Download periodic table II 016

1
Periodic Table II
Section 1
Electron structure and the
periodic table
2
Recall that the Periodic table is
arranged according to properties?
From
reactive
metals
To less
reactive metals
sodium
Mercury
3
Recall that the Periodic table is
arranged according to properties?
helium
…and
inert (noble)
gases
non-metals
To
semi-metals
carbon
silicon
4
The Periodic table is also arranged
according to electron arrangement
8
Valence
electrons
1, 2, 3
Valence
electrons
4 to 7
Valence electrons
5
The Periodic table is also arranged
according to electron arrangement
So the properties must depend on
electron arrangement!
Periods (horizontal) 1, 2, 3 etc.
Periods have elements with same number
of occupied energy levels (shells) OR
Have the same outer (valence) shell
6
Periods (horizontal) 1, 2, 3 etc.
Periods have elements with same number of occupied energy
levels (shells) OR Have the same outer (valence) shell
Example:
Elements in period 3
have electrons in
three energy shells
OR
the 3rd shell
is the valence level.
Al e- config. = 2 – 8 – 3
7
Periods (horizontal) 1, 2, 3 etc.
Periods have elements with same number of occupied energy levels
(shells) OR Have the same outer (valence) shell
Within a period
electrons vary from 1 to 8
Valence electrons
Also,
the properties change as the
atomic number increases
8
Groups (vertical) 1,2,…18
Groups have elements with the same
number of valence electrons
1
2
Valence electrons
1 or 2
3 4
5
8
7
6
Elements with the same number of valence
electrons make for similar properties in groups
9
10
Click here for
Crash course chemistry
on Periodic Table
Test Your Understanding
1. What aspect of an atoms structure allows you
to predict its properties? How does the
arrangement of elements on the table support
your conclusion?
2. Why do the transition elements have such
similar but unique properties?
3. What change in the atomic structure of the
elements in a period accompanies the change
in properties we observe?
11
12
HDYK?
13
Section 2
Periodic Trends
> Size of atoms
> Strength of atoms
2A) Periodic Trends: Size of atoms
Atoms get
larger down
the table
How can you explain
these trends?
(how do the structure
of atoms differ?
But get smaller left to right
14
Firstly, how are atoms measured?
Shooting X-rays at a sample reveals
where the nuclei are located, but not
the electrons.
Atomic nuclei
15
Firstly, how are atoms measured?
If you assume that
the outer shell of
each atom reaches
half way between
two atoms.
Then ½ the distance
between the two
nuclei is each
atom’s Radius.
D/2
radius
radius
16
17
Secondly, what determines the radius?
The size of an atom depends on two factors:
#1
The
number of
electron
shells.
More shells
take up
more space
Atoms get
bigger!
Periodic Trends: Size of atoms
As atomic number
increases going down a
group each atom has an
additional energy level,
further from the nucleus.
Radius goes UP since the
valence shell is now
further from the nucleus.
18
Secondly,
19
The size of an atom depends on two factors:
#2
The pull
from the
protons in
the nucleus
20
Secondly, what determines the radius?
The size of an atom depends on two factors:
A stronger
nuclear
charge will
pull
electrons
closer.
so atoms are smaller.
21
As Atomic number increases across a
period, atoms are gaining more protons
in their nuclei (greater nuclear charge)
Radius decreases - since the extra
positive charge from the protons pulls
the electrons closer to the nucleus
11+
Large
12+
13+
14+
15+
16+
17+
Small
Summary: Atomic radii
Click here for
Khan academy tutorial
on atomic radius
larger
smaller
22
23
Test your learning
1. How is the size of an atom expressed? Why?
2. Which atom is larger; H or He? Why?
3. Which is larger; and argon atom or a potassium
atom? Why?
4. Describe the overall trend in atomic radius for
elements on the periodic table.
24
25
Section 2B
Introduction to chemical reactions
Atomic “strength”
26
Why do atoms react with one
another?
Recall that
the group
18 Noble
gases have
8 valence
electrons?
Recall that
the group
18 Noble
gases are
very stable?
Atoms with 8 valence electrons are very
stable since they require a lot of energy to
remove any of their electrons.
27
Why do atoms react with one
another?
Nonmetal
atoms like
oxygen can
become stable
by gaining
some
additional
electrons.
The new ion that is formed is no longer neutral
28
Why do atoms react with one
another?
Ions like this
now called
“oxide “ are
negative since
they have two
more electrons
than protons
Negative ions like this are called an-ions
29
Why do atoms react with one
another?
Its new symbol
is O2Its new electron
configuration is
now 2-8 like
neon
Negative ions like this are called an-ions
30
Why do atoms react with one
another?
Metallic atoms like
sodium can become
stable by losing some
electrons.
31
Why do atoms react with one
another?
