periodictrends- chemistry for eight graders

History of the Periodic Table
• 1871 – Dimitri Mendeleev was the first scientist
to published an organized periodic table. He
arranged the elements according to: 1. Increasing
atomic mass 2. Elements w/ similar properties
were put in the same row
• 1913 – Moseley arranged the elements according
to: 1. Increasing atomic number 2. Elements
w/ similar properties were put in the same column

The Periodic Law
• Mendeleev understood the ‘Periodic Law’
which states:
• The properties of the elements are periodic
function of their atomic number.

The Periodic Law
• Atoms with similar properties appear in
groups or families (vertical columns) on the
periodic table.
• They are similar because they all have the
same number of valence (outer shell)
electrons, which governs their chemical
behavior.

Perio
d
Group
Metals
Non metals
Noble Gases
Representative

Group Names
Alkali
+1
Alkaline
Earth
Metals
+2 +3 -3 -2
Halogen
-1
Noble
Gases
0
H
1
He
2
Li
3
Be
4
B
5
C
6
N
7
O
8
F
9
Ne
10
Na
11
Mg
12
Al
13
Si
14
P
15
S
16
Cl
17
Ar
18

S & P block – Representative Elements
Metalloids (Semimetals, Semiconductors) – B,Si, Ge,
As, Sb, Te (properties of both metals & nonmetals)
Columns – groups or families Rows - periods
METALS
TRANSITION METALS
NONMETALS

Periodic Groups
• Elements in the same column have similar
chemical and physical properties
• These similarities are observed because
elements in a column have similar e-
configurations (same amount of electrons in
outermost shell)

Periodic Trends
• Periodic Trends – patterns (don’t always hold
true) can be seen with our current arrangement
of the elements (Moseley)
• Trends we’ll be looking at:
1. Atomic Size and Radius
2. Ionization Energy
3. Electronegativity
4. Electron Affinity
5. Metallic Property

Atomic Size
Atomic Size
• Size goes UP
Size goes UP on going down a
on going down a
group.
group.
• Because electrons are added
Because electrons are added
farther from the nucleus, there is
farther from the nucleus, there is
less attraction.
less attraction.
• Size goes DOWN
Size goes DOWN on going
on going
across a period.
across a period.

Atomic Radius
• Atomic Radius –
size of an atom
(distance from
nucleus to
outermost e-
)

Atomic Radius Trend
• Group Trend – As you go down a column,
atomic radius increases
As you go down, e-
are filled into orbitals that
are farther away from the nucleus (attraction
not as strong)
• Periodic Trend – As you go across a period
(L to R), atomic radius decreases
As you go L to R, e-
are put into the same
orbital, but more p+
and e-
total (more
attraction = smaller size)

Ionic Radius
• Ionic Radius –
size of an atom
when it is an
ion

Ionic Radius Trend
• Group Trend – As you go down a column, ionic
radius increases
• Periodic Trend – As you go across a period (L to
R), cation radius decreases,
anion radius decreases, too.
As you go L to R, cations have more attraction
(smaller size because more p+
than e-
). The anions
have a larger size than the cations, but also decrease
L to R because of less attraction (more e-
than p+
)

Ionic Radius
How do I remember this?????
The more electrons that are lost, the greater the
reduction in size.
Li+1
Be+2
protons 3
protons 3 protons 4
protons 4
electrons 2
electrons 2 electrons 2
electrons 2
Which ion is smaller?

Ionic Radius
How do I remember this???
The more electrons that are gained, the greater the
increase in size.
P-3
S-2
protons
protons 15
15 protons 16
protons 16
electrons 18
electrons 18 electrons 18
electrons 18

Ionization Energy
Ionization Energy
See Screen 8.12
See Screen 8.12
IE = energy required to remove an electron from an
IE = energy required to remove an electron from an
atom in the gas phase.
atom in the gas phase.
Mg (g) + 738 kJ ---> Mg
Mg (g) + 738 kJ ---> Mg+
+
(g) + e-
(g) + e-

Ionization Energy
ionization energy decreases
As you go down, atomic size is increasing (less
attraction), so easier to remove an e-
R), ionization energy increases
As you go L to R, atomic size is decreasing (more
attraction), so more difficult to remove an e-
(also, metals want to lose e-
, but nonmetals do
not)

