Download presentation

Presentation is loading. Please wait.

1
Periodic Trends Notes

2
Reason for lesson Today we are going add to our knowledge of the periodic table and introduce the important ideas of the periodic law, atomic radius, ionic radius electronegativity, ionization energy, electron affinity and shielding effects.

3
Periodic Trends The number of protons and electrons affects the properties of an atom. The periodic table is organized according to groups (up and down) and periods (right to left).

4
Periodic Trends 1. Atomic Mass:
increases as you move across the periodic table (from left to right) Increases as you move down a group

5
Periodic Trends 2. Atomic Radius: the size of the atom. If you compare two atoms that are directly above and below each other on the table, the one underneath will be larger because it has another shell added.

6
Atomic radius (cont.) Lithium Rubidium 3+ 37+

7
However, if you compare two atoms that are side-by-side on the table, the one on the left will be larger because when you add more electrons and protons, the electric force that pulls them together will increase, making the radius smaller. Atomic Radius (cont.)

8
Atomic Radius (cont.) Lithium Fluorine 3+ 9+

9
Atomic Radius (cont.)

10
Ionic radius

11
Ionic radius General trend increase moving down a group
Decreases moving across a period Na+ Mg 2+ Al3+ P3- S2- Cl-

12
Periodic Trends 3. Ionization Energy–the amount of energy that is required to remove an electron from a gaseous atom. If an atom will easily lose an electron (like the Alkali Metals), the ionization energy will be low. If it will not lose electrons easily (noble gases, Halogens), the ionization energy will be high.

13
Ionization Energy (cont.)
Ionization energy increases going left to right on the table, and decreases going down. Larger atoms lose electrons more easily because the electric force is less because the distance from the nucleus is greater.

14
Ionization Energy (cont.)
Lithium Rubidium 3+ 37+

15
Ionization Energy (cont.)
Across a period the atoms get smaller because the larger electric force in the nucleus pulls the valence shell in tighter allowing for a better hold on the outer electrons and a greater ionization energy.

16
Ionization Energy (cont.)
Lithium Fluorine 3+ 9+

17
Ionization Energy (cont.)

18
Multiple Ionization Energies
It is possible to remove all the electrons of an atom. To remove the second electron it will take more energy then to remove the first. To remove the third will take more energy then to remove the second. The table on the next slide gives first, second , third, etc ionization energies for different atoms.

19
Multiple Ionization Energies (cont.)
Element 1st 2nd 3rd 4th 5th Lithium 0.5 7.3 11.8 Beryllium 0.9 1.8 14.8 21.0 Boron 0.8 2.4 3.7 25.0 32.8

20
Review of Last Lesson Chapter 7 Ionization energy

21
Periodic Trends 4. Electron Affinity–the energy change when an electron is added to a gaseous atom. This is closely related to ionization energy, and increases going left to right, and decreases going down.

22
Electron Affinity (cont.)
Increases because as you move across the row, atoms become increasingly stable when they receive an electron. In chemical terms, energy is always released when stability is achieved. That is because when something is stable it is at a lower energy level.

23
Electron Affinity (cont.)
Lithium Fluorine 3+ 9+

24
Electron Affinity (cont.)
Decreases going down a group due to increasing number of electrons and shells therefore the effect is not as great when an atom gains one electron

25
Electron Affinity (cont.)
Lithium Rubidium 3+ 37+

26
Electron Affinity (cont.)

27
Periodic Trends 5. Electronegativity– Ability of an atom to attract electrons when in a molecule.

28
Electronegativity (cont.)
This increases going left to right, and decreases going down (with the exception of noble gases, which don’t make molecules with other atoms).

29
Electronegativity (cont.)
Linus Pauling assigned a value of 4.0 to the most electronegative element fluorine and all other elements were measured relative to F.

30
Electronegativity (cont.)
Why is Fluorine the most electronegative element? Which should be the least electronegative?

31
Electronegativity (cont.)
How does electronegativity effect molecules? Let’s use water as an example. Electronegativity is defined as the ability of an atom to attract an electron in a molecule or bond.

32
Electronegativity (cont.)
Water naturally has a bend in its structure. O H H If hydrogen and oxygen had the same electronegativity we would expect the electrons to be shared in the middle.

33
Electronegativity (cont.)
But they have different electronegativities. Oxygen has a larger electronegativity then hydrogen O H H Electrons are more attracted to the oxygen and spend more time closer to the oxygen.

34
Electronegativity (cont.)
This gives the oxygen a partial negative charge and the hydrogen a partial positive charge. d- d+ O d+ H H We call these forces “dipoles”. Depending on the shape of the molecule, some dipoles cancel out. If they do not cancel out, we call the bonds/molecule “polar”.*

35
Review of Electonegativity
Chapter 7 Electronegativity

36
Periodic Trends Reviewed

37
Periodic Trends. Up and to right, electrons are held more strongly.

38
Melting Point Trend

39
Melting Point Trend

40
Melting Point Trend The temperatures of changes of state are a measure of the strength of the intermolecular forces holding the molecules, ions or atoms together. The melting point trend across a period demonstrates that the intermolecular forces get stronger and peak around group 14 and then fall.

44
Why are the halogens and the alkali metals so reactive?
Starting Question Why are the halogens and the alkali metals so reactive?

45
Starting question Some atoms form crystals, Sodium Chloride is an example. What is another example? Iron oxide crystals will react with aluminum to form aluminum oxide crystals and free iron. The reaction will release so much heat it can melt steel. It is used in wielding. Fe2O3(s) + 2 Al(s) f. Al2O3(s) + 2 Fe(l). Why does the reaction above produce so much heat? The reason has to do with the radius of the aluminum and iron ions and the oxygen ion.

46
Answer to starting question
The radius of the aluminum ion is smaller than the radius of the iron ion, so the crystal fits together better. In this way, the strain in the crystal is released. The difference in energy, (called lattice energy) is released as heat.

Similar presentations

© 2023 SlidePlayer.com Inc.
All rights reserved.

You are watching: Periodic Trends Notes.. Info created by GBee English Center selection and synthesis along with other related topics.