1. Periodic Trends

2. Periodic Trends
Periodic trends are patterns that are seen throughout the Periodic Table due to the way the elements are arranged.
Trends occur both in groups and in periods.

3. # 1 - Atomic Radius Size of the atom
Determined by the number of energy levels occupied by electrons
We measure this by looking at distance between nuclei of bonded atoms, and then dividing by 2
This is because atoms don’t have clearly defined outer edges.

4. Atomic radius continued
Trends:
Groups trends – atomic size increases as you move down a group
Number of energy levels containing electrons is increasing
Period trends – atomic size decreases as you move from left to right across a period (within the same energy level)
Increased attraction between the nucleus and the electrons (pulls electrons in tighter).

5. Atomic Radius

6. The Shielding Effect Attraction between electrons & nucleus
As more electrons are added (energy levels), less attraction to nucleus

7. The Shielding Effect NUCLEUS (POSITIVELY CHARGED)
Electrons in this energy level are EASIEST to move
The more energy levels there are, the weaker the attraction becomes between those electrons and the nucleus. You can think of those outer electrons as being “shielded” from the attraction of the nucleus by the energy levels in between them.
1st Energy Level: Electrons in this level are closest to the nucleus. Therefore, they experience a STRONG attraction and are hardest to move.
2nd Energy Level: Electrons in this level are a little further from the nucleus. Therefore, their attraction is not as strong as the 1st level and are easier to move.
The electrons in these energy levels “absorb” most of the attraction from the nucleus, so outer electrons are “shielded”

8. #2 - Ionization Energy
Remember: Ions are atoms with charges (more or less electrons than protons.)
Ionization Energy is the energy needed to remove an electron thus creating an ion (charged particle)
Energy is required in this process to overcome the attraction of the nucleus for the electrons
The lower the energy, the easier it is to remove an electron
Group trend - This energy decreases as you move down a group
The valence electrons are further away from the nucleus
Period trend - This energy increases across a period
The atoms are getting slightly smaller. Therefore, electrons get closer to nucleus and thus the nucleus has a greater attraction on the electrons

9. Ionization Energy

10. # 3 - Electron Affinity
Energy change that results when an electron is added to an atom or ion
In other words, it’s the ability of an atom to attract and hold an extra electron
We measure Electron Affinity by the change in energy that occurs when an electron is added to an atom or ion

11. Electron affinity Negative e- affinity: Positive e- affinity:
Means an atom releases energy when it gains an electron
Positive e- affinity:
Means energy must be added to the atom for the electron to be added
The more negative the electron affinity, the easier it is to attract an electron.

12. Electron affinity Trends: Group – decreases as you move down a group
Period – increases as you move from left to right across the periodic table
Exceptions: group 18
Why?

13. #4 - Electronegativity
The tendency of an atom to attract e- to itself when it is chemically combined with another element
Electronegativity values range from 0.7 to 4.0
The larger the electronegativity value, the easier it is to attract an electron.

14. Electronegativity continued
Group trend - Values decrease going down a group
As you go down a group, more energy levels are added. Therefore, the nucleus’ ability to attract electrons from other atoms weakens.
Period trend - Values increase across a period
The positive charge of the nucleus gets bigger due to increase in atomic number. Therefore, it has a stronger “pull” on electrons.
Exception – group 18 is not included
Why?

15. Electronegativity