Transcript Chapter 6 Notes - Hingham Schools

Chapter 6
Atomic Structure
Hingham High School
Mr. Clune
Development of the Periodic Table
mid-1800s,
about 70 elements
Dmitri Mendeleev – Russian
chemist
Arranged elements in order of
increasing atomic mass
Thus, the first “Periodic Table”
Dimitri Mendeleev
Mendeleev Periodic Table
Mendeleev
 Left
blanks for undiscovered elements
 When
discovered, good prediction
 Problems
 Co
or inconsistencies?
and Ni; Ar and K; Te and I
New way
 Henry
Moseley – British physicist
 Arranged elements according to
increasing atomic number
 The arrangement today
 Symbol, atomic number & mass
Henry Moseley
Periodic table
Horizontal rows = Period
There are 7 periods
Periodic law:
Vertical
column =
Group
(or family)
Similar physical & chemical prop.
Identified by number & letter
Horizontal rows are called periods
 There are 7 periods

Vertical columns called groups
Elements are placed in columns by
similar properties
Also called families
Periodic Law
When elements are arranged in
order of increasing atomic number,
there is a periodic repetition of
their physical and chemical
properties.
Groups of Elements
The vertical columns in the
Periodic Table. Elements in
each group have similar
properties.
Groups of Elements
Example: Group 11
Copper
Silver
Gold
Metals - Metalloids – NonMetals
Metals
•80% of the elements
•Good Conductor
•Heat
•Electricity
•High luster or sheen
•Reflect light
Metals
•Ductile
(Pulled into wire)
•Malleable
(Hammered into sheets)
Non-Metals
•Non-Metals are not Metals
(Properties unlike Metals)
•Bad Conductors
(Except Carbon)
•Brittle
(Shatter)
Non-Metals
•Most are Gases
Oxygen, Nitrogen, Hydrogen!!
•Some are Solid
(Sulfur and Phosphorus)
•Some are liquid
Bromine is a dark brown liquid.
Areas of the periodic table
 Group
A elements = representative
elements
 Wide
range of phys & chem prop.
 Metals: electrical conductors, have
luster, ductile, malleable
Metalloids
•Staircase separating
Metals and Nonmetals
•Properties similar to
Metals and Nonmetals
•Sometimes acts like a metal.
•Sometimes acts like a nonmetal.
Silicon
•Pure Silicon - Nonmetal.
•Bad Conductor !!!
•Add Boron
•Conducts
electricity
•SemiConductor
Squares in a Periodic Table
•Atomic Number
•Electron
Configuration
•Chemical Symbol
•Name
•Mass Number
14
Si
Silicon
28.086
2
8
6
Noble Gases
Alkaline Earth Metals
Alkali Metals
Metalloids
NonMetals
Transition
Metals
Elements
Halogens
Electron Configuration
in
Groups
H
Li
1
3
Na
11
K
19
Rb
37
Cs
55
Fr
87
1s1
1s22s1
1s22s22p63s1
Group 1 –
1
_s
1s22s22p63s23p64s1
1s22s22p63s23p64s23d104p65s1
1s22s22p63s23p64s23d104p65s24d10
5p66s1
1s22s22p63s23p64s23d104p65s24d105p66
s24f145d106p67s1
Group 18 –
16
_p
1s2 He 2
Ne
2
2
6
1s 2s 2p
10
1s22s22p63s23p6 Ar18
1s22s22p63s23p64s23d104p6 Kr
36
1s22s22p63s23p64s23d104p65s24d105p6 Xe
54
1s22s22p63s23p64s23d104p65s24d10 Rn
5p66s24f145d106p6 86
s1
S- block
s2  Alkali metals all end in s1
 Alkaline earth metals all end
in s2
 really should include He,
but it fits better later.
 He has the properties of the
noble gases.
Transition Metals -d block
s1
d1
d2
d3
d5
s1
d5
d6
d7
d8
d9
10
d
The P-block
p1 p2
p3
p4
p5
p6
B C N O F Ne
F - block

