1. Trends & the Periodic Table
Objective:
I will be able to:
name and describe the primary trends of the Periodic Table.

2. Trends
More than 20 chemical properties change in a predictable way based location of elements on the Periodic Table.
The major properties we will be using:
density
melting point/boiling point
atomic radius
ionization energy
electronegativity
: - anyone know where we can find these numbers?!

3. Electron Configuration Trend
The number of energy levels increases as we go down a group. Fr 7 Cs 6 Rb 5 K 4
2-8-1 Na 3
2-1 Li 2 1 H
Configuration
Element
Period
Group 1

4. Periodic Trends
Add 5 arrows to your sheet
You will be filling in your handout as we progress through the PowerPoint.
This is for a grade!

5. Periodic Trends
Increasing number of energy levels

6. Atomic Radius Trend

7. Atomic Radius Trend
Compare
Li: Group 1 Period Cs: Group 1 Period 6

8. Atomic Radius Trend Increasing Atomic Radius
Increasing number of energy levels
Increasing Atomic Radius

9. Atomic Radius Trend
As we go across a period from left to right, elements gain electrons, but the radius is getting smaller!
2-8 Ne
VIIIA or 18
2-7 F
VIIA or 17
2-6 O
VIA or 16
2-5 N
VA or 15
2-4 C
IVA or 14
2-3 B
IIIA or 13
2-2 Be
IIA or 2
2-1 Li
IA or 1
Configuration
Element
Family

10. Atomic Radius
Across a period

11. Why does this happen..
As you go from left to right, you gain more protons (the atomic number increases)
You have greater “proton pulling power”
Remember the nucleus is (+) and the electrons are (-) so they get pulled towards the nucleus
The more protons you
have, the more Proton Pulling Power

12. Proton Pulling Power
                                                                                                                                                                                                                  

13. Atomic Radius Trend Decreasing Atomic Radius
Increasing number of energy levels
Increasing Atomic Radius
Decreasing Atomic Radius

14. Proton Pulling Power and Effective nuclear charge (Zeff)
We can “measure” the Proton Pulling Power by determining the “Effective Nuclear Charge”
It is the charge actually felt by valence electrons
Electrons (e-) are attracted to the (+) nucleus.
3 factors that effect this:
The more protons in the nucleus, the greater the Zeff
The more distance between the nucleus and electrons the smaller the Zeff
The more repulsion between electrons the smaller the Zeff

15. What the inner electrons do….
They shield the charge felt by the valence electrons.

16. Proton Pulling Power
H and He are only elements whose valence electrons feel full nuclear charge (pull)
NOTHING TO SHIELD THEM

17. Atomic Radius Trend Increased Electron Shielding
Increasing number of energy levels
Increasing Atomic Radius
Decreasing Atomic Radius
Increased Electron Shielding

18. Look at all the shielding Francium's one valence electron has
Look at all the shielding Francium's one valence electron has. It barely feels the proton pull from the nucleus. No wonder it will lose it’s one electron the easiest. No wonder it’s the most reactive metal

19. Ionization Energy
ionization energy: the amount of energy required to remove an electron from an atom to form a cation
1st ionization energy: energy required to remove the most loosely held valence electron (e- farthest from nucleus)

20. ionization energy
Cs valence electron is farther away from nucleus than Li’s valence electron.
Electrostatic attraction is much weaker so it is easier to steal an electron away from Cs.
THEREFORE, Li has a higher Ionization energy then Cs.

21. Increased Ionization Energy (harder to remove an electron)
Decreasing Atomic Radius
Increased Electron Shielding
Increasing number of energy levels
Increasing Atomic Radius
Decreased Ionization Energy (easier to remove an electron)

22. Electronegativity
electronegativity: the ability of an atom to attract electrons when the atom is part of a chemical compound
Linus Pauling won a Nobel Prize in Chemistry for his work on chemical bonds and was the first to describe electronegativity. Electronegativity units are named after him.
Electronegativity Unit = Pauling

23. Electronegativity
Metals at the far left of the Periodic Table have low values.
Nonmetals at the far right have high values.
Noble gases tend not to form bonds, so they don’t have electronegativity values
Most and least electronegative elements:
Fluorine = 4.0 Paulings
Francium = 0.7 Paulings
Transition metals electronegative values are irregular.

