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Section 5.3 – Electron Configuration and Periodic Properties

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Presentation on theme: "Section 5.3 – Electron Configuration and Periodic Properties"— Presentation transcript:

1 Section 5.3 – Electron Configuration and Periodic Properties
Honors chemistry

2 Atomic Radius Determined by the distance from the nucleus to the edge of the outer orbital. Edge of outer orbital not well defined Use identical bonded atoms – ½ of the distance between the nuclei

3 Trends in Atomic Radii

4 Trends in Atomic Radii Decrease across the period
Due to increasing positive charge of the nucleus Increase as you go down the group Exception Ga to Al – Ga smaller due to increased nuclear charge (first addition of d electrons)

5 Problems Which of these elements; Li, Rb, K or Na has the smallest radius? Largest? Which of these elements; O, Se, S and Po has the smallest atomic radius? Largest?

6 Ionization Energy Atom + energy → Atom+ + e-
First electron removed – First Ionization Energy (IE1) Second electron removed – Second Ionization Energy (IE2) etc. Group 1 – lowest ionization energy Group 18 - highest ionization energy Ionization energies increase across the period due to increased nuclear charge. Ionization energies decrease down a group due to further distance from nucleus and electron shielding

7 Ionization Energies

8 Ionization Energies

9 Successive Ionization Energies

10 Why? Each successive electron feels a stronger nuclear attraction
This information lead to the understanding of the stability of the noble gas configuration+

11 Practice Choose the element with the higher IE1: Ca or Ba Ca or Br
Ca or K Ca or Mg

12 Electron Affinity Atom + e- → Atom- + (- energy) IMPORTANT!!!!!
Negative energy means energy lost by system Positive energy means energy gained by system Sign indicates direction not numerical value!!!!

13 Electron Affinity Values

14 Electron Affinity Trends
Generally become larger (look at as absolute value) as you move across the period. Exception – Group 15 due to half filled p orbitals Generally become smaller as you move down a group due to: Greater Nuclear Attraction Greater Atomic Radius

15 Second Electron Affinities
Very difficult to add an electron to an anion (negative ion) Second Electron Affinities are all positive

16 Ionic Radii Cation (positive ion) Anion
Smaller atomic radius than atom Due to: electrons being removed increased effective nuclear charge Anion Larger atomic radius than atom Due to electrons being added decreased effective nuclear charge Greater repulsion of electrons

17 Ionic Radii

18 Valence Electrons Available to be lost, gained or shared when compounds are formed. In outer main energy levels For Main Group Elements – s and p orbitals

19 Bonded Atoms Very rarely are electrons shared equally
Usually attracted more to one atom This will effect the chemical properties of the compound!!! Measure of attraction – called electronegativity Based on a 4.0 scale – F = 4.0. Developed by Linus Pauling

20 Electronegativity

21 Electronegativity Trends
Increase across a period. Tend to decrease or stay the same down a group. If a noble gas does not form compounds – it does not have an electronegativity If a noble gas does form compounds – it will have a high electronegativity

22 Summary of Trends

23 Summary of Trends

24 D-Block These elements tend to vary less and with less regularity than Main Group Elements. Still electrons in d orbitals are often responsible for characteristics of elements in the d-block Atomic radius tends to decrease across the block Ionization energies generally increase across both the d and f-blocks

25 D and F Blocks Tend to lose electrons from outer shell!!!!
That means the valence electrons come from the ns shell not the (n-1)d shell Generally these elements from 2+ ions. Electronegativities D-block - between 1.1 and 2.54 (Only groups 1 and 2 are lower) F-block – between 1.1 and 1.5

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