Transcript Periodic Table 09

New topic
The Periodic Table
The how and why
I.
The Modern Table
A. Organization of the P.T.
 Elements
are still grouped by properties
 Similar properties are in the same
column
 Order is in increasing atomic number
 Added a column of elements Mendeleev
didn’t know about.
 The noble gases weren’t found because
they didn’t react with anything.
 Horizontal
rows are called periods
 There are 7 periods
 Vertical
columns are called groups.
 Elements are placed in columns by
similar properties.
 Also called families
 The
1
2
elements in these groups 18
are called the representative
13 14 15 16 17
elements
VIIIB
IIB
VIIB
VIB
VB
13 14 15 16 17
3A 4A 5A 6A 7A
IB
VIIIA
VIIA
VIA
VA
IVA
IIIA
IIIB
1 2
1A 2A
IVB
IIA
IA
Other Systems
3 4 5 6 7 8 9 10 11 12
3B 4B 5B 6B 7B 8B 8B 8B 1B 2B
18
8A
B. Metals
Metals
Luster – shiny.
 Ductile – drawn into wires.
 Malleable – hammered into sheets.
 Conductors of heat and electricity.

C. Transition metals

The Group B
elements
Dull
D.
 Brittle
 Nonconductors
- insulators

Non-metals
Metalloids or Semimetals
Properties of both
 Semiconductors

 These
are called the inner
transition elements and they
belong here
 Group
1 are the alkali metals
 Group 2 are the alkaline earth metals
 Group
17 is called the Halogens
 Group 18 are the noble gases
Why?
 The
part of the atom another atom
sees is the electron cloud.
 More importantly the outside orbitals
 The orbitals fill up in a regular pattern
 The outside orbital electron
configuration repeats
 So.. the properties of atoms repeat.
H
Li
1
3
Na
11
K
19
Rb
37
Cs
55
Fr
87
1s1
1s22s1
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p63s23p63d104s24p65s1
1s22s22p63s23p63d104s24p64d105s2
5p66s1
1s22s22p63s23p63d104s24p64d104f145s
25p65d106s26p67s1
Periodic trends
Identifying the patterns
What we will investigate
 Atomic
size
• how big the atoms are
 Ionization energy
• How much energy to remove an
electron
 Electronegativity
• The attraction for the electron in a
compound
 Ionic size
• How big ions are
What we will look for
 Periodic
trends• How those 4 things vary as you go
across a period
 Group trends
• How those 4 things vary as you go
down a group
 Why?
• Explain why they vary
The why first
 The
positive nucleus pulls on
electrons
 Periodic trends – as you go across a
period
• The charge on the nucleus gets
bigger
• The outermost electrons are in the
same energy level
• So the outermost electrons are pulled
stronger
The why first
 The
positive nucleus pulls on
electrons
 Group Trends
• As you go down a group
–You add energy levels
–Outermost electrons not as attracted by
the nucleus
Shielding
 The
electron on the
outside energy level has
to look through all the
other energy levels to
see the nucleus
+
 The
Shielding
electron on the
outside energy level has
to look through all the
other energy levels to
see the nucleus
 A second electron has
the same shielding
 In the same energy level
(period) shielding is the
same
+
Shielding
 As
the energy levels
changes the shielding
changes
 Lower down the group
• More energy levels
• More shielding
• Outer electron less
attracted
+
Three
No shielding
One
Two
shields
shield
shields
Atomic Size
 First
problem where do you start
measuring
 The electron cloud doesn’t have a
definite edge.
 They get around this by measuring
more than 1 atom at a time
Atomic Size
}
Radius
Atomic
Radius = half the distance
between two nuclei of molecule
Trends in Atomic Size
Influenced
by two factors
Energy Level
Higher energy level is further
away
Charge on nucleus
More charge pulls electrons in
closer
Group trends
 As
we go down a
group
 Each atom has
another energy
level
 More shielding
 So the atoms get
bigger
H
Li
Na
K
Rb
Periodic Trends
 As
you go across a period the radius
gets smaller.
 Same shielding and energy level
 More nuclear charge
 Pulls outermost electrons closer
Na
Mg
Al
Si
P
S Cl Ar
Rb
K
Atomic Radius (nm)
Overall
Na
Li
Kr
Ar
Ne
H
10
Atomic Number
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 is called the first
ionization energy
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 First
H
He
Li
Be
B
C
N
O
F
Ne
1312
2731
520
900
800
1086
1402
1314
1681
2080
Second
Third
5247
7297
1757
2430
2352
2857
3391
3375
3963
11810
14840
3569
4619
4577
5301
6045
6276
What determines IE
 The
greater the nuclear charge the
greater IE.
 Increased shielding 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 of
More shielding
So outer electron less attracted
Periodic trends
 All
the atoms in the same period
• Same shielding.
• Increasing nuclear charge
 So IE generally increases from left to
right.
 Exceptions at full and 1/2 full orbitals
First Ionization energy
He
 He
H
has a greater IE
than H
 same shielding
 greater nuclear
charge
Atomic number
First Ionization energy
He
Li has lower IE than
H
 more shielding
 outweighs greater
nuclear charge

