1. (Chemical Periodicity; WS – Periodic Properties)
Monday October 15, 2012
(Chemical Periodicity; WS – Periodic Properties)

2. Bell Ringer Monday,
Henry Moseley studied x­ray spectra of several elements in a row of the Periodic Table. He found that each element had one more proton than the element immediately to its left. Which statement BEST describes Moseley's contribution to the modern Periodic Table?
recognizing that elements in the periodic table have similar properties
arranging elements in the periodic table based on their properties
predicting the properties of missing elements in the periodic table
providing a basis for ordering elements in the Periodic Table

3. Bell Ringer Monday,
Which of the following describes the properties of noble gases?
A. They have high boiling points
B. They are highly reactive .
C . They  are  generally  unreactive.
D They are solid at room  temperature.
The number of valence electrons in Group 2
elements is __________ .
A 0
B 1
C 2
D 3

4. Announcements

5. Assignment Currently Open Summative or Formative?
Date Issued
Date Due
Date Into GradeSpeed
Final Day
QUIZ 6 S3
10/5
10/19
QUIZ 7 S4
10/12
10/26
WS – The Periodic Law F1
10/17

6. Question What is meant by the term “periodicity?”
In general terms, periodicity is defined as the quality of recurring at regular intervals .
Chemistry defines periodicity as similarities in chemical properties
recurring at regular intervals .
This leads to the Periodic Law, which states that the physical and chemical properties of the elements are periodic functions of their atomic numbers.

7. Five Properties of Elements that Display Periodicity
Atomic Radius Ionization Energy Electron Affinity Ionic Radius Electronegativity

8. On your metric staircase, this would be 9 steps below milli-
Atomic Radius
Atomic radius may be defined as one-half the distance between the nuclei of identical atoms that are bonded together.
A picometer (pm) is a metric unit of length equal to one trillionth of a meter.
On your metric staircase, this would be 9 steps below milli-
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1x10-12 meter

9. Atomic Radius
Atomic radii decrease from left to right across a period because of the increasing positive charge of the nuclei (more and more p+) pulling the e- closer and closer.
Generally speaking, the size of the atoms get smaller as you move from left to right across a period.
smaller
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10. Atomic Radius
Atomic radii increase from top to bottom down a group because of e- occupying sublevels in successively higher main energy levels.
Generally speaking, the size of the atoms get larger as you move from top to bottom down a group.
larger
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12. Even thought the lines in each period are not identical, they show a similarity that creates a pattern. This is NOT randomness!

