3.4 Periodic table and trends in atomic properties

The periodic table organizes elements based on atomic number and electron configuration. It reveals patterns in atomic properties, helping predict chemical behavior. Elements in the same group share similar traits due to their valence electrons.

Atomic properties like radius, ionization energy, and electron affinity follow trends across periods and groups. These trends result from the interplay between effective nuclear charge and electron shielding, influencing elements' reactivity and bonding tendencies.

Organization of the Periodic Table

Arrangement and Atomic Number

Top images from around the web for Arrangement and Atomic Number

• 

  Elements and The Periodic Table | Introductory Chemistry – Lecture & Lab View original

  Is this image relevant?

• 

  File:Periodic table of elements showing electron shells.png - Wikipedia View original

  Is this image relevant?

• 

  File:Periodic Table and electron configurations.svg - Wikibooks, open books for an open world View original

  Is this image relevant?

• 

  Elements and The Periodic Table | Introductory Chemistry – Lecture & Lab View original

  Is this image relevant?

• 

  File:Periodic table of elements showing electron shells.png - Wikipedia View original

  Is this image relevant?

1 of 3

Top images from around the web for Arrangement and Atomic Number

• 

  Elements and The Periodic Table | Introductory Chemistry – Lecture & Lab View original

  Is this image relevant?

• 

  File:Periodic table of elements showing electron shells.png - Wikipedia View original

  Is this image relevant?

• 

  File:Periodic Table and electron configurations.svg - Wikibooks, open books for an open world View original

  Is this image relevant?

• 

  Elements and The Periodic Table | Introductory Chemistry – Lecture & Lab View original

  Is this image relevant?

• 

  File:Periodic table of elements showing electron shells.png - Wikipedia View original

  Is this image relevant?

1 of 3

• The periodic table is arranged in order of increasing atomic number
  • Atomic number corresponds to the number of protons in the nucleus of an atom
• Elements are organized into periods (rows) and groups (columns) based on their electron configurations and chemical properties

Valence Electrons and Chemical Behavior

• Elements in the same group have similar electron configurations in their outermost energy level (valence electrons)
  • Exhibit similar chemical behavior due to similar valence electron configurations
• As you move from left to right across a period, the number of valence electrons increases
  • Elements become more nonmetallic in character (carbon, nitrogen, oxygen)

Periodic Table Blocks

• The periodic table is divided into four blocks (s, p, d, and f) based on the subshell in which the last electron is added to the electron configuration
  • s-block elements (alkali metals and alkaline earth metals)
  • p-block elements (main group elements)
  • d-block elements (transition metals)
  • f-block elements (lanthanides and actinides)

Periodic Trends in Atomic Properties

Atomic Radius

• Atomic radius is the distance from the nucleus to the outermost stable electron orbital in an atom
• Generally decreases from left to right across a period (lithium to neon)
  • Increasing effective nuclear charge pulls electrons closer to the nucleus
• Increases from top to bottom within a group (lithium to cesium)
  • Additional electron shells increase distance between nucleus and outermost electrons

Ionization Energy

• Ionization energy is the minimum energy required to remove an electron from a neutral atom in the gaseous state
• Generally increases from left to right across a period (sodium to argon)
  • Stronger attraction between electrons and nucleus makes removal more difficult
• Decreases from top to bottom within a group (fluorine to iodine)
  • Increased distance between outermost electrons and nucleus reduces attraction

Electron Affinity

• Electron affinity is the energy change that occurs when an electron is added to a neutral atom in the gaseous state
• Generally becomes more negative from left to right across a period (beryllium to chlorine)
  • Increased attraction for additional electrons due to higher effective nuclear charge
• Becomes less negative from top to bottom within a group (fluorine to iodine)
  • Larger atomic size reduces attraction for additional electrons

Interplay of Effective Nuclear Charge and Shielding Effect

• Periodic trends result from the interplay between increasing number of protons (effective nuclear charge) and shielding effect of inner electrons
  • Effective nuclear charge increases across a period and down a group
  • Shielding effect of inner electrons increases down a group but remains relatively constant across a period

Predicting Element Properties

Similar Properties within Groups

• Elements in the same group tend to have similar chemical properties due to their similar valence electron configurations
  • Alkali metals (Group 1) are highly reactive and form +1 cations
  • Halogens (Group 17) are highly reactive and form -1 anions

Metallic and Nonmetallic Character

• Metallic character decreases from left to right across a period and increases from top to bottom within a group
  • Metals (left side of periodic table) are good conductors of heat and electricity, malleable, and ductile
• Nonmetallic character increases from left to right across a period and decreases from top to bottom within a group
  • Nonmetals (right side of periodic table) are poor conductors, brittle, and have high electronegativities

Reactivity Trends

• The reactivity of metals generally increases from top to bottom within a group (lithium to cesium)
  • Lower ionization energies facilitate loss of electrons
• The reactivity of nonmetals generally decreases from top to bottom within a group (fluorine to iodine)
  • Decreased electron affinity reduces tendency to gain electrons

Gradual Changes within Periods

• Elements in the same period tend to have a gradual change in properties as the atomic number increases
  • Increasing ionization energy (lithium to neon)
  • Increasing electronegativity (sodium to argon)

Atomic Properties and Reactivity

Influence on Chemical Bonding

• The trends in atomic radius, ionization energy, and electron affinity influence the formation of chemical bonds and reactivity of elements
  • Smaller atomic radii and higher ionization energies favor covalent bonding (carbon)
  • Larger atomic radii and lower ionization energies favor ionic bonding (sodium chloride)

Reactivity of Metals and Nonmetals

• Elements with low ionization energies and low electronegativities (alkali metals) readily lose electrons to form cations
  • Highly reactive due to ease of electron loss (sodium reacts vigorously with water)
• Elements with high ionization energies and high electronegativities (halogens) readily gain electrons to form anions
  • Highly reactive due to strong attraction for electrons (chlorine is a strong oxidizing agent)

Transition Metals

• Elements with intermediate ionization energies and electronegativities (transition metals) can form multiple oxidation states
  • Exhibit a wide range of reactivities depending on the specific element and oxidation state
  • Iron can form both +2 and +3 oxidation states, each with different chemical properties

Predicting Bonding and Stability

• Periodic trends in atomic properties can be used to predict the type of bonding (ionic, covalent, or metallic) that occurs between elements
  • Larger differences in electronegativity favor ionic bonding (sodium chloride)
  • Similar electronegativities favor covalent bonding (water, carbon dioxide)
• Trends also influence the stability of the resulting compounds
  • Compounds with large differences in electronegativity tend to be more stable (potassium fluoride)
  • Compounds with similar electronegativities may be less stable or reactive (hydrogen peroxide)