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The Periodic Table and Chemical Periodicity

The Periodic Table
and
Chemical Periodicity

I. The Periodic Table

A. Review

1. The most successful scheme for organizing the elements was
devised by which Russian chemist?
_______________________________
2. What are the rows and columns of the periodic table called?
_______________________________________________
3. Going from left to right along a row in the periodic table, how are the elements arranged?
_______________________________________________

4. Elements belong to what three general sections of the periodic
table?
____________________________________

5. What is the defining chemical property of the noble gases?
______________________________________________________
_____________________________________________

6. What is a valence electron?
________________________________________________________

7. What is the name given to the elements in Group 7A or 17?
__________________________________


8. What is the name given to the elements in Groups 1A (or 1)
and 2A (or 2)?
_________________________________________________________







B. Divisions

of

the

Periodic

Table












1. Hydrogen has unique properties.

2. A diagonal bold stepped line divides the periodic table into two parts.
Elements to the right of this line are __________________. Elements to the left are ____________. Many elements adjacent to this line (ex. Si and Ge) have both metallic and nonmetallic properties are called ___________.

3. The columns labeled p1 through p6 are called the p-block because the last valence electron enters a p orbital.

4. The columns labeled s1 and s2 are called the s-block because
the last valence electron enters a s orbital. The circled groups 3
and 4 (the s and p blocks) are called the ___________________________.

5. This is the d-block or _________________________. The columns are labeled d1 through d10 because the last valence electron enters a
d-orbital.

6. This grouping is the f-block. This block contains the __________________ or rare earth elements (top row) and the __________________ (bottom
row). The columns are labeled f1 through f14.
II. Using the Periodic Table to Write Electron Configurations

Use the periodic table to write the electron configurations (long form unless otherwise noted) for the following elements:

Be _____________________

Mg _____________________

Ca _____________________

Sr _____________________

Ba _____________________

Ra _____________________ (use shortened notation)


What is the trend in the electron configurations of the elements above?
What group do the element belong to?


F _____________________

Cl _____________________

Br _____________________

I _____________________

At _____________________ (use shortened notation)


What is the trend in the configurations above? What is the name of the group that these elements belong to?


Sc _____________________

Pb _____________________ (use shortened notation)
I


II. Periodic Trends

A. Atomic Size

- according to the quantum mechanical model, an atom does not have a sharply defined boundary

- atomic radius cannot be measured directly

- defined as one half the distance between the nuclei of two like atoms
























To understand the trend in atomic size, must consider:

___________________________ - positive charge felt by the outermost electrons in an atom; approximately equal to the atomic number minus the number of electrons in inner, complete levels

_________________________ - effect of the inner electrons on the reducing the attraction between the nucleus and the outer electrons

Effective nuclear charge (ENC):








Li Be B C

ENC=







N O F Ne

ENC=








Na

ENC=






As “effective nuclear charge” ____________________, outer electrons are pulled in more tightly, and atomic radius _____________________.
































Atomic radii of the main group elements are shown above. What trends do you see across the periods? down the groups?


In general: Atomic radii

1) ___________________________________________________________

2) ___________________________________________________________




B. Ionization Energy

________________________________ - the energy required to overcome the attraction of the nuclear charge and remove an electron from a gaseous atom

Write the Aufbau orbital diagram for Na, Cl, and Ar.


Na ____ ____ ____ ____ ____ ____
1s 2s 2p 2p 2p 3s


F ____ ____ ____ ____ ____
1s 2s 2p 2p 2p


Ar ____ ____ ____ ____ ____ ____ ____ ____ ____
1s 2s 2p 2p 2p 3s 3p 3p 3p

Questions:

1) Based on the filled orbitals, which element would seem more stable?


2) Comparing the electron configurations of Na and F, which element would most likely lose an electron, and which one most likely gain an electron? Why?

First ionization energy - energy required to remove the first outermost electron

Second ionization energy - energy required to remove the next electron

Third ionization energy - energy required to remove the next electron











Table of ionization energies p.402

































In general, ionization energy:

1) ___________________________________________________________

2) ___________________________________________________________




C. Ionic Size

__________________ - any atom or group of atoms with a positive charge

__________________ - any atom or group of atoms with a negative charge


Ex.

Mg ___________________________

Mg2+ ___________________________


















In general, ionic size:


1)________________________________________________________

2)____________________________________________________________
____________________________________________________________

D. Electronegativity

_____________________________ - tendency for the atoms of an element to attract electrons when they are chemically combined with other elements































In general, electronegativity:

1)___________________________________________________

2)___________________________________________________
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