24. Increased Electronegativity
Increased Ionization Energy (harder to remove an electron)
Decreasing Atomic Radius
Increased Electron Shielding
Decreased Electronegativity
Decreased Ionization Energy (easier to remove an electron)
Increasing number of energy levels
Increasing Atomic Radius

25. Electronegativity

26. Reactivity of Metals and Nonmetals
The reactivity of metals is judged by how easily they give up electrons (they are e- losers).
The reactivity of non-metals is judged by how easily gain electrons (they are e- gainers).

27. Noble Gases are Nonreactive Increasing number of energy levels
Increasing Atomic Radius
Decreasing Atomic Radius
Increased Electron Shielding
Decreased Ionization Energy (easier to remove an electron)
Increased Ionization Energy (harder to remove an electron)
Decreased Electronegativity
Increased Electronegativity
Most Reactive Metal
Most Metallic Metal
More metallic
Most Reactive Nonmetal
Metal Reactivity
Nonmetal Reactivity

28. Periodic Table Group Names
Group 1 = Alkaline Metals
Group 2 = Alkaline Earth Metals
Group 13 = Boron family
Group 14 = Carbon family
Group 15 = Nitrogen family
Group 16 = Oxygen family
Group 17 = Halogens
Group 18 = Noble Gases

29. Ions
ion: an atom or group of atoms that have a positive or negative charge
Positive and negative ions form when electrons are transferred between atoms.
Octet Rule: atoms tend to gain, lose or share electrons so as to have 8 electrons in the valence shell

30. How do you know if an atom gains or loses electrons?
Think about Lewis Dot Structures of ions.
Lewis Dot Structures are Symbols of atoms with dots to represent the valence-shell electrons.
Transition Elements

31. How do you know if an atom gains or loses electrons?
Atoms form ions to get a valence # of 8 (or 2 for H).
Metals tend to have 1, 2, or 3 valence electrons.
It’s easier to lose electrons.
Nonmetals tend to have 5, 6, or 7 valence electrons.
It’s easier to add electrons.
Potassium will give its e- to chlorine
Noble gases already have 8 (or 2 for He) so they don’t form ions very easily.
1e e-

32. Positive ions (cations)
Cations are formed by loss of electrons.
Cations are always smaller than parent atom because they have one less energy level. 2e 8e 8e 8e 8e 2e 2e Ca Ca
Ca+2

33. Negative ions (anions)
Anions are formed by gaining electrons.
Anions are always larger than the parent atom because the proton pulling power is less due to more electrons being pulled.

34. Allotropes
allotrope: one or more different molecular forms of an element in the same physical state.
Allotropes have different structures and properties.
O2 and O3 - both gas phase
O2 (oxygen) - necessary for life
O3 (ozone) - toxic to life
Graphite, coal, diamond, fullerene:
All are carbon in solid form.

35. Melting Point and Thermal conductivity
melting point: the temperature at which a substance changes from a solid to a liquid
At this temperature the vibrations of the atoms are strong enough to break the bonds holding the atoms together.
Most metals generally possess a high melting point.
Most non-metals possess low melting points.
Exceptions:
The non-metal carbon possesses the highest boiling point of all the elements. (MP=6,381 0F, BP=7,281 0F)
The semi-metal boron also possesses a high melting point. (MP=3,769 0F, BP=7,101 0F)
The metal gallium melts slightly above room temperature. (MP=86 0F, BP=4,352 0F)

36. Melting Point and Thermal conductivity
thermal conductivity: a property of a material that determines the rate at which it can transfer heat
Thermal conductivity is:
lowest at the bottom corners of the periodic table and
highest in the upper center of the periodic table.

37. Websites Crash Course #22 – Types of Chemical Bonds
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Crash Course #23 – Polar and Non-Polar Molecules
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of-Solids-Liquids-and-Gases-in-general-use-today