Periodic Properties of Elements

• The Modern Periodic Table arranges elements by increasing atomic number.
• Element properties are periodic functions of their atomic numbers.
• Periodicity is the recurrence of properties at regular intervals.
• Definite intervals occur after differences of 2, 8, 18, or 32 in atomic numbers in the modern periodic table.
• Properties appearing at regular intervals are called 'periodic properties'.
• Periodicity is due to elements having similar valence shell electronic configurations after definite atomic number intervals.
• Element properties depend on electron number and arrangement in shells.
• Property increase or decrease in a period/subgroup results from gradual changes in electronic configuration.

Periodic Properties

• Atomic radii
• Ionization potential
• Electron affinity
• Electronegativity
• Non-metallic and metallic character.
• Density
• Melting and boiling point
• Nature of oxides, oxy-acids, hydrides

Atomic Radius

• The distance between the nucleus and outermost shell.

Ionization Potential (I.P.)

• The energy needed to remove an electron from an isolated gaseous atom's outermost shell.

Electron Affinity (E.A.)

• Energy released when a gaseous atom accepts an electron to form an anion.

Electronegativity (E.N.)

• The tendency of an atom to attract electrons when combined.

Non-metallic & Metallic Character

• Non-metal: gains one or more electrons
• Metal: loses one or more electrons

Atomic Size

• Atomic size refers to the distance between the nucleus and outermost shell of an atom

Factors Affecting Atomic Size

• Number of Shells:
  • More shells increase atomic size.
  • Atomic size increases due to the distance between the nucleus and the outermost shell.
• Nuclear Charge:
  • Greater nuclear charge decreases atomic size.
  • Increased nuclear charge attracts electrons with greater force.
  • Nuclear charge is the positive charge on the nucleus, equivalent to the atomic number.

Atomic Size Trends Across a Period (Left to Right)

• Atomic size generally decreases across a period from left to right
• Number of Shells:
  • Stays the same as you move across the periodic table
• Nuclear Charge:
  • Increases as you move across the periodic table
• In period 2, Lithium has the largest, and Fluorine has the smallest atomic radius.
• Neon has a larger atomic radius due to its full outer shell leading to repulsion.
• The nuclear pull is not seen over valence shell electrons

Atomic Size Trends Down a Group

• Number of Shells:
  • Increases down a group.
  • New shells are added with increasing atomic number which increases atomic size.
• Nuclear Charge:
  • The effect of the nuclear charge increasing is dominated by number of shells increasing
• Atomic Size:
  • Increases down a group

Ionization Potential

• The amount of energy required to remove a loosely bound electron from the outer shell of an isolated gaseous atom.
• Removing the 1st electron is called the 1st ionization potential, and removing a 2nd electron is the 2nd ionization potential.

Factors Affecting Ionization Potential

• Atomic Size:
  • Larger atomic size leads to decreased ionization potential.
  • The nuclear attraction on outer electrons decreases in larger atoms.
• Nuclear Charge:
  • Greater nuclear charge increases ionization potential.
  • The nuclear attraction on the outer electrons increases and are more firmly held.

Ionization Potential Trends Across a Period

• Ionization potential generally increases across a period (left to right).
• Atomic radii decreases across a period
• Nuclear charge increases across a period
• Helium has the highest ionization potential and caesium has the lowest
• Neon in period 2 has maximum I.P. due to its completely filled outer orbit
• Metals tend to have low ionization potentials compared to non-metals

Ionization Potential Trends Down a Group

• Ionization potential decreases down a group.
• Atomic radii increases down a group
• Nuclear charge increases down a group, but is dominated by atomic radii increase

Electron Affinity

• The amount of energy released when a gaseous atom accepts an electron to form an anion.
• Electron affinity is represented with a negative sign.

Factors Affecting Electron Affinity

• Atomic Size:
  • Larger atomic size decreases electron affinity.
  • Smaller atoms accept electrons more readily because the nucleus has greater attraction.
• Nuclear Charge:
  • Greater nuclear charge increases electron affinity.
  • Increased nuclear charge makes it easier to gain electrons

Electron Affinity Trends Across a Period

• Electron affinity generally increases across a period (left to right).
• Atomic radii decreases across a periods
• Nuclear charge increases across a period
• Neon has an electron affinity of zero since its stable configuration makes it difficult to accept electrons.
• Halogens have the highest electron affinity, and alkali metals have the lowest electron affinity
• The more electronegative element has a higher E.A.

Electron Affinity Trends Down a Group

• Electron affinity decreases down a group
• Atomic radii increases down a group
• Nuclear charge increases down a group

Electronegativity

• Electronegativity refers to the tendency of an atom to attract electrons to itself when combined in a compound.
• Ionic bond formation occurs between atoms with large electronegativity differences.
• Covalent bond formation occurs between atoms with similar electronegativities

Factors Affecting Electronegativity

• Atomic Size:
  • Larger atomic radii causes electronegativity to decrease
• Nuclear Charge:
  • Greater nuclear charge causes electronegativity to increase
  • Electronegativity is affected by similar factors as electron affinity.

Electronegativity Trends Across a Period

• Electronegativity increases across a period from left to right
• Atomic radii decreases across a period
• Nuclear charge increases across a period.
• Elements are normally non-metallic and non-metals are often electronegative
• Fluorine is the most electronegative element, whilst caesium is the least electronegative
• Noble gases are complete octet and don't attract electrons.

Electronegativity Trends Down a Group

• Electronegativity decreases as you descend a group.
• Atomic radii increases down a group
• Nuclear charge increases down a group

Metallic & Non-Metallic Character

• Metallic Character:
  • Involves the tendency of an atom to lose electron(s)
  • Metals lose one or more electrons when supplied with energy.
• Non-Metallic Character
  • Involves the tendency of an atom to gain electron(s).
  • Non-metals gain one or more electrons when supplied with energy.

Factors Influencing Metallic & Non-Metallic Character

• Atomic Radii:

  • Increasing atomic radii increases metallic character and decrease non-metallic character.

• Ionization Potential:

  • Decreasing ionization potential increases metallic character and decrease non-metallic character.
  • Metallic atoms have Low-ionisation potential value & tend to Lose-electrons.
  • Non-metallic atoms have High-ionisationpotential value & tend to Gain-electrons.

• Metals are good reducing agents, and non-metals.

• Metals have greater tendency to lose electrons, and greater reactivity.

• Non-metals have greater tendency to gain electrons, and greater reactivity.

Trends in Character Across a Period

• Moving from left to right across a period:

  • Metallic character generally decreases
  • Non-metallic character increases

• Atomic Radii:

  • Decreases across a period

• Ionization Potential:

  • Increases across a period.

Trends in Character Down a Group

• Top of the group are non-metal

• Elements at the bottom of the group are

  • Having largest size and lowest I.P.
  • Loosely held electrons will form ions from metals most readily & thus are more reactive.

• Atomic Radii:

  • Increases down a group

• Ionization Potential:

  • It Decreases down a group