Periodic Properties of the Elements – Chapter 7 Study Notes

Periodic Properties of the Elements

Development of the Periodic Table

The periodic table organizes elements based on recurring chemical properties. Dmitri Mendeleev and Lothar Meyer independently recognized that elements could be grouped by their properties, leading to the modern periodic table.

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

• Later, Ernest Rutherford discovered the nuclear structure of the atom.

• Henry Moseley established the concept of atomic number (number of protons) as the basis for periodic properties.

Periodicity

Periodicity refers to the repetitive pattern of properties of 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

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 in which electrons reside

Cations vs. Anions

• Cations are smaller than their parent atoms because the outermost electron is removed and electron repulsions are reduced.

• Anions are larger than their parent atoms due to added electrons and increased repulsions.

Isoelectronic Series

An isoelectronic series consists of ions with the same number of electrons. Within such a series:

• Ionic size decreases as nuclear charge increases.

• Increasing nuclear charge leads to decreasing ionic radius as atomic number increases.

Ion	Protons	Electrons	Ionic Radius (pm)
N3−	7	10	146
O2−	8	10	140
F−	9	10	133
Na+	11	10	102
Mg2+	12	10	72
Al3+	13	10	53

Ionization Energy

Definition and Trends

Ionization energy (I) 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.

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

It requires more energy to remove each successive electron, especially after all valence electrons have been removed (core electrons).

Periodic Trends in Ionization Energy

• I1 generally increases to the right across a period.

• I1 generally decreases down a group.

Factors influencing ionization energy:

• Smaller atoms have higher I values.

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

There are irregularities in the general trend due to electron configurations.

Electron Affinity

Definition and Trends

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

$\mathrm{Cl} + e^- \rightarrow \mathrm{Cl}^-$

• Typically exothermic (negative value) for most elements.

• Across a period, electron affinity generally increases (becomes more negative) to the right, with exceptions.

• Usually, not much change in a group.

Metals, Nonmetals, and Metalloids

Classification and Properties

• Metals tend to form cations.

• Nonmetals tend to form anions.

• Hydrogen is a nonmetal with special properties.

Properties of Metals

• Usually shiny

• Conduct heat and electricity

• Ductile, flexible

• Solid at room temperature (except mercury)

• Low ionization energies; form cations easily

Metal Chemistry

• Compounds with nonmetals tend to be ionic.

• Metal oxides are basic and react with acids.

Properties of Nonmetals

• Found on the right side of the periodic table

• Can be solid, liquid, or gas

• Solids are dull, brittle, poor conductors

• Large negative electron affinity; form anions readily

Nonmetal Chemistry

• Substances with 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
Solids are malleable and ductile	Solids are brittle, hard, some are soft
Good conductors of heat and electricity	Poor conductors of heat and electricity
Most metal oxides are ionic and basic	Most nonmetal oxides are molecular and 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 (used in computer chips)

Group Trends

Overview of Group Properties

Elements in a group share similar properties. Notable groups include:

• Group 1A: Alkali metals

• Group 2A: Alkaline earth metals

• Group 6A: Oxygen group

• Group 7A: Halogens

• Group 8A: Noble gases

• Hydrogen is a nonmetal

Alkali Metals

• Soft, metallic solids

• Found only in compounds in nature

• Exhibit typical metallic properties (luster, conductivity)

• Low densities and melting points

• Low ionization energies

Element	Electron Configuration	Melting Point (°C)	Density (g/cm³)	Atomic Radius (pm)	I1 (kJ/mol)
Lithium	[He]2s1	181	0.53	128	520
Sodium	[Ne]3s1	98	0.97	166	496
Potassium	[Ar]4s1	63	0.86	203	419
Rubidium	[Kr]5s1	39	1.53	220	403
Cesium	[Xe]6s1	28	1.88	244	376

• Reactions with water are famously exothermic.

• Characteristic flame colors due to electronic transitions.

Alkaline Earth Metals

• Higher densities and melting points than alkali metals

• Low ionization energies (not as low as alkali metals)

• Readily form +2 cations by losing 2 valence electrons

• Beryllium does not react with water; magnesium reacts only with steam; others react readily with water

• Reactivity increases down the group

Element	Electron Configuration	Melting Point (°C)	Density (g/cm³)	Atomic Radius (pm)	I1 (kJ/mol)
Beryllium	[He]2s2	1287	1.85	112	900
Magnesium	[Ne]3s2	650	1.74	160	738
Calcium	[Ar]4s2	842	1.55	197	590
Strontium	[Kr]5s2	777	2.63	215	549
Barium	[Xe]6s2	727	3.51	222	503

Group 6A: Oxygen Group

• Oxygen, sulfur, selenium: nonmetals

• Tellurium: metalloid

• Polonium: metal

• Increasing metallic character down the group

• Oxygen forms 2− anion; polonium likely to have a positive charge

Element	Type
Oxygen	Nonmetal
Sulfur	Nonmetal
Selenium	Nonmetal
Tellurium	Metalloid
Polonium	Metal

Allotropes of Oxygen

• Oxygen gas: O2 (dioxygen)

• Ozone gas: O3

Group 7A: Halogens

• Typical nonmetals

• Highly negative electron affinities; exist as anions in nature

• React directly with metals to form metal halides

Element	Electron Configuration	Melting Point (°C)	Density (g/cm³)	Atomic Radius (pm)	I1 (kJ/mol)
Fluorine	[He]2s22p5	-220	1.70	57	1681
Chlorine	[Ne]3s23p5	-101	3.12	102	1251
Bromine	[Ar]4s24p5	-7	3.12	120	1140
Iodine	[Kr]5s25p5	114	4.94	139	1008

Group 8A: Noble Gases

• Very large ionization energies

• Positive electron affinities (cannot form stable anions)

• Relatively unreactive

• Found as monatomic gases

Element	Boiling Point (K)	Density (g/L)	Atomic Radius (pm)	I1 (kJ/mol)
Helium	4.2	0.18	31	2372
Neon	27.1	0.90	38	2081
Argon	87.3	1.78	71	1521
Krypton	120.0	3.75	88	1351
Xenon	165.0	5.89	108	1170

Hydrogen

• 1s1 electron configuration, similar to alkali metals

• Higher ionization energy than any metal

• Nonmetal; occurs as colorless diatomic gas H2(g)

• Reacts with most nonmetals to form molecular compounds, not ionic compounds