Electronegativity
• Electronegativity-
tendency of an
atom to attract e-

Electronegativity Trend
electronegativity decreases
As you go down, atomic size is increasing, so less
attraction to its own e-
and other atom’s e-
R), electronegativity increases
As you go L to R, atomic size is decreasing, so there
is more attraction to its own e-
and other atom’s e-

Electron Affinity
Electron Affinity
A few elements
A few elements GAIN
GAIN electrons to form
electrons to form
anions
anions.
.
Electron affinity is the energy change
Electron affinity is the energy change
when an electron is added:
when an electron is added:
A(g) + e- ---> A
A(g) + e- ---> A-
-
(g) E.A. = ∆E
(g) E.A. = ∆E

Electron Affinity of Oxygen
Electron Affinity of Oxygen
∆
∆E is
E is EXO
EXOthermic
thermic
because O has an
because O has an
affinity for an e-.
affinity for an e-.
[He] 
 
  
O atom
EA = - 141 kJ
+ electron
O [He] 
 
  
- ion

Electron Affinity of Nitrogen
Electron Affinity of Nitrogen
∆
∆E is
E is zero
zero for N
for N-
-
due
due
to electron-
to electron-
electron
electron
repulsions.
repulsions.
EA = 0 kJ
[He] 
  
N atom 
[He] 
  
N-
ion 
+ electron

Reactivity
• Reactivity – tendency of an atom to react
• Metals – lose e-
when they react, so metals’
reactivity is based on lowest Ionization Energy
(bottom/left corner) Low I.E = High Reactivity
• Nonmetals – gain e-
when they react, so
nonmetals’ reactivity is based on high
electronegativity (upper/right corner)
High electronegativity = High reactivity

Metallic Character
• Properties of a Metal – 1. Easy to shape
2. Conduct electricity 3. Shiny
• Group Trend – As you go down a column, metallic
character increases
R), metallic character decreases (L to R, you are
going from metals to non-metals

Effective Nuclear Charge,
Effective Nuclear Charge,
Z*
Z*
• Z* is the nuclear charge experienced by
Z* is the nuclear charge experienced by
the outermost electrons.
the outermost electrons.
• Explains why E(2s) < E(2p)
Explains why E(2s) < E(2p)
• Z* increases across a period owing to
Z* increases across a period owing to
incomplete shielding by inner electrons.
incomplete shielding by inner electrons.
• Estimate Z* by --> [
Estimate Z* by --> [ Z - (no. inner electrons)
Z - (no. inner electrons) ]
]
• Charge felt by 2s e- in Li
Charge felt by 2s e- in Li Z* = 3 - 2 = 1
Z* = 3 - 2 = 1
• Be
Be Z* = 4 - 2 = 2
Z* = 4 - 2 = 2
• B
B Z* = 5 - 2 = 3
Z* = 5 - 2 = 3 and so on!
and so on!

Effective
Effective Nuclear Charge, Z*
Nuclear Charge, Z*
• Atom Z* Experienced by Electrons in
Valence Orbitals
• Li +1.28
• Be -------
• B +2.58
• C +3.22
• N +3.85
• O +4.49
• F +5.13
Increase in Z*
Increase in Z*
across a period
across a period

Shielding
• The less attracted to the nucleus, the more
shielded, thus lesser effective nuclear
charge.
• The effective nuclear charge on those outer
electrons is less, and so the outer electrons
are less tightly held.

Effective Nuclear Charge
• What keeps electrons from simply flying off
into space?
• Effective nuclear charge is the pull that an
electron “feels” from the nucleus.
• The closer an electron is to the nucleus, the
more pull it feels.
• As effective nuclear charge increases, the
electron cloud is pulled in tighter.

General Periodic Trends
General Periodic Trends
• Atomic and ionic size
Atomic and ionic size
• Ionization energy
Ionization energy
• Electron affinity
Electron affinity
• Electronegativity
Electronegativity
Higher effective nuclear charge.
Electrons held more tightly
Smaller orbitals.
Electrons held more
tightly.

periodictrends- chemistry for eight graders

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periodictrends- chemistry for eight graders