Transcript Periodic trends

```Unit 5
The Periodic Table
The how and why
Newlands -1865
Arranged
known elements
according to properties & order
of increasing atomic mass
Law of Octaves – pattern of
chemical & physical properties
repeated every 8 elements
Mendeleev - 1869
1st periodic table (63 elements)
 Ordered by increasing atomic mass
 Predicted pattern of missing elements
 Started new rows and lined up
columns to organize elements with
similar properties
 Rearranged elements so similar
properties would line up correctly
 Created
The Modern Table
 Moseley-
determined the atomic number
for each known element.
 Elements are still grouped by properties
 Similar properties are in the same
column
 Ordered by increasing atomic number
 Added a column of elements Mendeleev
didn’t know about – noble gases
Periodic Law
When
elements are arranged in
order of increasing atomic
number, elements with similar
properties appear at regular
intervals
 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
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
1A
 The
2A
elements in the A groups 8A
0
are called the representative
3A 4A 5A 6A 7A
elements
Transition metals

The Group B
elements
 These
are called the inner
transition elements and they
belong here
Three Classes of Elements
Metals
Nonmetals
Metalloids
Metals
Metals
Ductile – drawn into wires
 Malleable – hammered into sheets
 All solid at room temperature (except
Hg- Mercury)
 Conductors of heat and electricity
 Families
– 1 - Alkali
– 2 - Alkaline Earth
– Transition (B groups)

 Group
1A are the alkali metals
 VERY reactive because one valence e• Found as compounds in nature
• Not including H!
 Group
2A are the alkaline earth metals
 Still highly reactive but not as much so
 as alkali metals (2 valence e-)
Transition Metals
 The
weird ones…
 May lose different #s of valence
electrons depending on the element
with which it reacts
 Less reactive than alkali or alkaline
earth metals
 Good conductors of electricity &
heat, ductile, malleable
Inner Transition Metals
1st
row = lanthanides
• Shiny metals similar in reactivity
to alkaline earth metals
2nd row = actinides
• Unstable nuclei – all radioactive
Non-metals
Non-metals
 Most
are gases, some solid, and 1
liquid (Br)
 More variation than metals
 Families
–Halogens (Group 17 or 7A)
–Noble Gases (Group 18 or 8A)
 Group
7A is called the Halogens
 Most reactive non-metals – 7 valence
React frequently with alkali metals
Group 8A are the noble gases
Low reactivity, very stable, inert
Metalloids or Semimetals
Metalloids
 Border
the staircase between
metals and nonmetals
 Properties – similar to metals and
nonmetals
Part 2
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
What we will look for
 Periodic
trends
• How those things vary as you go
across a period
 Group trends
• How those things vary as you go
down a group
 Why?
• Explain why these variations exist
Atomic Size
 Where
do you start measuring?
 The electron cloud doesn’t have a
definite edge.
 Scientists focused first on diatomic
elements -- measured more than 1
atom at a time
Atomic Size
}
Atomic
between two nuclei of molecule
Atomic Size - Periodic Trends
 The
positive nucleus pulls on electrons
 Periodic trend
• As you move across a period, elements
have more protons
• The charge on the nucleus gets bigger
• The outermost electrons of each
element are in the same energy level
• So there is more pull on the outermost
electrons as you move across
Periodic Trends
 As
you go across a period, the radius
gets smaller.
 Same outermost energy level
 More nuclear charge
 Pulls outermost electrons closer
Na
Mg
Al
Si
P
S Cl Ar
Atomic Size – Group Trends
 The
positive nucleus pulls on electrons
 Group Trend
• As you go down a group, you add
energy levels
• Outermost electrons not as attracted by
the nucleus
Shielding
 Increasing
numbers of
electrons between the
nucleus and the valence
electrons tends to
decrease the force
between the nucleus &
the valence electrons
+
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
 Moving down the group
• More energy levels
• More shielding
• Outer electron less
attracted
+
Three
No shielding
One
Two
shields
shield
shields
Group trends
 As
we go down a
group
• Each atom has
another energy
level
• More shielding
• The atoms get
bigger
H
Li
Na
K
Rb
Rb
K
Overall
Na
Li
Kr
Ar
Ne
H
10
Atomic Number
Atomic size increases,
IONIZATION ENERGY
Alkali
metals are more stable if
they lose an electron
Example
• Sodium ([Ne] 3s1)
• Getting rid of the 3s1 electron
makes sodium more stable and
creates a sodium ion (Na1+)
Ionization Energy
The
amount of energy
required to completely
remove an electron from a
neutral atom.
The energy required for the
1st electron is called the first
ionization energy
Ionization Energy
2nd ionization energy is the
energy required to remove the
second electron
Always greater than 1st IE
The 3rd IE is the energy required to
remove a third electron
Greater than 1st or 2nd IE
The
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
Group trends
As
you go down a group first IE
decreases
• Valence e- farther from nucleus
• More shielding
Periodic trends
 First
IE increases from left to right
across a period
• Increased nuclear charge from added
proton
• Electron shielding not an issue b/c
valence are all in same energy level
 Exceptions at full and 1/2 full orbitals
• Lower IE b/c offer stability to atom
Ionization energy
INCREASE
How to remember?
LO
HI
LO
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
Atomic number
First Ionization energy
Electronegativity
Electronegativity
 There’s
an electron tug of war
between atoms in a compound
 The tendency for an atom to attract
electrons to itself when it is
chemically combined with another
element
 How “greedy”
 Large electronegativity means the
atom pulls the electron towards itself
Group Trend
As
you move down a group
• More shielding
• Less attraction for electrons
• Lower electronegativity
Periodic Trend
As
you move across a period
from left to right,
• Nuclear charge increases
• Greater electronegativity
Electronegativity
INCREASE
How to remember?
LO
HI
LO
All 3 trends
LO
HI
LO
```