Periodic Properties of the Elements – CHEM 1010 Study Notes

Periodic Properties of the Elements

Objectives

This chapter explores the periodic table, electron configurations, and the periodic trends that govern atomic and ionic properties. Understanding these concepts is essential for predicting chemical behavior and reactivity.

• Write electron configurations for elements.

• Define and explain atomic radius, ionic radius, ionization energy, and electron affinity.

• Describe periodic trends across periods and within groups.

• Predict physical properties of elements based on periodic trends.

9.2 The Periodic Table – Dmitri Mendeleev

Development and Structure

The periodic table organizes elements by increasing atomic number and recurring chemical properties. Dmitri Mendeleev created the first periodic table, arranging elements by mass and properties, and predicted the existence of undiscovered elements.

• Modern table: 118 elements arranged by atomic number.

• Groups (columns): Elements with similar properties.

• Periods (rows): Elements with increasing atomic number.

9.4 Electron Configurations and Valence Electrons

Definitions and Principles

Electron configuration describes the arrangement of electrons in an atom's orbitals. Valence electrons are those in the outermost shell and are crucial for chemical reactivity.

• Main group elements: Valence electrons are in the outermost shell.

• Transition elements: Outermost d electrons are included in the valence count.

• Core electrons: Electrons in lower energy shells.

• Electron configuration notation: Sublevels listed in order of filling, with electron count as superscript.

Example: For Si (Z=14): $1s^2 2s^2 2p^6 3s^2 3p^2$

Shorthand notation: Use noble gas in brackets for core electrons. For Rb (Z=37): $[Kr] 5s^1$

Orbital Diagrams and Filling Order

• Energy levels fill from lowest to highest: $s \rightarrow p \rightarrow d \rightarrow f$

• Aufbau principle: Electrons occupy lowest energy orbitals first.

Symbol	Number of Electrons	Electron Configuration	Orbital Diagram
Li	3	1s22s1	[↑↓][↑]
Be	4	1s22s2	[↑↓][↑↓]
B	5	1s22s22p1	[↑↓][↑↓][↑]
C	6	1s22s22p2	[↑↓][↑↓][↑][↑]

Blocks of the Periodic Table

• s-block: Groups 1A and 2A, plus He.

• p-block: Groups 3A to 8A.

• d-block: Transition metals (Groups 3B to 2B).

• f-block: Lanthanides and actinides.

Valence Electrons and Ion Formation

• Group 1A: 1 valence electron, forms 1+ ions.

• Group 2A: 2 valence electrons, forms 2+ ions.

• Group 7A: 7 valence electrons, forms 1- ions.

• Group 6A: 6 valence electrons, forms 2- ions.

Group	Common Ion Charge
1A	+1
2A	+2
6A	-2
7A	-1

Transition Metal Electron Configurations

Special Cases and Irregularities

Transition metals (d-block) and inner transition metals (f-block) may have irregular electron configurations due to sublevel energy differences.

• 4s sublevel fills before 3d due to lower energy.

• Some transition metals have experimentally determined configurations (e.g., Cr, Cu).

Element	Expected Configuration	Experimental Configuration
Cr	[Ar] 4s2 3d4	[Ar] 4s1 3d5
Cu	[Ar] 4s2 3d9	[Ar] 4s1 3d10

9.6 Periodic Trends in the Size of Atoms

Atomic Radius

Atomic radius is the average distance from the nucleus to the outermost electron. It can be measured as van der Waals radius (nonbonding) or covalent radius (bonding).

• Atomic radius decreases across a period (left to right): Effective nuclear charge increases, pulling electrons closer.

• Atomic radius increases down a group: Valence shell is farther from nucleus.

Effective nuclear charge ($Z_{eff}$): Net positive charge attracting a particular electron.

9.7 Periodic Trends in Ionic Radius

Ionic Radius

Ionic radius refers to the size of an ion. Trends are similar to atomic radius but depend on ion charge and electron configuration.

• Ion size increases down a group: Higher valence shell, larger ion.

• Cations are smaller than neutral atoms; anions are larger than neutral atoms.

• Larger positive charge: Smaller cation.

• Larger negative charge: Larger anion.

• Isoelectronic species: Ions with the same electron configuration.

Ionization Energy (IE)

Definition and Trends

Ionization energy is the minimum energy required to remove an electron from an atom or ion in the gas phase. It is an endothermic process.

• First ionization energy ($IE_1$): Energy to remove the first electron from a neutral atom.

• Second ionization energy ($IE_2$): Energy to remove an electron from a 1+ ion.

• IE decreases down a group (valence electrons farther from nucleus).

• IE increases across a period (effective nuclear charge increases).

$M(g) + IE_1 \rightarrow M^+(g) + e^-$

$M^+(g) + IE_2 \rightarrow M^{2+}(g) + e^-$

General Trends in First Ionization Energy

• Larger $Z_{eff}$: Higher ionization energy.

• Greater distance from nucleus: Lower ionization energy.

• Exceptions: Between groups 2A/3A and 5A/6A due to half-filled p orbitals.

9.8 Electron Affinity (EA)

Definition and Trends

Electron affinity is the energy change when a neutral atom in the gas phase gains an electron. It is usually exothermic (negative value), but can be endothermic for some elements.

• EA generally increases (becomes more negative) across a period.

• EA decreases down a group (less negative).

• Highest EA: Halogens.

• Alkali earth metals and noble gases: EA is positive (endothermic).

$M(g) + e^- \rightarrow M^-(g) + EA$

Group	Electron Affinity (kJ/mol)
1A (Na)	-53
7A (Cl)	-349
8A (Ne)	+0

Properties of Metals and Nonmetals

Classification and Trends

Metals	Nonmetals
Malleable and ductile	Brittle in solid state
Shiny, reflect light	Dull, nonreflective
Conduct heat and electricity	Insulators
Form cations in solution	Form anions and polyatomic anions
Lose electrons (oxidized)	Gain electrons (reduced)

• Metallic character: Decreases left to right across a period, increases down a group.

Summary of Periodic Trends

• Atomic radius: Increases down a group, decreases across a period.

• Ionization energy: Decreases down a group, increases across a period.

• Electron affinity: Generally increases (more negative) across a period.

• Metallic character: Increases down a group, decreases across a period.

Practice and Application

• Valence electrons: Determine by group number for main group elements.

• Ion charge prediction: Based on electron configuration and group position.

• Periodic trends: Use to predict element properties and reactivity.

Additional info: These notes are based on CHEM 1010 Week 4 lecture slides and cover Chapter 9.2, 9.4, 9.6, 9.7, 9.8, and 9.9 from a standard General Chemistry textbook.