Trends in the periodic table
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Trends in the periodic table. Atomic radius. Atomic radii trends and explanations. Atomic radius decreases across a period because each successive element has one more proton in its nucleus and one more electron is added to the same valence shell.

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Trends in the periodic table

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Trends in the periodic table

Trends in the periodic table


Atomic radius

Atomic radius


Atomic radii trends and explanations

Atomic radii trends and explanations

  • Atomic radius decreases across a period because each successive element has one more proton in its nucleus and one more electron is added to the same valence shell.

  • Therefore, each electron experiences a greater effective nuclear charge so is attracted more strongly to the nucleus resulting in a smaller atomic radius.

  • Atomic radius increases down a group because as we progress down a group the valence electrons are found in another shell much further from the nucleus.


Ionic radius

Ionic radius


Atomic vs ionic radius

Atomic vs ionic radius

  • Cations: When cations form, all the valence electrons are removed from the outer shell, so the ions have one less shell than the atom. This results in a smaller radius than the atom.

  • Anions: When anions form, electrons are added to the existing valence shell. Greater repulsion between valence electrons results in a larger radius than the atom.


Ionic radii trends and explanations

Ionic radii trends and explanations

  • Ionic radius increases down a group for the same reason atomic radius increases.

  • Ionic radius decreases across a period for the cationsfor the same reason atomic radius increases.

  • However, here is a big jump in ionic radius between cations and anions because the anions have one more shell than the cations.

  • The trend resumes for the anions.


1 st ionization energy

1st ionization energy


1 st ionization energy definition

1st ionization energy definition

  • 1st ionization energy = amount of energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms (unit is kJ/mol).

  • Endothermic process.

  • Example equation:

    Na(g) Na+(g) + e-


Trends and anomalies

Trends and anomalies

  • 1st I.E increases across period because as effective nuclear charge increases, valence electrons are held more tightly, so more energy is needed to remove an electron.

  • Anomalies exist in trend due to more stable electron configurations which require more energy to remove electrons from half full or completely full subshells.

  • 1st I.E increases down a group because valence electrons are found in a shell much further from the nucleus and are also shielded from the nucleus so are not held as tightly and require less energy to remove.


Successive ionization energies e g sodium

Successive ionization energiese.g sodium


Explanation of sodium successive i e

Explanation of sodium successive I.E

Use the graph on the previous slide to explain the successive ionization energies for sodium. In your answer you should:

  • Describe what 1st, 2nd, 3rdetcionisation energy means.

  • Describe the overall trend.

  • Explain the jumps in I.E by referring to which shells and sub-shells electrons are being removed from.


Electronegativity

Electronegativity

  • A measure of the attraction an atom has for electrons in a bond.

  • The four most electronegative elements are F, O, N and Cl.

  • These elements have a high attraction for electrons in a bond because their atomic radius is relatively small and they have a high effective nuclear charge.

  • Essentially F, O, N and Cl nuclei attract a pair of bonding electrons more strongly than the nuclei of other elements they are bonded to.

  • Their nuclei are closer to the bonding electrons and they have a higher effective nuclear charge to attract those electrons.


Electronegativity and polarity of bonds

Electronegativity and polarity of bonds

  • The difference in electronegativity helps to determine whether a bond will be polar or not.


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