Sodium’s electron
configuration was 2-8-1
Now becomes 2-8
Stable like neon.
32
Why do atoms react with one
another?
The symbol for the
sodium ion is Na1+
since it lose an electron it
now has less electrons
than protons and carries
an excess positive
charge.
Positive ions are called
Cat-ions
Forming an ion:
So atoms lose or gain electrons to form ions
Why?
Na+ = 11 P+ + 10 e= + 1 charge
O 2- = 8 P+ + 10 e= -2 charge
But, why are some atoms losers,
and other atoms gainers?
33
Atomic Strength: Ionization Energy
34
Energy is required to remove an electron from an atom.
This is called it ionization energy (energy to ion-ize)
Mg + 738 kJ ---> Mg+ + eEnergy
ion and electron
Notice energy is added to Mg to remove one electron.
This is called the FIRST ionization energy because we
removed only one of the outer electrons
Atomic Strength: Ionization Energy
35
Energy is required to remove an electron from an atom.
This is called it ionization energy (energy to ion-ize)
Mg+ + 1451 kJ ---> Mg2+ + eIt takes even more energy to remove a second electron.
This is the SECOND ionization energy.
Atomic Strength: Ionization Energy
36
Energy is required to remove an electron from an atom.
This is called it ionization energy (energy to ion-ize)
Metals like magnesium will lose electrons since
their ionization energies are relatively low.
Atomic Strength: Ionization Energy
37
Energy is required to remove an electron from an atom.
This is called it ionization energy (energy to ion-ize)
In other words, metals are “weak”
Atomic Strength: Ionization Energy
38
Energy is required to remove an electron from an atom.
This is called it ionization energy (energy to ion-ize)
F + 1681kJ ---> F+ + eRemoving an electron from a nonmetal requires
too much energy
Atomic Strength: Ionization Energy
39
Energy is required to remove an electron from an atom.
This is called it ionization energy (energy to ion-ize)
F + 1681kJ ---> F+ + eNonmetals like fluorine are too strong to lose
electrons
Atomic Strength: Ionization Energy
40
Energy is required to remove an electron from an atom.
This is called it ionization energy (energy to ion-ize)
F + 1681kJ ---> F+ + eIn fact, nonmetals gain electrons from metals
and become stable.
Atomic Strength: Ionization Energy
41
Energy is required to remove an electron from an atom.
This is called it ionization energy (energy to ion-ize)
F + e-  F- + 328 kJ
In fact energy is released when F gains and electron
Atomic Strength: Ionization Energy
42
Energy is required to remove an electron from an atom.
This is called it ionization energy (energy to ion-ize)
F + e-  F- + 328 kJ
This is called electron affinity
43
Trends in Ionization Energy
Ionization energy increases across a period because the
positive nuclear charge increases (more protons).
Period
3
11+
losers
496 kJ
12+
13+
14+
1012 kJ
15+
16+
17+
1251 kJ
As the nuclear charge gets stronger, they exert a
greater pull on the outer electrons
Trends in Ionization Energy
44
Ionization energy increases across a period because the
positive nuclear charge increases (more protons).
losers
Ionization energy increases
gainers
Metals (left side) are weak - lose electrons more easily: losers
Nonmetals (right side ) strong - don’t lose electrons: gainers
45
Going down a group
Atoms get larger: this
means there is a
greater distance
from the nucleus to
valence electrons.
So they get weaker.
Ionization
energy
decreases
and Ionization energy
decreases.
46
Shielding Effect
*Each additional energy
level adds electrons
that repel the valence
electrons
Ionization energy
decreases
Atoms with lots
of inner shells.
47
Like radius, ionization energy
follows the diagonal trend.
Non-metal atoms are strong
And small.
Metal atoms are weak
And large
48
49
Periodicity: a trend that occurs
at regular intervals
Click here for Khan academy Tutorial
Ionization energy trends
Notice how ionization energy increases,
then decreases with each new period.
50
What does diagram this illustrate?
Learning check
1. What is ionization energy? What does it indicate
about an element?
2. Which will have a higher ionization energy H or
He? Why?
3. Which will have a larger ionization energy Ar or K?
Why?
4. Describe the overall trend in ionization energy on
the periodic table.
51
52
53
Section 3
Other important trends
54
Ion Sizes
Forming
a cation.
Metal ions are SMALLER than the
atoms from which they come.
Loss of entire energy level so size
DECREASES.
55
Forming
an anion
Nonmetal ions are
LARGER than the
atoms from which
they come.
Additional electrons
repel each other
more so size
INCREASES
(“swells up”)
56
Think about it
Which of the following is the smallest:
O2- ion, Ne atom, or Mg2+ ion?
Why?
Hint: Each has 10 electrons.
But: How many protons?
