1.7 Periodic Trends in Atomic Properties (1)

8/2/2019 1.7 Periodic Trends in Atomic Properties (1)

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1.7 PERIODIC TRENDS IN ATOMIC PROPERTIES

In order for one to understand the trends of the periodic table, they must understand effectivenuclear charge. Effective nuclear charge is the attractive force that a nucleus has forelectrons.

ATOMIC RADIUS – the size of an atom (the distance between the nucleus and valence electrons).

The atomic radius of an atom is usually measured in picometres.

Across a period: Atomic Radius from right to left.REASON : a proton is added in each element on the same energy level. This increases thepositive charge of the nucleus, adding to the power of the nucleus to pull electrons inward.

How about the electrons?

Down a group: Atomic Radius as you move down the group.

REASON : Each energy level is overlapped by another energy level for each period. For example,the second energy level is a greater distance from the nucleus than the first energy level.

How about the electrons?

Figure 1: Periodic trend of atomic radii. Atomic radii increases as you move towards the left anddownwards in the periodic table.



IONIC RADIUS – the size of an ion (the distance between the nucleus of the ion and the valenceelectrons)

For Metal Cations: Ionic Radius is for a metal ion than the atomicradius of the metal atomREASON : Consider what would happen if an alkali metal lost one electron. There would be more

protons than electrons. The attractive force of the nucleus is greater and is shared betweenfewer electrons.

For Non-metal Anions: Ionic Radius for a non-metal ion than the atomicradius of the non-metal atomREASON : Consider what would happen if a non-metal gained one electron. The charge of thenucleus is the same, but there is more repulsion among the electrons. The effective nuclearcharge is shared between more electrons

Figure 2: Periodic trend of ionic radii. Atomic radii increases as you move downwards in the periodictable.



IONIZATION ENERGY – the energy required to remove an electron from its outer shell

Across a period: Ionization Energy from left to right across a period.REASON: Increased number of protons increases the energy needed to pull electrons away fromthe nucleus. More energy is needed for smaller atoms than larger ones.

Down a group: Ionization Energy as you move down the group.

REASON: (1) Distance from nucleus is greater(2) Inner electrons shield the attraction of the nucleus to the valence electrons

Figure 3: A chart representing the periodic trend of ionization energy. Ionization energy increases asyou move towards the right and downwards in the periodic table.



ELECTRON AFFINITY – energy change that occurs when an electron is added to a neutral atom

ELECTRONEGATIVITY – attraction of an atom for an electron within a chemical bond

Across a period: Electron Affinity from left to right across a period, exceptfor Noble GasesREASON : Increased number of protons within the same energy level causes increased attraction

to the incoming electron

Down a group: Electron Affinity moving down a group.REASON : Larger atom means nucleus is farther from the outer shell. Weaker effective nuclearcharge attracts the incoming electron with a lower force.

Figure 4: Various diagrams showing how the size of the atom affects its abilities to attract electrons.

REACTIVITY – ability of an atom to react

Starting from the middle, reactivitymoving outwards to the ends of the periodic table (except forNoble Gases)REASON: Moving outward, atoms need to give fewer electrons toachieve an octet

On the metal side, reactivity moving down thegroup and right to leftREASON: Large metals give up electrons easier, as they are

farther away from nucleus

On the non-metal side, reactivity moving upthe group.REASON: Smaller non-metals hold the incoming electrons morestrongly

Figure 5: Reactivity of Metals



Figure 6: Summary of Periodic Trends including atomic radius, electron affinity, ionization energy andreactivity of metals (metallic character) and reactivity of non-metals (non-metallic character)

HOMEWORK: Read Pages 36-41 and do questions #1-9 on Page 41. Try Question 10.

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1.7 Periodic Trends in Atomic Properties (1)