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 foundational tool in chemistry, organizing elements according to recurring properties. Its development was shaped by the work of Dmitri Mendeleev and Lothar Meyer, who independently recognized patterns among the elements.

• Mendeleev's Table: Arranged elements by increasing atomic mass and grouped them by similar chemical properties.

• Predictions: Mendeleev predicted the existence and properties of undiscovered elements, such as "eka-silicon" (later identified as germanium).

Property	Mendeleev's Prediction (Eka-Silicon)	Observed (Germanium)
Atomic weight	72	72.59
Density (g/cm³)	5.5	5.35
Melting point (°C)	High	947
Color	Dark gray	Grayish white
Formula of oxide	XO₂	GeO₂
Boiling point of chloride (°C)	A little under 100	84

Additional info: Mendeleev's use of chemical properties allowed him to predict missing elements and their characteristics.

Atomic Number

The concept of atomic number refined the organization of the periodic table. Henry Moseley established atomic number as the number of protons in the nucleus, which became the basis for periodic properties.

• Mendeleev: Used atomic mass for arrangement.

• Moseley: Used atomic number, enabling identification of gaps and discovery of new elements.

Periodicity

Periodicity refers to the repeating patterns of properties among elements as a function of atomic number. Key periodic properties include:

• Sizes of atoms and ions

• Ionization energy

• Electron affinity

• Chemical property trends within groups

These trends are largely explained by effective nuclear charge.

7.2 Effective Nuclear Charge

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

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

• The nuclear charge is "screened" by core electrons.

Calculation:

• $Z_{eff} = Z - S$

• Where $Z$ is the atomic number and $S$ is the screening constant (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. Two important measures are:

• Nonbonding atomic radius (van der Waals radius): Half the shortest distance between nuclei during atomic collisions.

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

Trends:

• Bonding atomic radius decreases from left to right across a period ($Z_{eff}$ increases).

• Bonding atomic radius increases from top to bottom within a group (principal quantum number $n$ increases).

Sizes of Ions—Ionic Radii

Ionic radii are determined by interatomic distances in ionic compounds. Ionic size depends on:

• Nuclear charge

• Number of electrons

• Orbitals occupied by valence electrons

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

When forming ions, electrons are added or removed from the highest energy level:

• Cations: Electrons are lost from the highest principal quantum number ($n$).

• Example: $\text{Li: } 1s^2 2s^1 \rightarrow \text{Li}^+: 1s^2$

• Example: $\text{Fe: } 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^6 \rightarrow \text{Fe}^{2+}: 1s^2 2s^2 2p^6 3s^2 3p^6 3d^6$

• Anions: Electrons are added to fill $ns^2 np^6$.

• Example: $\text{F: } 1s^2 2s^2 2p^5 \rightarrow \text{F}^-: 1s^2 2s^2 2p^6$

Isoelectronic Series and Ion Size

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

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

7.4 Ionization Energy and Electron Affinity

Ionization energy (I) 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.

• Higher ionization energy means it is more difficult to remove an electron.

Element	$I_1$	$I_2$	$I_3$	$I_4$	$I_5$	$I_6$
Na	496	4562	6912	9544	13354	16613
Mg	738	1451	7733	10542	13630	18020
Al	578	1820	2740	11577	14842	18374
Si	786	1577	3232	4356	16091	19805
P	1012	1907	2914	4956	6274	21267
S	999	2252	3357	4556	7004	8495
Cl	1251	2298	3821	5157	7298	9371
Ar	1521	2666	3931	5771	7230	11995

Periodic Trends for First Ionization Energy ($I_1$)

• $I_1$ generally increases across a period.

• $I_1$ generally decreases down a group.

• s- and p-block elements show a larger range of $I_1$ values; d-block increases slowly; f-block shows small 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).

Factors Influencing Ionization Energy

• Smaller atoms have higher ionization energies.

• Ionization energy 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:

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

• Typically exothermic (negative EA) for most elements.

General Trend: Electron affinity generally increases across a period, with exceptions:

• Group 2A: s sublevel is full

• Group 5A: p sublevel is half-full

• Group 8A: p sublevel is full

Additional info: For these exceptions, electron affinity may be positive (unfavorable addition of electron).

7.5 Metals, Nonmetals, and Metalloids

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

Metals

• Form cations

• Shiny luster, conduct heat and electricity

• Malleable and ductile

• Solids at room temperature (except mercury)

• Low ionization energies

Metal Chemistry

• Compounds with nonmetals tend to be ionic

• Metal oxides are basic and react with acids

Nonmetals

• Form anions

• Found on the right side of the periodic table

• Can be solid, liquid, or gas

• Dull, brittle, poor conductors

• Large negative electron affinity

Nonmetal Chemistry

• Substances containing only nonmetals are molecular compounds

• Most nonmetal oxides are acidic

Comparison Table: Metals vs Nonmetals

Metals	Nonmetals
Shiny luster; various colors	No luster; various colors
Malleable and ductile	Brittle; some hard, some soft
Good conductors of heat/electricity	Poor conductors of heat/electricity
Oxides are ionic solids, basic	Oxides are molecular, acidic
Form cations in solution	Form anions/oxyanions in solution

Metalloids

• Exhibit properties intermediate between metals and nonmetals

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

Group Trends

Elements in the same group share similar properties, with trends observed within each group:

• Group 1A: Alkali metals

• Group 2A: Alkaline earth metals

• Group 6A: Oxygen group

• Group 7A: Halogens

• Group 8A: Noble gases

• Hydrogen: Not a metal despite its electron configuration

7.6 Group 1A: Alkali Metals

Alkali metals are soft, metallic solids with typical metallic properties. They are found only in compounds in nature.

• Low densities and melting points

• Low ionization energies

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

Alkali Metal Chemistry

• Reactions with water are highly exothermic

Differences in Alkali Metal Chemistry

• Lithium forms an oxide: $4\text{Li} + \text{O}_2 \rightarrow 2\text{Li}_2\text{O}$ (Oxide ion: O2−)

• Sodium forms a peroxide: $2\text{Na} + \text{O}_2 \rightarrow \text{Na}_2\text{O}_2$ (Peroxide ion: O22−)

• K, Rb, Cs form superoxides: $\text{M} + \text{O}_2 \rightarrow \text{MO}_2$ (Superoxide ion: O2−)

Flame Tests

• Alkali metals produce characteristic flame colors due to electronic transitions.

• Example: Li (red), Na (yellow), K (violet)