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Drill #6:11/11 & 11/12/13. What is the difference between Mendeleev’s and Moseley’s versions of the Periodic Table? Determine # of valence electrons using Periodic Table: Lithium - Beryllium (Be) - Neon - Phosphorus -. Periodic Trends Graphing Activity. Octet Rule. - PowerPoint PPT Presentation
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Drill pd 3 11/17/14
Rank in order of increasing ionization energy. Si, C, S, F, Ca
Rank in order of decreasing atomic radii. F, S, N
Rank in order of increasing atomic radius As, Te, P, Al
Rank in order of increasing electronegativity. Ge, P, As, Cs
HINT: Look at group trends first, then period trends
Answers
Ca, Si, S, C, F F, N, S P, Al, As, Te Cs, Ge, As, P
Drill 4A/4B 11/17& 18/2014 Differentiate between ionization energy
and electronegativity. What types of periodic trends are you
noticing for atomic radius, ionization energy and electronegativity?
Ionization energy is the energy needed to REMOVE one of the atom’s outermost electrons while an atom’s electronegativity is its ability to ATTRACT electrons to itself in a chemical bond.
Announcement
Conclusions have been graded. You have one week to turn back into me for ½ credit.
PBIS stuff
SWBAT:
Analyze the trends of ionization energy, electronegativity and atomic radius by completing a graphing activity.
Homework
-2 Review and Reinforcement – pd 4A and 4B
Graph
PERIODIC TRENDS GRAPHING ACTIVITY
GRAPH DATA
Construct 1 graph for ionization energy, electronegativity and atomic radius. Answer questions on the Worksheet.
Octet Rule
Atoms tend to gain, lose, or share electrons so that each atom has a full outermost energy level, which typically consists of 8 electrons (called an octet)
Periodic Table Patterns
Nuclear Charge
Nuclear charge-the charge in the nucleus or the number of protons
Across – increases Down – increases
Atomic Mass
Weighted average (based on mass and percent abundance of each naturally occurring isotope)
Across – increases Down – increases Remember there are some exceptions to
this – like Te and I
Atomic Radius
One-half the distance between the nuclei of identical atoms that are bonded together.
Ionization Energy (IE)
Ionization energy is the energy needed to REMOVE one of the atom’s outermost electrons
Electronegativity
An atom’s electronegativity is its ability to attract electrons to itself in a chemical bond
Trend for Atomic Radius
Atomic radius decreases as you go from the left to the right across a period. This is because there is an increased nuclear charge (protons) but no additional “shielding” electrons come between valence electrons and nucleus.
Atomic radius increases as you go down a group. Increased distance of electrons and additional electron energy levels shield valence electrons.
Shielding Effect
The shielding effect is the reduction of attractive force between the nucleus (+) and its outer electrons (-) due to the blocking affect of the inner electrons
Nucleus
Shielding electrons
Valence electrons ‘shielded’ by inner electrons
Snowman analogy
Trend for Ionization Energy
Ionization energy increases as you go from the left to the right across a period this means it’s MORE DIFFICULT to remove
an electron because of increased positive nuclear charge.
Ionization energy decreases as you go down a group this means it’s EASIER to remove an
electron because as atomic radius increases, outermost electrons are further away from nucleus.
Trend for Electronegativity
Electronegativity increases as you go from the left to the right across a period
Electronegativity generally decreases as you go down a group in the periodic table
** Because the noble gases form very few compounds, they do not have an electronegativity value.
Trend for Electronegativity (cont) The highest electronegativity values
are located in the upper-right hand corner of the periodic table Fluorine has the highest electronegativity
value of 4.0 The lowest electronegativity values are
located in the lower-left hand corner of the periodic table
It is NOT an amount of energy
Ionization Energy (IE) - review Ionization energy – this is the energy
needed to remove one of the atom’s outermost electrons
Metals have low IE Nonmetals have high IE
Trend for IE
1) IE’s decrease as you move down a group why?
Why?
The electrons in larger atoms are held less strongly by the nucleus, therefore as you move down a group, the IE decreases
Trend for IE
2) IE’s increase as you move from the left to the right across a period why?
WHY?
As you move across a period, there is a decrease in atomic radius, therefore causing the IE to increase
Removing electrons from positive ions With sufficient energy, electrons can be
removed from positive ions. The energy required to remove a second
electron from an atom is called its second IE, and so on…
The second IE will be greater than the first IE, the third IE greater than the second, etc. because as electrons are removed, fewer electrons remain to shield the attractive nuclear force.
Electron Affinity (EA)
Neutral atoms can acquire electrons. An atom’s electron affinity is the
energy change that occurs when an electron is acquired by a neutral atom.
Trend for EA
They change in irregular ways across a period and down a group
General rules: nonmetals have higher electron
affinities than do metals (this means it’s usually easier to add an electron to a nonmetal)
The halogens gain electrons most readily – higher electron affinities.
