1. Periodic Trends

2. Which alkali metal has an atomic radius of 238 pm?
Based on the alkali metals above, how does atomic size change within a group (Li – Fr)?
What seems to be happening within each period?
What is happening when we go from period to period?
potassium
Increases
Radius decreases across the period
Radius increases from the previous period

3. Graphing Periodic Trends Activity
Use the information in the tables to complete the graph.
Using colored pencil or pen, list under the symbol (in this order!)
the Atomic Radius,
First Ionization Energy, and
Electronegativity.
Use a different color for each property. Example: Write all of the Atomic radius values in red, all of the First ionization energies in green, and all of the Electronegativities in blue.

4. 186
496
0.9

5. Graphing Periodic Trends Activity
Observe the trends in each property as you go down the Alkali metal group, and as you go across Period 3.
Answer each of the statements written above and below the chart, completing each statement with the observed trend (increase or decrease)

6. Periodic Trends Objective:
I will be able to explain the trends for electronegativity, ionization energy, and atomic radius on the Periodic Table.

7. Electronegativity (electron affinity)
Periodic Trends
Atomic Size/Radius
Ionization Energy
Electronegativity (electron affinity)

8. Vocabulary
atomic radius: a measure of the size of an atom, usually the mean distance from the center of the nucleus to the boundary of the surrounding cloud of electrons
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)

9. Vocabulary
electronegativity: a measure of how strongly atoms attract bonding electrons to themselves, the higher the electronegativity, the greater an atom's attraction for electrons
effective nuclear charge: the net charge an electron experiences in an atom with multiple electrons (Zeff)
valence electrons: the number of electrons in the s and p orbitals of the outermost energy level

10. Alkali Metals – Group 1
Effective Nuclear Charge:

11. 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

12. Proton Pulling Power Number of energy levels are equal.
Zeff is greater for neon, electrons are pulled closer to the nucleus.

13. Equation: Zeff = Z - S Zeff = effective nuclear charge
Z = # of protons (nuclear charge)
S = screening electrons (inner electrons)
Ex. Oxygen: 1s22s22p4
Zeff = 8 – 2 = + 6
Ex. Fluorine: 1s22s22p5
Zeff = 9 – 2 = + 7
Fluorine has a greater effective nuclear charge.

14. Calculate “effective nuclear charge” for the following elements.
                                                                                                                                                                             +7 +1

15.  You try one! Write this out and solve in your notes.
What is Zeff for Mg?
Zeff = Z – S
1s22s22p63s2
Zeff = Z – S = 12 – 10 = 2

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

17. The Analogy with Atomic Radius
Prowler = Nucleus
His Charm = Effective Nuclear Charge
Ladies = Electrons

18. Shielding Electrons Using Desks in Your Classroom
Teacher’s Desk = Nucleus
= Electron

19. Atomic Radius (notice any patterns?)

20. Atomic Radius Trends in atomic radius:
Going down groups, radius increases:
each new period adds an energy level (electron shells), greatly increasing the radius.
Going across periods, radius decreases:
More valence electrons in the outer shell (energy level) and more protons in the nucleus create a stronger attraction so the electrons are pulled closer to the nucleus.
The nucleus has an increased effective nuclear charge with more protons.

21. Atomic Radius
Atomic radius: defined as ½ distance between neighbouring nuclei in a molecule or compound
Affected by:
1. Number of energy levels
2. Proton Pulling Power (the more protons, the stronger the pull on the electrons)

22. Atomic Radius Trends

23. 

24. First Ionization Energy
Ionization energy is the energy required to remove an electron from an atom in the gas phase.

25. Ionization Energies (notice any patterns?)

26. Ionization Energies (notice any patterns?)
Noble Gases require the most ionization energy, Alkali Metals require the least. The more valence electrons, the higher the ionization energy.

27. Ionization Energy Going down groups, ionization energy decreases:
Electrons further from the nucleus experience more screening of core electrons, so there is less effective nuclear charge.
Going across periods, ionization energy increases:
Increased effective nuclear charge holds electrons more tightly.

28. Electronegativity
Electronegativity is the measure of attraction of an atom for a shared electron in a compound.
Electronegativity can be used to predict the type of bond atoms will form in a chemical compound.

29. Trend for Electronegativity

30. Trends in Electronegativity:
Going down groups, electronegativity decreases:
the electrons are further from the nucleus, so there is less effective nuclear charge.
Going across periods, electronegativity increases :
as atoms have more valence electrons in the same energy level and more protons in nucleus they have a higher effective nuclear charge.

32. 

33. How do you know if an atom gains or loses electrons?
Think about Lewis Dot Structures of ions.
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.
Noble gases already have 8 (or 2 for He) so they don’t form ions very easily.

34. Ions
ion: an atom or group of atoms that has a positive or negative charge
Positive and negative ions form when electrons are transferred between atoms.

35. 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

36. Negative ions or (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.

37. 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.

38. 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 melting and boiling points 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)

39. 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.

40. 

41. Websites Crash Course #22 – Types of Chemical Bonds
Crash Course #23 – Polar and Non-Polar Molecules