
 The larger the ionization energy, the more difficult it is to remove an electron
 Sharp increase in ion. E when core electron is removed
Table 7.2, p.271
Prac. Prob. 7.5
Periodic Trends in Ionization Energy
 Ionization energy decreases down a group
o Outer electrons removed most readily
 Ionization energy generally increases across a period
o Increased Zeff
o Two exceptions
 Removing first p electron  s2p0
 Removing fourth p electron s2p3
Fig. 7.10
Prac. Prob. 7.6
Electron Configurations of Ions
 These are derived from the electron configurations of elements with the required number of electrons added or
removed from the most accessible orbital.

Li:
[He]2s1
becomes Li+:
[He]
2
5

F:
[He]2s 2p
becomes F :
[He]2s22p6 = [Ar]
 Transition metals tend to lose the valence shell electrons first and then as many d electrons as are required to
reach the desired charge on the ion.

Thus electrons are removed from 4s before the 3d, etc.
These are derived from the electron configurations of elements with the required number of electrons added or
removed from the most accessible orbital.

Li:
[He]2s1
becomes Li+:
[He]
2
5

F:
[He]2s 2p
becomes F :
[He]2s22p6 = [Ar]
 Transition metals tend to lose the valence shell electrons first and then as many d electrons as are required to
reach the desired charge on the ion.

Thus electrons are removed from 4s before the 3d, etc.
Prac. Prob. 7.7
7.5 Electron Affinities
 Electron affinity
 Can be exothermic (-) or endothermic (+)
o More (-) indicates more favorable addition of electron
 Use electron configurations to determine (-) or (+)
o Ex. Addition of electron to Ar vs. Cl
 Does not vary greatly down a group
Fig. 7.12
7.6 Metals, Nonmetals, and Metalloids
Metals
 Metallic character
o Increases down a group
o Decreases across a period
 Shiny, lustrous, malleable, ductile
 Metallic oxides form basic solids
o Most metal oxides are basic
metal oxide + water  metal hydroxide
o metal oxide + acid  salt + water
o Most neutral metals are oxidized to form characteristic cations
 Group 1 forms M+
 Group 2 forms M2+
 Transition metals vary
 Low ionization energies
Prac. Prob. 7.8
Nonmetals
 In general: nonlustrous, poor conductors of heat and electricity, lower melting points
 Diatomic molecular gases
 Reaction with metals: loss of electrons to form ionic compounds
o Metal + nonmetal  salt
o Most nonmetal oxides are acidic
Nonmetal oxide + water  acid
Prac. Prob. 7.9
o Nonmetal oxide + base  salt + water
Sample Prob. 7.9
Metalloids
 Properties intermediate between metals and nonmetals
 Semiconductors, integrated circuits, computer chips
7.7 Group Trends for the Active Metals
Group 1: The Alkali Metals
 Alkali metals are soft
 Tend to lose a single electron
 Reactivity increases down a the group
 Melting point decreases down the group
 React with H to form hydrides
o H is present as the hydride ion, Ho React with water  MOH and H2 gas
Very exothermic!
o Produce different oxides with O2
 Oxides (O2-): Li
 Peroxides (O22-): Na, K, Rb, Cs
 Superoxides (O21-): K, Rb, Cs
 Emit characteristic flame colors
Prac. Prob. 7.10
Group 2A: The Alkaline Earth Metals
 Harder, denser, higher m.p. than Group 1
 Not as reactive as Group 1
 Tend to lose 2 electrons
 Reactivity increases down the group
o Be – no rxn with water or steam
o Mg – only reacts with steam
o Ca and below – react with water to give a hydroxide and hydrogen
o Can all react with O2
 Characteristic flame colors
Transition Metals – good electrical conductors
7.8 Group Trends for Selected Nonmetals
Hydrogen
 A unique element
 Most often colorless, diatomic gas
 Can gain an electron  H- (hydride ion)
 Can lose an electron  H+ (proton)
Group 6A: The Oxygen Group
 Metallic character increases down group
 Two important forms of oxygen: O2, O3
o O2 and O3 are allotropes
o O3 prepared from O2 (use electricity)
o O3 is pungent and toxic
o O2 is a strong oxidizing agent
o Sulfur can react with metals
 Most common form, S8
 Tends to form S2- in compounds
Group 7A: The Halogens
 “salt formers”
 tend to gain 1 electron
 F is one of most reactive elements
 All exist as diatomic molecules (except At)
 Cl is the most industrially useful – produced by electrolysis of brine
 Cl and water  HCl and HOCl
 H compounds of halogens are all strong acids except HF
 React with most metals  metal halide
Group 8: The Noble Gases
 All nonmetals and monatomic
 Unreactive (inert) due to full outer shell
 1962 – first compounds prepared (XeF2, XeF4, XeF6) – none for He, Ne,
 2000- HArF first compound with Ar formed; KrF2 only known Kr compound
 Rn – too radioactive to study
*Most ionic compounds are solids
*Most covalent (molecular) compounds are liquids or gases (esp. with oxygen)