1. Electron Configuration and Atomic Properties
Section 11.4
Electron Configuration and Atomic Properties
PERIOD 6 Group #2

2. Background
- Electron; covers a broad area it also behaves like cloud in an
atom
to understand how the principal energy levels fill with electrons in
atoms beyond hydrogen.
to learn about valence electrons and core electrons.
to learn about the electron configurations of atoms with Z < 18.
to understand the general trend in properties in the periodic table.

3. Electron Configuration
The electron arrangement in atoms from Z = 1 to Z = 18 can be described by placing electrons in principal energy level
Principal energy level – energy denoted by principal quantum number n. ex. – n = 1 , n = 2, n = 3….
The 1st element in the each period of the Periodic Table represents a new principal energy level.
Ex. – hydrogen has lone electron in 1s orbital / so represents 1s^1 because it is in electron arrangement
Electron arrangement – can be called electron configuration in an atom

6. Question #1
For the first 18 elements, what is the sublevel order and the elements?

7. Answer #1
1s 2 s 2p 3s 3p are the sublevels for the first 18 elements.

8. Orbital Diagram Also known as the box diagram
Gives more detailed information about electron configuration
Arrows represent an electron spinning in a particular direction.
For each element’s orbital – takes one electron before any orbital in each sub level can receive another electron.
Pauli Exclusion Principle –

9. Question #2
What is the electron configuration and orbital diagram for silicon?

10. Example of Orbital Diagram Answer #2
Silicon – 1s^2 2s^ ^ s^ p^2

11. Metal & Non-metal & Metalloids Chart
All URL on “Reference” page
Metals
Non-metals
Metalloids
Gallium crystals
from left to right He, Ne, Ar, Kr, Xe
Boron
Iron (Fe)
from left to right; F, Cl, Br, l
Silicon
Copper(Cu)
Sulfur
Germanium

12. Metals & Nonmetals & Metalloids
Metals and nonmetals have relatively large ionization energies.
Metals gain electrons
Lower left region of the periodic table have the lowest ionization energies (are the most chemically active metals)
Upper right region of the periodic table have the highest ionization energies (are the most chemically active nonmetals)
Metalloids – found along the stair-steps line between the metals and nonmetals.
Metalloids have properties of both metals and nonmetals.

13. Question #3
Name the physical properties of the metals.

14. Answer #3
A lustrous appearance, the ability to change shape without breaking, and excellent conductivity of heat and electricity.

15. Atomic Size Increases down a group
Decreases to the right of the period.
The average distance of the electrons from the
nucleases increases.
Atoms get bigger as electrons are added to larger principal energy levels.
The sizes of atoms vary
The atoms in a particular period all have their outermost electrons in a given principal level

16. Question #4
Which element has the smallest atomic size?

17. Answer #4
The element with the smallest atomic size is helium.

18. Ionization Energy

19. Ionization Energy
The energy required to remove an electron from and individual atom in the gas phase.
Increases up a group and to the right of a period.
Metals have relatively low ionization energies.
Ionization energies tend to decrease in going from the top to the bottom of a group.

20. Question #5
What is the definition of ionization energy and how does it related to the elements?

21. Answer #5
It is the energy to remove an electron from an individual atom in the gases phase of an element and the energies tend to decrease going down a group.

22. Quiz
What is the electron configuration for Cerium (Ce; atomic #58)? Include a box diagram.
What is the Pauli Exclusion Principle and how does it relate to the orbital diagram?
Out of the metals and non-metals, which has relatively large ionization energy?
In which pattern so metals tend to lose electrons? Explain.
Define atomic mass and atomic size.

23. Answer key
1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6 5s^2 4d^10 5p^6 6s^2 4f^1
2 electrons in the same level cannot be in the same quantum state so in the orbital diagram, the arrows cannot be in the same direction, and each orbital can only hold up to 2 electrons.
Nonmetals
Metals tent to lose electrons going down a group because as we go down a group, the electron being removed resides on average farther and farther from the nucleus.
Atomic mass is the mass of an atom and atomic size is the size of the atom.

24. Reference https://teach.lanecc.edu/gaudias/scheme.gif