H
Li
Atomic number
First Ionization energy
He
Be has higher IE
than Li
 same shielding
 greater nuclear
charge

H
Be
Li
Atomic number
First Ionization energy
He
B has lower IE than Be
 same shielding
 greater nuclear charge
 By removing an
electron we make s
orbital full

H
Be
B
Li
Atomic number
First Ionization energy
He
H
Be
C
B
Li
Atomic number
First Ionization energy
He
N
H
C
Be
B
Li
Atomic number
First Ionization energy
He
 Breaks
N
H
C O
Be
the
pattern because
removing an
electron gets to
1/2 filled p orbital
B
Li
Atomic number
First Ionization energy
He
N F
H
C O
Be
B
Li
Atomic number
Ne
First Ionization energy
He
 Ne
N F
H
C O
Be
has a lower
IE than He
 Both are full,
 Ne has more
shielding
B
Li
Atomic number
Ne
First Ionization energy
He

N F
Na has a lower
IE than Li
Both are s1
 Na has more
shielding

H
C O
Be
B
Li
Na
Atomic number
First Ionization energy
Web elements
Atomic number
Driving Force
 Full
Energy Levels are very low
energy
 Noble Gases have full orbitals
 Atoms behave in ways to achieve
noble gas configuration
Ionic Size
 Cations
are positive ions
 Cations form by losing electrons
 Cations are smaller than the atom
they come from
 Metals form cations
 Cations of representative elements
have noble gas configuration.
Ionic size
 Anions
are negative ions
 Anions form by gaining electrons
 Anions are bigger than the atom they
come from
 Nonmetals form anions
 Anions of representative elements
have noble gas configuration.
Configuration of Ions
 Ions
of representative elements have
noble gas configuration
 Na is 1s22s22p63s1
 Forms a 1+ ion - 1s22s22p6
 Same configuration as neon
 Metals form ions with the
configuration of the noble gas before
them - they lose electrons
Configuration of Ions
 Non-metals
form ions by gaining
electrons to achieve noble gas
configuration.
 They end up with the configuration of
the noble gas after them.
Group trends
 Adding
energy level
 Ions get bigger as
you go down
H1+
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-
Size of Isoelectronic ions
 Iso
- same
 Iso electronic ions have the same #
of electrons
 Al3+ Mg2+ Na1+ Ne F1- O2- and N3 all have 10 electrons
 all have the configuration 1s22s22p6
Size of Isoelectronic ions
 Positive
ions have more protons so
they are smaller
Al+3
Na+1
Mg+2
Ne
F-1
O-2
N-3
Electronegativity
Electronegativity
 The
tendency for an atom to attract
electrons to itself when it is
chemically combined with another
element.
 How “greedy”
 Big electronegativity means it pulls
the electron toward it.
Group Trend
 The
further down a group
• More shielding
• more electrons an atom has.
 Less attraction for electrons
 Low electronegativity.
Periodic Trend
 Metals
- left end
 Low nuclear charge
 Low attraction
 Low electronegativity
 Right end - nonmetals
 High nuclear charge
 Large attraction
 High electronegativity
How to answer why questions
 Trend
• Periodic
• Group
 Reason
• Nuclear charge
• Energy level and shielding
 Result
• What happens to which electron
& Shielding
Energy Levels
Nuclear Charge
Ionization energy, electronegativity
INCREASE
Atomic size increases,
Ionic size increases
Other topics
 Reactivity
• Alkali metals most reactive
• Nobel gases least reactive
• Noble gases found as free elements
• Gr 1,2,17 - never found free
 Metallic properties
• Group
• Period
Properties of Groups
Hydrogen
•By itself
•Can gain or lose electrons
•Also shares electrons
•1 and 2
•Metallic characteristics
•Lose electrons,low electroneg and IE
•Never found in atomic state (compounds)
•Reactivity inc as you go down gr
•BP/MP no trend
Properties of Groups
 15
•
•
•
•
•
•
Nonmetallic to metallic change
N and P gain 3 electrons-nonmetal
As and Sb metalloids
Bi loses elecrons
N-fixing bacteria, less reactive than P
P- 4 bonds, more reactive than N
Properties of Groups
 16
• O & S gain 2 electrons nonmetals
• S loses 2 or 4 electrons
• O most important
– Very reactive
– Diatomic except photosynthesis
• Se & Te metalloids
• Po metal
Properties of Groups
 17
•
•
•
•
•
•
•
Halogens in free state
Gain 1 electron called halides
Nonmetals, w/ metal characteristics
All three states of matter at RT
Occur in nature as compounds
F most reactive halogen
BP/MP increases from top to bottom
Properties of Groups
 18
•
•
•
•
•
Monatomic molecules
He 2 valence electrons
The rest = 8 valence electrons
Some react w/ Florine
BP/MP increase from top to bottom