Chapter 7: Periodic Properties of the Elements – Study Notes

Chapter 7: Periodic Properties of the Elements

7.1 Development of the Periodic Table

The periodic table is a systematic arrangement of elements based on recurring chemical and physical properties. Its development was a major milestone in chemistry, allowing scientists to predict the properties of undiscovered elements.

• Dmitri Mendeleev and Lothar Meyer independently proposed that elements should be grouped according to periodic trends in their properties.

• Mendeleev's table was based on atomic mass, the most fundamental property known at the time.

• He used chemical properties to organize the table and predicted the existence and properties of missing elements (e.g., eka-silicon, later discovered as germanium).

Table 7.1: Comparison of Eka-Silicon (Predicted by Mendeleev) and Germanium (Observed)

Property	Mendeleev's Prediction (Eka-Silicon)	Observed Germanium
Atomic weight	72	72.59
Density (g/cm³)	5.5	5.35
Specific heat (J/g·K)	0.30	0.32
Melting point (°C)	High	947
Color	Dark gray	Grayish white
Formula of oxide	EO2	GeO2
Density of oxide (g/cm³)	4.7	4.70
Formula of chloride	ECl4	GeCl4
Boiling point of chloride (°C)	A little under 100	86

Additional info: Mendeleev's predictions closely matched the observed properties of germanium, validating the periodic table's predictive power.

Atomic Number and Periodicity

• About 35 years after Mendeleev, Ernest Rutherford discovered the nuclear atom.

• Henry Moseley established the concept of atomic number (number of protons) as the basis for periodicity using X-ray experiments.

• Periodicity refers to the repetitive pattern of element properties based on atomic number.

Key properties discussed in this chapter:

• Sizes of atoms and ions

• Ionization energy

• Electron affinity

• Group chemical property trends

7.2 Effective Nuclear Charge (Zeff)

Many atomic properties depend on the attraction between valence electrons and the nucleus, as well as repulsion from other electrons. The effective nuclear charge is the net positive charge experienced by valence electrons.

• Electrons are attracted to the nucleus and repelled by other electrons.

• The screening effect reduces the full nuclear charge felt by outer electrons.

• The effective nuclear charge is calculated as:

$Z_{eff} = Z - S$

• Where Z is the atomic number and S is the screening constant (usually the number of core electrons).

• Trend: $Z_{eff}$ increases from left to right across a period.

7.3 Sizes of Atoms and Ions

The size of an atom is determined by the space occupied by its electrons, which is described by probability distributions.

• Nonbonding atomic radius (van der Waals radius): Half the shortest distance between nuclei during a collision of atoms.

• Bonding atomic radius (covalent radius): Half the distance between nuclei in a bond.

Trends in Bonding Atomic Radius:

• Decreases from left to right across a period (due to increasing $Z_{eff}$).

• Increases from top to bottom within a group (due to increasing principal quantum number, n).

Sizes of Ions—Ionic Radii

• Determined by interatomic distances in ionic compounds.

• Ionic size depends on:

  • Nuclear charge

  • Number of electrons

  • Orbitals in which valence electrons reside

• Cations are smaller than their parent atoms (loss of electrons reduces repulsion).

• Anions are larger than their parent atoms (gain of electrons increases repulsion).

Electron Configurations of Ions

• Cations: Electrons are lost from the highest energy level (largest n value).

• Anions: Electron configurations are filled to $ns^2np^6$ (octet rule).

• Example: $\text{F}^-: 1s^2 2s^2 2p^6$

Isoelectronic Series and Ion Size

An isoelectronic series consists of ions with the same number of electrons. Within a series, ionic size decreases as nuclear charge increases.

	O2-	F-	Na+	Mg2+	Al3+
Protons	8	9	11	12	13
Electrons	10	10	10	10	10
Ionic Radius (Å)	1.26	1.19	1.16	0.86	0.68

7.4 Ionization Energy and Electron Affinity

Ionization Energy (I)

Ionization energy is the minimum energy required to remove an electron from a gaseous atom or ion.

• First ionization energy ($I_1$): Energy to remove the first electron.

• Second ionization energy ($I_2$): Energy to remove the second electron.

• Each successive ionization energy is higher than the previous one.

• Large jumps in ionization energy occur when removing core (inner-shell) electrons.

Table 7.2: Successive Ionization Energies (kJ/mol) for Na to Ar

Element	$I_1$	$I_2$	$I_3$	$I_4$	$I_5$	$I_6$
Na	496	4562	6910	9544	13354	16613
Mg	738	1451	7733	10542	13630	18020
Al	578	1817	2745	11577	14842	18379
Si	786	1577	3232	4361	16091	19805
P	1012	1907	2914	4964	6274	21267
S	1000	2251	3357	4556	7004	8495
Cl	1251	2298	3822	5158	6549	9444
Ar	1521	2666	3931	5771	7238	11811

Periodic Trends in Ionization Energy

• $I_1$ generally increases across a period (left to right).

