Download Periodic Trends Review

Chemistry
Unit 4 Review
Know these things:
1. How to write the electron configuration for exceptional elements.
2. Write the electron configuration for ions.
3. Relate an electron configuration to an element’s position on the periodic table.
4. Find the number of valence electrons for an element.
5. The definitions and trends for reactivity, atomic radius, ionization energy,
electron affinity, and electronegativity.
A number of physical and chemical properties of elements can be predicted from
their position in the Periodic Table. Among these properties are Ionization Energy,
Electron Affinity and Atomic/ Ionic Radii.
These properties all involve the outer shell (valence) electrons as well as the inner shell
(shielding) electrons. Electrons are held in the atom by their electrostatic attraction
to the positively charged protons, the nuclear charge, Z. However, not all electrons
in an atom experience the same nuclear charge. Those closest to the nucleus
experience the full nuclear charge and are held most strongly. As the number of
electrons between the nucleus and the valence electrons increases, the apparent
nuclear charge decreases, due to the"shielding" of these inner shell electrons. Going
down a group increases the value of n, and increases the number of inner shell
electrons. This leads to better shielding and a weaker attraction between the nucleus
and the outer shell electrons. Going across a period leads to a larger nuclear charge, as
the number of protons increases. There is also an increase in the number of valence
electrons, but electrons in the same shell are poor at shielding each other. Going
across a row generally leads to a stronger interaction between the nucleus and the
valence electrons. The type of orbital holding the shielding electrons is also
important. The s orbitals are said to be penetrating - they have electron density close
to the nucleus. The best shielding comes from s orbitals, followed by p, d and f.
Ionization Energy
Ionization energy is the energy required to remove an electron from a gaseous atom in its
ground state. This is related to how "tightly" the electron is held by the nucleus. The higher
the ionization energy, the more difficult it is to remove the electron. For a many-electron
atom, the energy required for the reaction:
Energy + X (g) → X+ (g) + e-
First Ionization Energy (I1)
Since this requires an input of energy, it is an endothermic reaction, with a positive
energy value. The energy required for the reactions:
Energy + X+ (g) → X 2+ (g) + eEnergy + X 2+ (g) → X 3+ (g) + e-
Second Ionization Energy (I2)
Third Ionization Energy (I3)
When an electron is removed from an atom the repulsion between the remaining
electrons decreases. The nuclear charge remains constant, so more energy is
required to remove another electron from the positively charged ion. This means
that, I1 < I2 < I3 < ..., for any given atom. Going down a group the electrons become
increasingly easy to remove, since they are at an increasing distance from the
nucleus, with increasing numbers of shielding inner electrons. So, the ionization
energies decrease. Going across a period the ionization energies generally increase.
Electrons in the same set of orbitals do not shield each other very well but the
nuclear charge increases, making the electrons more difficult to remove.
Z
3
4
Atom
Li
Be
electron configuration
1s2 2s1
1s2 2s2
I1 (kJ/mol)
520
899
This trend is observed for Li and Be. The 1s electrons are the screening electrons for
both Li and Be. The nuclear charge for Li is +3. The nuclear charge for Be is +4 and
the 2s electrons in Be doesn't screen the nuclear charge very effectively. So an
electron in Be is harder to remove than an electron in Li. However, there are some
notable exceptions to this trend.
Z
Atom
electron configuration
I1 (kJ/mol)
2
1
3
Li
1s 2s
520
4
Be
1s2 2s2
899
2
2
1
5
B
1s 2s 2p
800
It is easier to remove an electron from B (Z = +5) than from Be (Z = +4).
Why? (Hint: what type of electron is being removed?)
For the next few elements the trend is followed. As Z increases, the I1 increases.
Z
Atom
electron configuration
I1 (kJ/mol)
3
Li
1s2 2s1
520
2
2
4
Be
1s 2s
899
5
B
1s22s2 2p1
800
2
2
2
6
C
1s 2s 2p
1090
7
N
1s2 2s2 2p3
1400
8
O
1s2 2s2 2p4
1310
There is another break between N and O.
Question 1: (2 points)
Fill in the energy diagram for each of these atoms:
N __ __ __ __ __
1s 2s 2px 2py 2pz
O __ __ __ __ __
1s 2s 2px 2py 2pz
How do the electrons being removed differ from each other?
The final two elements in this period follow the trend as Z increases, I1 increases.
Z
Atom
electron configuration
I1 (kJ/mol)
9
F
1s2 2s2 2p5
1680
2
2
6
10
Ne
1s 2s 2p
2080
More Questions:
2. Match the following electron configurations with the appropriate ionization
energies (I1). (3 points)
a) 1s2 2s2 3p6 3s2 3p6 4s2 3d10 4p6 5s1
b) 1s2 2s2 3p6 3s2 3p6 4s2
c) 1s2 2s2 3p6 3s2 3p6 4s2 3d104p6
i) 1356 kJ/mol
ii) 595 kJ/mol
iii) 409 kJ/mol
3. Which of the following would have the largest I1? Explain thoughtfully. (2
points)
Na
K
Li
Cs
4. What are the general periodic trends (i.e. trends as you go across a period) for:
a. ionization energy? (1 pt)
b. electronegativity? (1 pt)
c. atomic radius? (1 pt)
d. What is the major underlying concept that explains all of these trends?