+8
+10
oxide
Neon
8 protons
10 protons
+12
magnesium
12 protons
57
Think about it
Which of the following is the smallest:
O2- ion, Ne atom, or Mg2+ ion?
+8
+10
+12
Neon
oxide    middle
Largest
   magnesium
smallest
8 protons
10 protons
12 protons
As nuclear charge increases, the size will decrease
Test your understanding
1. What is the electron configuration for a magnesium
atom?
2. In terms of electrons, what will likely happen to
magnesium during a chemical reaction?
3. What will be its new electron configuration?
4. Draw the Bohr diagram for the magnesium atom and
the magnesium cation.
5. Which will be larger? Why?
58
Atomic # 12 = Mg
9 An atom of an element has a total of 12 electrons. An ion of the
same element has a total of 10 electrons. Which statement describes
the charge and radius of the ion?
+
-
12 P + 10 e
(1) The ion is positively charged and its radius is
smaller than the radius of the atom.
(2) The ion is positively charged and its radius is
larger than the radius of the atom.
(3) The ion is negatively charged and its radius
is smaller than the radius of the atom.
(4) The ion is negatively charged and its radius
is larger than the radius of the atom.
59
Loser metal
60
61
Chemical strength: Electronegativity
measure of the ability of an atom in a
molecule to attract electrons to itself.
Attraction for a pair of electrons
shared in a chemical bond
 Chemical bond 
Electronegativity: same as trends
in Ionization energy
62
Why no values for group 18?
Increases: stronger nucleus = stronger force
Decreases:
More
Energy
Levels
=
Weaker
force
decreases
increases
63
64
Which has a higher: Circle one in each pair
1st ionization
energy?
Mg or Ca ?
Al or S ?
Which is more
Cs or Ba ?
electronegative?
F or Cl ?
Na or K ?
Sn or I ?
Why?
HDYK?
65
Ionization Energy
66
Introduction to Chemical Reactivity
Atoms lose or gain
electrons to become
stable like group 18
elements
(8 valence e- ‘s)
Ex: Ne 2 - 8
What’s the magic number?
67
What elements sit beyond group 18?
Nonmetals gain
To become like
Noble gases
F: 2-7, becomes 2-8
Metals lose
To become
Like
Noble gases
Na: 2-8-1, becomes 2-8
68
Metals – group 1,2 and transitional, etc.
Chemically weak, tend to lose outer
electrons to stronger atoms
Ex: Mg 2 - 8 - 2 loses 2e–becomes 2 - 8 like Ne
–Becomes a positive ION Mg2+
–ION = charged “atom”
–CATION – a positively charged ion
69
gain 1 e-
Nonmetals – groups 15, 16, 17
Chemically strong tend to gain electrons
from weaker atoms
Ex: Fluorine atom F 2 - 7 gains 1 ebecomes 2 - 8 like Ne
• Becomes a negative fluoride ION: F1• ANION - a negative ion
70
71
13 When a lithium atom forms an Li+ ion,
the lithium atom
(1) gains a proton
(2) gains an electron
(3) loses a proton
(4) loses an electron
37 What is the total number of electrons in a
Cu+ ion?
(1) 28
(3) 30
(2) 29
(4) 36
72
F gains e-
73
Activity / Reactivity
Metals –”losers” must lose electrons during reactions
weakest (lowest electronegativity) are most active (francium)
Nonmetals – gainers must gain electrons during reactions
strongest (highest electronegativity) are most active (fluorine)
More
Active
metals
Fr
F
More
Active
nonmetals
74
Metallic / nonmetallic “character”
More
(loser)
metallic
character
More
(gainer)
Nonmetallic
character
75
33 As the elements in Group 17 on the Periodic Table are
considered from top to bottom, what happens to the atomic
radius and the metallic character of each successive element?
(1) The atomic radius and the metallic character both increase.
(2) The atomic radius increases and the metallic character
decreases.
(3) The atomic radius decreases and the metallic character
increases.
(4) The atomic radius and the metallic character both
decrease.
76
For questions 73 – 76 Refer to the table below and your knowledge of chemistry
Hint: standard temperature = 00C
77
181
98
64
78
1) How many elements are solids at STP?
2) Which element is a noble gas?
3) Letter Z below corresponds to a different element. Elements G,Q, L
and Z are in the same group on the periodic table as shown to the right.
4) Based on the trend in melting points for G, Q and L, estimate the
melting point of Element Z
79
80
Can you explain this trend
in ionization energy?
+19
1) K + 419 kJ ---> K+ + eEnergy
2) Na + 496 kJ ---> Na+ + eEnergy
+11
3) Mg + 738 kJ ---> Mg+ + eEnergy
4) Mg+1 + 1451 kJ ---> Mg+2 + eEnergy
Hint: compare the structures on the right 
+13