Chapter 7: Periodic Properties of the Elements – Study Notes

Development of the Periodic Table

Historical Background

The periodic table is a systematic arrangement of elements based on recurring chemical properties. Dmitri Mendeleev and Lothar Meyer independently proposed grouping elements by their properties, leading to the modern periodic table.

• Mendeleev is credited for using chemical properties to organize the table and predicting missing elements, such as germanium.

• Periodic Law: Properties of elements recur periodically when arranged by increasing atomic number.

Comparison of Predicted and Observed Properties

Mendeleev predicted properties of unknown elements (e.g., eka-silicon) that matched closely with those of germanium discovered later.

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)	0.305	0.32
Melting point (°C)	High	High
Color	Dark gray	Grayish white
Formula of oxide	EO2	GeO2
Density of oxide (g/cm3)	4.7	4.7
Formula of chloride	ECl4	GeCl4
Boiling point of chloride (°C)	A little under 100	86

Atomic Number

Definition and Importance

The atomic number is the number of protons in the nucleus of an atom. Mendeleev's table was originally based on atomic masses, but Henry Moseley later established atomic number as the fundamental property for periodicity.

• Atomic number (Z): Determines the identity of an element.

• Periodic properties are based on atomic number, not atomic mass.

Periodicity

Definition

Periodicity refers to the repetitive pattern of chemical and physical properties of elements as a function of atomic number.

• Key properties discussed: atomic and ionic sizes, ionization energy, electron affinity, and group trends.

Effective Nuclear Charge

Concept and Calculation

Effective nuclear charge (Zeff) is the net positive charge experienced by valence electrons, accounting for both attraction to the nucleus and repulsion by other electrons.

• Calculated as: $Z_{eff} = Z - S$ where Z is the atomic number and S is the screening constant (approximate number of inner electrons).

• Example: For Na ([Ne]3s1), $Z_{eff} = 11 - 10 = 1$ (simple estimate).

• Zeff increases across a period (due to increasing nuclear charge).

• Zeff increases down a group (due to electron cloud diffusion).

Atomic Size

Bonding and Nonbonding Atomic Radius

The bonding atomic radius (covalent radius) is half the distance between nuclei in a bond. The nonbonding atomic radius (van der Waals radius) is half the shortest distance between nuclei during atomic collisions.

• Bonding atomic radius decreases from left to right across a period (Zeff increases).

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

Example Calculation

Given atomic radii: C = 0.76 Å, S = 1.05 Å, H = 0.31 Å. Bond lengths in CH3SH:

• C–S bond: 0.76 Å + 1.05 Å = 1.81 Å

• C–H bond: 0.76 Å + 0.31 Å = 1.07 Å

• S–H bond: 1.05 Å + 0.31 Å = 1.36 Å

Sizes of Ions

Factors Affecting Ionic Size

Ionic size is determined by interatomic distances in ionic compounds and depends on:

• Nuclear charge

• Number of electrons

• Orbitals occupied by electrons

Cations vs. Anions

• Cations are smaller than their parent atoms (loss of outer electron, reduced electron repulsion).

• Anions are larger than their parent atoms (gain of electron, increased electron repulsion).

Electron Configurations of Ions

Rules for Cations and Anions

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

• Example: Fe ([Ar]4s23d6) → Fe2+ ([Ar]3d6) (loss of two 4s electrons).

• Anions: Electron configurations are filled to ns2np6. Example: F ([He]2s22p5) → F− ([He]2s22p6).

Isoelectronic Series and Ion Size

Definition and Trend

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

Ion	Radius (Å)
O2−	1.26
F−	1.19
Na+	1.16
Mg2+	0.86
Al3+	0.68

Ionization Energy (I)

Definition

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

• First ionization energy (I1): Energy to remove the first electron.

• Second ionization energy (I2): Energy to remove the second electron, and so on.

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

Trends in Ionization Energy

• I1 generally increases across a period.

• I1 generally decreases down a group.

• s- and p-block elements show a larger range of I1 values; d- and f-block elements show smaller variations.

Irregularities in Trends

• Trend is not followed when the added valence electron enters a new (higher energy) sublevel.

• Trend is not followed when the added electron is the first to pair in an orbital (electron repulsions lower energy).

Factors Influencing Ionization Energy

• Smaller atoms have higher ionization energies.

• Depends on effective nuclear charge and average distance of the electron from the nucleus.

Electron Affinity

Definition and Trends

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

$\text{Cl} + e^- \longrightarrow \text{Cl}^-$

• Typically exothermic (negative value for most elements).

• Not much change in a group; generally increases across a period.

• Exceptions: Group 2A (full s sublevel), Group 5A (half-full p sublevel), Group 8A (full p sublevel) – these have positive electron affinities (X− is unstable).

Metals, Nonmetals, and Metalloids

Classification

• Metals: Tend to form cations, have shiny luster, conduct heat/electricity, are malleable/ductile, and are solids at room temperature (except mercury).

• Nonmetals: Tend to form anions, may be solid, liquid, or gas, are dull and brittle, poor conductors, and have large negative electronegativities.

• Metalloids: Exhibit properties intermediate between metals and nonmetals; several are semiconductors (e.g., silicon).

Metal Chemistry

• Compounds formed between metals and nonmetals are usually ionic.

• Metal oxides tend to be basic: $\text{Metal oxide} + \text{Water} \rightarrow \text{Metal hydroxide}$ $\text{Na}_2\text{O}(s) + \text{H}_2\text{O}(l) \rightarrow 2\text{NaOH}(aq)$

• Insoluble metal oxides react with acids: $\text{Metal oxide} + \text{Acid} \rightarrow \text{Salt} + \text{Water}$ $\text{NiO}(s) + 2\text{HNO}_3(aq) \rightarrow \text{Ni(NO}_3)_2(aq) + \text{H}_2\text{O}$

Nonmetal Chemistry

• Substances containing only nonmetals are molecular compounds (e.g., SO3, CO2, NO2).

• Nonmetals tend to gain electrons when reacting with metals (form anions).

• Nonmetal oxides tend to be acidic: $\text{Nonmetal oxide} + \text{Water} \rightarrow \text{Acid}$ $\text{CO}_2(g) + \text{H}_2\text{O}(l) \rightarrow \text{H}_2\text{CO}_3(aq)$

• Nonmetal oxides react with bases: $\text{CO}_2(g) + 2\text{NaOH}(aq) \rightarrow \text{Na}_2\text{CO}_3(aq) + \text{H}_2\text{O}$

Metalloids

• Have properties of both metals and nonmetals.

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

Summary Table: Properties of Metals, Nonmetals, and Metalloids

Property	Metals	Nonmetals	Metalloids
Electrical Conductivity	High	Low	Intermediate
Luster	Shiny	Dull	Variable
Malleability/Ductility	Yes	No	Variable
Typical Ion Formed	Cation	Anion	Variable
State at Room Temp	Solid (except Hg)	Solid, liquid, or gas	Solid

Practice Exercises

• Arrange atoms/ions by increasing/decreasing atomic or ionic radius using periodic trends.

• Predict chemical behavior based on group and period trends.

• Write balanced equations for reactions of metals, nonmetals, and their oxides.

Additional info: These notes are based on lecture slides for Chapter 7 of "Chemistry: The Central Science" and cover all major periodic properties relevant to a General Chemistry course.