13. Ionization Energy
Ionization energy may be defined as the energy required to remove one electron from a neutral atom of an element.
This property determines the extent to which an atom is capable of losing electrons in order to begin the chemical bonding process.
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14. Ionization Energy increases
In general, ionization energies of the main-group elements increase across each period. This increase is caused by increasing nuclear charge. A higher charge more strongly attracts electrons in the same energy level.
Generally speaking, the energy required for an atom to lose electrons increases as you move from left to right across a period.
increases
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15. Ionization Energy decrease
Among the main-group elements, ionization energies generally decrease down the groups. Electrons removed from atoms of each succeeding element in a group are in higher energy levels, farther from the nucleus.
Generally speaking, the energy required to lose electrons get less as you move down a group. Therefore, they are removed more easily.
decrease
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17. Electron Affinity .
Electron affinity is defined as the energy change that occurs when an electron is acquired by a neutral atom.
Most atoms release energy when they acquire an electron - the quantity of energy released is represented by a negative number.
Some atoms must be “forced” to gain an electron by the addition of energy - the quantity of energy absorbed is represented by a positive number.
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18. For example, adding an electron to a
Electron Affinity
In general, as electrons add to the same p sublevel of atoms with increasing nuclear charge, electron affinities become more negative across each period within the p block, although there are exceptions.
For example, adding an electron to a
nitrogen atom gives a half-filled p sublevel. This occurs much more easily than forcing an electron to pair with another electron in an orbital of the already half-filled p sublevel of a oxygen atom.
negativity increases
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19. the size effect predominates.
Electron Affinity
As a general rule, electrons add with greater difficulty down a group. This pattern is a result of two competing factors. The first is a slight increase in effective nuclear charge down a group, which increases electron affinities. The second is an increase in atomic radius down a group, which decreases electron affinities.
decrease
In general,
the size effect predominates.
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20. Ions
Ions are created when neutral atoms lose or gain electrons, and the number of protons no longer matches the number of electrons.
The atom is no longer electrically neutral and has an overall + or – charge.
If an atoms loses one or more electrons, it now has more protons than electrons and takes on a positive charge.
If an atoms gains one or more electrons, it now has more electrons than protons and takes on a negative charge.
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21. Ions Positive and negative ions have specific names.
A positive ion is known as a cation.
A cation is formed by the loss of one or more electrons always leads to a decrease in atomic radius because the removal of the highest-energy-level electrons results in a smaller electron cloud.
Also, the remaining electrons are drawn closer to the nucleus by its unbalanced positive charge.
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22. Ions Positive and negative ions have specific names.
A negative ion is known as an anion.
The formation of an anion by the addition of one or more electrons always leads to an increase in atomic radius.
This is because the total positive charge of the nucleus remains unchanged when an electron is added to an atom or an ion.
So the electrons are not drawn to the nucleus as strongly as they were before the addition of the extra electron.
The electron cloud also spreads out because of greater repulsion between the increased number of electrons.
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23. Ions
Within each period, the metals at the left tend to form cations and the non-metals at the upper right tend to form anions.
Lose electrons to form positive cations
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Gain electrons to form negative anions

24. Ionic Radius
Cationic radii decrease across a period because the electron cloud shrinks due to the increasing nuclear charge acting on electrons in the same main energy level. Starting with Group 15, in which atoms assume stable noble-gas configurations by gaining three electrons, anions are more common than cations.
Anionic radii decrease across each period for the elements in Groups 15–18. The reasons for this trend are the same as the reasons that cationic radii decrease from left to right across a period.
decrease
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25. Ionic Radius
As they are in atoms, the outer electrons in both cations and anions are in higher energy levels as one reads down a group. Therefore, just as there is a gradual increase of atomic radii down a group, there is also a gradual increase of ionic radii.
increase
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27. Electronegativity
The outer-most, highest energy level electrons in an atom’s electron cloud are called its “valence” electrons. Valence electrons are the ones that participate in chemical bonding – the force that hold elements together to form compounds. When two atoms are bonded together, one of the atoms may “tug” harder at the participating valence electrons harder than the other atom does. Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons.
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28. Electronegativity
Electronegativity values tend to increase across each period, although there are exceptions.
increase
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29. Electronegativity
Electronegativity values tend to either decrease down a group or remain about the same.
Decrease/remain same
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30. least electronegative
Electronegativity
The alkali and alkaline-earth metals are the least electronegative elements. In compounds, their atoms have a low attraction for electrons.
least electronegative
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31. Electronegativity
Nitrogen, oxygen, and the halogens are the most electronegative elements. Their atoms attract electrons strongly in compounds.
most electronegative
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32. Electronegativity
The Noble gases are unusual in that some of them do not form compounds and therefore cannot be assigned electronegativity values.
But when a Noble gas does form a compound, its electronegativity is rather high, similar to the values for the halogens.
are inter

33. Electronegativity
The combination of the period and group trends in electronegativity results in the highest values belonging to the elements in the upper right of the periodic table. The lowest values belong to the elements in the lower left of the table.
highest
lowest
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34. Electronegativity
The alkali metals cesium and francium have the lowest electronegativity values at 0.7. The halogen fluorine has the highest electronegativy value at 4.0.
highest
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lowest

37. Worksheet
Periodic Properties