Trends Across the Periodic Table
PERIODIC TABLE OF THE ELEMENTS1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 181
H1.008
2
He4.003
3
Li6.941
4
Be9.012
5
B10.81
6
C12.01
7
N14.01
8
O16.00
9
F19.00
10
Ne20.18
11
Na22.99
12
Mg24.31
13
Al26.98
14
Si28.09
15
P30.97
16
S32.07
17
Cl35.45
18
Ar39.95
19
K39.10
20
Ca40.08
21
Sc44.96
22
Ti47.87
23
V50.94
24
Cr52.00
25
Mn54.94
26
Fe55.85
27
Co58.93
28
Ni58.69
29
Cu63.55
30
Zn65.39
31
Ga69.72
32
Ge72.61
33
As74.92
34
Se78.96
35
Br79.90
36
Kr83.80
37
Rb85.47
38
Sr87.62
39
Y88.91
40
Zr91.22
41
Nb92.91
42
Mo95.94
43
Tc[98.9]
44
Ru101.1
45
Rh102.9
46
Pd106.4
47
Ag107.9
48
Cd112.4
49
In114.8
50
Sn118.7
51
Sb121.8
52
Te127.6
53
I126.9
54
Xe131.3
55
Cs132.9
56
Ba137.3
57-70*
71
Lu175.0
72
Hf178.5
73
Ta180.9
74
W183.8
75
Re186.2
76
Os190.2
77
Ir192.2
78
Pt195.1
79
Au197.0
80
Hg200.6
81
Tl204.4
82
Pb207.2
83
Bi209.0
84
Po[209]
85
At[210]
86
Rn[222]
87
Fr[223]
88
Ra[226]
89-102**
103
Lr[262]
104
Rf[261]
105
Db[262]
106
Sg[263]
107
Bh[264]
108
Hs[265]
109
Mt[268]
110
Uun[269]
111
Uuu[272]
112
Uub[277]
114
Uuq[289]
116
Uuh[289]
118
Uuo[293]
*lanthanoids 57
La138.9
58
Ce140.0
59
Pr140.9
60
Nd144.2
61
Pm[145]
62
Sm150.4
63
Eu152.0
64
Gd157.3
65
Tb158.9
66
Dy162.5
67
Ho164.9
68
Er167.3
69
Tm168.9
70
Yb173.0
**actinoids 89
Ac[227]
90
Th232.0
91
Pa231.0
92
U238.0
93
Np[237]
94
Pu[244]
95
Am[243]
96
Cm[247]
97
Bk[247]
98
Cf[251]
99
Es[252]
100
Fm[257]
101
Md[258]
102
No[259]
Version Date 29 February 2000
Shielding Effect Stays the SameAtomic Radius decreasesIonization Energy IncreasesElectronegativity Increases
Trends Down the Periodic Table
PERIODIC TABLE OF THE ELEMENTS1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 181
H1.008
2
He4.003
3
Li6.941
4
Be9.012
5
B10.81
6
C12.01
7
N14.01
8
O16.00
9
F19.00
10
Ne20.18
11
Na22.99
12
Mg24.31
13
Al26.98
14
Si28.09
15
P30.97
16
S32.07
17
Cl35.45
18
Ar39.95
19
K39.10
20
Ca40.08
21
Sc44.96
22
Ti47.87
23
V50.94
24
Cr52.00
25
Mn54.94
26
Fe55.85
27
Co58.93
28
Ni58.69
29
Cu63.55
30
Zn65.39
31
Ga69.72
32
Ge72.61
33
As74.92
34
Se78.96
35
Br79.90
36
Kr83.80
37
Rb85.47
38
Sr87.62
39
Y88.91
40
Zr91.22
41
Nb92.91
42
Mo95.94
43
Tc[98.9]
44
Ru101.1
45
Rh102.9
46
Pd106.4
47
Ag107.9
48
Cd112.4
49
In114.8
50
Sn118.7
51
Sb121.8
52
Te127.6
53
I126.9
54
Xe131.3
55
Cs132.9
56
Ba137.3
57-70*
71
Lu175.0
72
Hf178.5
73
Ta180.9
74
W183.8
75
Re186.2
76
Os190.2
77
Ir192.2
78
Pt195.1
79
Au197.0
80
Hg200.6
81
Tl204.4
82
Pb207.2
83
Bi209.0
84
Po[209]
85
At[210]
86
Rn[222]
87
Fr[223]
88
Ra[226]
89-102**
103
Lr[262]
104
Rf[261]
105
Db[262]
106
Sg[263]
107
Bh[264]
108
Hs[265]
109
Mt[268]
110
Uun[269]
111
Uuu[272]
112
Uub[277]
114
Uuq[289]
116
Uuh[289]
118
Uuo[293]
*lanthanoids 57
La138.9
58
Ce140.0
59
Pr140.9
60
Nd144.2
61
Pm[145]
62
Sm150.4
63
Eu152.0
64
Gd157.3
65
Tb158.9
66
Dy162.5
67
Ho164.9
68
Er167.3
69
Tm168.9
70
Yb173.0
**actinoids 89
Ac[227]
90
Th232.0
91
Pa231.0
92
U238.0
93
Np[237]
94
Pu[244]
95
Am[243]
96
Cm[247]
97
Bk[247]
98
Cf[251]
99
Es[252]
100
Fm[257]
101
Md[258]
102
No[259]
Version Date 29 February 2000
Shielding Effect IncreasesAtomic Radius IncreasesIonization Energy DecreasesElectronegativity Decreases
Drill 11/19/14
Comparison of Na and Cl atoms1. Both atoms are in the same period.
Explain the differences in their sizes.
2. Predict the charge that an ion of each would have. Explain your answer.
3. Compare the ionization energy required to remove the first electron from each of these atoms.
4. Draw the ion for each atom. Make sure to represent their sizes relative to their original sizes.
Ignore sizes of atoms:
Answers to Drill1. Cl is smaller than Na because the increased
positive nuclear charge exerts a stronger pull on the e- shrinking the cloud.
2. A Na ion would have a charge of +1 and Cl ion, a charge of -1 because Na must lose an e- to have a full octet and Cl must gain an e-.
3. The IE of Cl would be higher because nonmetals hold on to their valence e- more tightly.
4. The Cl ion is now larger than the Na ion. Sodium lost its outermost orbital. The positive nuclear charge is felt more strongly.