Explain carefully. (1 pts)
6. What is the relationship between:
a. ionization energy and electronegativity? ( 1 pt)
b. ionization energy and atomic radius (1 pt)
c. electronegativity and atomic radius (1 pt)
7. Complete the table. (10 pts)
Element
Noble Gas Configuration
Predicted charge on its ion
Cs
Cl
Ga
At
S
8. Define Ionization Energy. (2 pts)
_____________________________________________________________________________
_____________________________________________________________________________
9. Write two separate equations to show the first and second ionization energy
of calcium. (Remember state symbols are important as they from part of the
definition). (2 pts)
First Ionization
_____________________________________________________
Second Ionization
_____________________________________________________
10. Which of the following elements (one from each pair) would you expect to
have the highest first ionization energy? Explain your answers.
Li and Fr
(2 pts)
_____________________________________________________
_____________________________________________________
K and Kr
(2 pts)
_____________________________________________________
_____________________________________________________
11. Define the term, shielding. (2 pts)
_______________________________________________________________________
_______________________________________________________________________
12. Consider the table of the first four ionization energies for element A shown
below.
IONIZATION ENERGY IN
KJ/MOL
1ST
578
2ND
1817
3RD
2745
4TH
11580
(i)
In which group does A appear on the periodic table? _______________(1
pt)
(ii)
Predict the formula of the compound A forms with chlorine. ________ (1
pt)
13. (3 pts) Arrange the following species in order of increasing size. Rb+, Y3+, Br-,
Kr, Sr2+ and Se2-.
SMALLEST
LARGEST
________________________________________________________________________
14. Define isoelectronic? Give examples of two different cations and two
different anions which are all isoelectronic. (3 pts)
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
15. Define electron affinity. (2 pts)
_______________________________________________________________________
________________________________________________________________________
16. Write an equation to summarize the process of second electron affinity of
oxygen. (Remember state symbols are important as they from part of the
definition). (2 pts)
________________________________________________________________________
17. Consider the table of ionization energies for element X shown below.
IONIZATION ENERGY IN
1ST
2ND
3RD
4TH
5TH
KJ/MOL
496
4562
6912
9544
(i)
In which group will X be found? (1) __________ (1 pt)
(ii)
Explain your answer. (2 pts)
13353
________________________________________________________________________
________________________________________________________________________
(iii)
Predict the formula of X’s oxide. ___________ (1 pt)
18. Explain thoughtfully why magnesium tends only to form a +2 ion. (2 pts)
_______________________________________________________________________
________________________________________________________________________
19. Why do elements in the same group react in similar ways? Explain well. (2
pts)
________________________________________________________________________
________________________________________________________________________
20. Identify any isoelectronic species in the following list. Fe2+, Sc3+, Ca2+, F-,
Co2+, Co3+, Sr2+, Cu+, Zn2+ and Al3+. (3 pts)
________________________________________________________________________
21. Arrange the following atoms into order of increasing first ionization energy.
Sr, Cs, S, F and As. (3 pts)
LOWEST
HIGHEST
________________________________________________________________________
22. For the following pairs of atom or ions, fill in the blank with either < or > to
correctly order the identified periodic trend. (12 pts)
Atomic radius
Cd
Ionic Radius
Si __ Cl
K+ __ Cs+
Br __ I
F- __ Ne
Au __ Hf
Ar __ K+
Mg __
Al 3+ __ N3-
Electron Affinity
Cl __ I
O __ F
Na __ Mg
23. Fill in the blank with > or <. (3 pts)
Atomic radius
Ionization energy
Na ______ S
Na ______ S
H __ He
Electron affinity
Na ______ S
24. What is a period? How many are there in the periodic table? (3 points)
25. What is a group (also called a family)? How many are there in the periodic
table? (3 points)
26. State the number of valence electrons in an atom of: (4 points)
a. oxygen
b. barium
c. iodine
d. phosphorous
27. Name the following families: (3 points)
GROUP
FAMILY
NAME
GROUP1
GROUP 2
GROUP 17
GROUP 18
Noble gases
28. What is the name given to the group of elements that have the following valence
shell electron configurations? (3 points)
VALENCE SHELL s2
s2p6
s2p5
s1
ELECTRON
CONFIGURATION
Alkali
FAMILY NAME
metals
29. Which alkali metal belongs to the sixth period? (1 point)
30. Which halogen belongs to the fourth period? (1 point)
31. What element is in the fifth period and the eleventh group? (1 point)
32. For each of the following pairs, identify which of the two has the larger value. (6
points)
CATEGORY
ATOMS OR IONS
WHICH HAS A LARGER
VALUE?
Electron affinity
F or Cl
Electron affinity
S or Se
Ionization Energy Na or Al
Atomic Radius
I or Br
Ionic Radius
Na+ or K+
Ionization Energy Fr or Ca