inner transition elements
f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
1
2
3
4
5
6
7
 Each
row (or period) is the
energy level for s and p orbitals.
1
Transition Elements
2
3
4
5
6
7
Inner
Transition Elements
Atomic
Radius = half the distance
between two nuclei of a diatomic
molecule.
Atomic Size
}
Radius
Trends in Atomic Size
 Influenced
by three factors:
1. Energy Level
 Higher
energy level is further away.
2. Charge on nucleus
 More
 3.
charge pulls electrons in closer.
Shielding effect
(blocking effect?)
Group Trends
 As
we go down a
group...
 each atom has
another energy
level,
 so the atoms get
bigger.
H
Li
Na
K
Rb
Periodic Trends
 As you go across a period, the radius
gets smaller.
 Electrons are in same energy level.
 More nuclear charge.
 Outermost electrons are closer.
Na
Mg
Al
Si
P
S Cl Ar
Trends in Atomic Size
Size generally increases
Trends in Ionization Energy
The
amount of energy required to
completely remove an electron
from a gaseous atom.
Removing one electron makes a
1+ ion.
The energy required to remove
the first electron is called the first
ionization energy.
Ionization
Ion – A charged atom.
Gains or loses electrons
Cation – Ion with a positive charge
Loses electron(s)
Anion – Ion with a negative
charge
Gains electrons
+
Na
e-
Cl
Ionization Energy
The second ionization energy is
the energy required to remove the
second electron.
Always
greater than first IE.
The third IE is the energy required
to remove a third electron.
Greater than 1st or 2nd IE.
Symbol
H
He
Li
Be
B
C
N
O
F
Ne
First Second
Third
1312
2731
520
900
800
1086
1402
1314
1681
2080
11810
14840
3569
4619
4577
5301
6045
6276
5247
7297
1757
2430
2352
2857
3391
3375
3963
What determines IE
The greater the nuclear charge,
the greater IE.
Greater
distance from nucleus
decreases IE
Filled and half-filled orbitals have
lower energy, so achieving them
is easier, lower IE.
Group trends
As
you go down a group,
first IE decreases because...
The
electron is further away.
Periodic trends
All
the atoms in the same period
have the same energy level.
But,
So
increasing nuclear charge
IE generally increases from left
to right.
Trends in Ionization Energy
Energy generally
increases
Trends in Ionic Size
 Cations
form by losing electrons.
 Cations are smaller that the atom
they come from.
 Metals form cations.
 Cations of representative elements
have noble gas configuration.
Ionic size
Anions
form by gaining electrons.
Anions are bigger that the atom
they come from.
Nonmetals form anions.
Anions of representative elements
have noble gas configuration.
Group trends
Adding
level
Ions
energy
get bigger
as you go down.
Li1+
Na1+
K1+
Rb1+
Cs1+
Periodic Trends
Across
the period, nuclear charge
increases so they get smaller.
Energy level changes between
anions and cations.
Li1+
B3+
Be2+
C4+
N3-
O2-
F1-
Anions
decreases
Size increases
Trends in Ionic Size
Size of Cations
increases
Electronegativity
 The tendency for an atom to attract
electrons to itself when it is
chemically combined with another
element.
 How fair is the sharing?
 Big electronegativity means it pulls
the electron toward it.
 Atoms with large negative electron
affinity have larger electronegativity.
Group Trend
The
further down a group, the
farther the electron is away, and
the more electrons an atom has.
More willing to share.
Low electronegativity.
Periodic Trend
Metals
are at the left of the table.
They let their electrons go easily
Low electronegativity
At the right end are the
nonmetals.
They want more electrons.
Try to take them away from others
Trends
in
Periodic Table
Atomic Size Decreases
Ionization Energy Increases
Electronegativity Increases
Nuclear Charge Increases
Shielding is Constant
Shielding Increases
Nuclear Charge Increases
Electronegativity Decreases
Ionization Energy Decreases
Ionic Size Increases
Atomic Size Increases
Cation Size Decreases
Anion Size
Decreases
Homework
Worksheets 6-2, 6-3
Due: 11/22/04
Test: 11/23/04