• $I_1$ generally decreases down a group (top to bottom).

• s- and p-block elements show a larger range of $I_1$ values; d- and f-block elements show smaller variations.

Irregularities: The trend is not followed when the added valence electron enters a new sublevel or is the first to pair in an orbital (due to electron repulsions).

• Example: Oxygen's $I_1$ is lower than nitrogen's due to electron pairing in the 2p orbital.

Factors Influencing Ionization Energy:

• Smaller atoms have higher ionization energies.

• Depends on effective nuclear charge and average electron-nucleus distance.

Electron Affinity (EA)

Electron affinity is the energy change when an electron is added to a gaseous atom:

$\text{Cl}(g) + e^- \rightarrow \text{Cl}^-(g) \qquad EA = -349\ \text{kJ}$

• Typically exothermic (negative value), but not always.

General Trend:

• EA generally increases (becomes more negative) across a period, with exceptions:

  • Group 2A: s sublevel is full

  • Group 5A: p sublevel is half-full

  • Group 8A: p sublevel is full (noble gases)

• For many elements, EA is positive (adding an electron is unfavorable).

7.5 Metals, Nonmetals, and Metalloids

Elements are classified as metals, nonmetals, or metalloids based on their physical and chemical properties.

Metals

• Most elements are metals.

• Properties:

  • Shiny luster

  • Good conductors of heat and electricity

  • Malleable and ductile

  • Solids at room temperature (except mercury)

  • Low ionization energies; form cations easily

Metal Chemistry:

• Form ionic compounds with nonmetals.

• Metal oxides are basic and react with acids.

Nonmetals

• Found on the right side of the periodic table.

• Properties:

  • Can be solid, liquid, or gas

  • Solids are dull, brittle, poor conductors

  • Large negative electron affinities; form anions readily

Nonmetal Chemistry:

• Substances containing only nonmetals are molecular compounds.

• Most nonmetal oxides are acidic.

Comparison of Metals and Nonmetals

Metals	Nonmetals
Shiny luster; various colors	No luster; various colors
Solids are malleable and ductile	Solids are usually brittle
Good conductors of heat and electricity	Poor conductors of heat and electricity
Most metal oxides are basic	Most nonmetal oxides are acidic
Tend to form cations in aqueous solution	Tend to form anions or oxyanions in aqueous solution

Metalloids

• Have properties intermediate between metals and nonmetals.

• Several are electrical semiconductors (e.g., silicon in computer chips).

7.6 Group Trends

Elements in the same group (vertical column) have similar properties. Important groups include:

• Group 1A: Alkali metals

• Group 2A: Alkaline earth metals

• Group 6A: Oxygen group

• Group 7A: Halogens

• Group 8A: Noble gases

• Hydrogen (unique properties)

7.6 Group 1A: Alkali Metals

• Soft, metallic solids; not found in elemental form in nature.

• Exhibit typical metallic properties (luster, conductivity).

• Low densities and melting points; low ionization energies.

Table 7.4: Properties of Alkali Metals

Element	Electron Configuration	Melting Point (°C)	Density (g/cm³)	Atomic Radius (Å)	$I_1$ (kJ/mol)
Lithium	[He]2s1	181	0.53	1.28	520
Sodium	[Ne]3s1	98	0.97	1.66	496
Potassium	[Ar]4s1	63	0.86	2.03	419
Rubidium	[Kr]5s1	39	1.53	2.20	403
Cesium	[Xe]6s1	28	1.88	2.44	376

• Reactions with water are highly exothermic.

• Flame tests can be used to identify alkali metals by their characteristic colors (due to electronic transitions).

Differences in Alkali Metal Chemistry:

• Lithium reacts with oxygen to form an oxide: $4\text{Li} + \text{O}_2 \rightarrow 2\text{Li}_2\text{O}$ (oxide ion: $\text{O}^{2-}$)

• Sodium forms a peroxide: $2\text{Na} + \text{O}_2 \rightarrow \text{Na}_2\text{O}_2$ (peroxide ion: $\text{O}_2^{2-}$)

• Potassium, rubidium, and cesium form superoxides: $\text{M} + \text{O}_2 \rightarrow \text{MO}_2$ (superoxide ion: $\text{O}_2^-$)

Additional info: The periodic properties discussed here are foundational for understanding chemical reactivity, bonding, and the organization of the periodic table. Mastery of these trends is essential for predicting the behavior of elements in chemical reactions.