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The Periodic Table
Chapter 5
Video clip
What Does Periodic Mean?
To answer: Turn to the picture on pg. 133, Fig. 1
What things occur periodically? Moon phases, magazine publications
Keep these in mind as we learn about the periodic table
History
Russian scientist Dmitri Mendeleev taught chemistry in terms of properties.
Mid 1800 - molar masses of elements were known.
Wrote down the elements in order of increasing mass.
Found a pattern of repeating properties.
The Table
Compare this table with our modern table
The Table
Mendeleev’s Table Grouped elements in columns by similar
properties in order of increasing atomic mass. Found some inconsistencies - felt that the
properties were more important than the mass, so switched order Example: Iodine after Tellurium
Found some gaps. Must be undiscovered elements. Predicted their properties before they were
found.
Interestingly…
Mendeleev never won a Nobel Prize. He was nominated shortly before his death,
but lost to Henri Moissan, who discovered Fluorine in 1906
Periodic Law
Why could most of the elements be arranged in the order of increasing mass, but a few could not? Henry Moseley- discovered that atomic number
should be the basis for the table, not mass Example: I = 53, Te = 52
Physical and chemical properties of the elements are periodic functions of their atomic numbers
The Modern Table 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. So a new group was formed
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
1A
2A 3A 4A 5A 6A7A
8A0
The elements in the A groups are called the representative elements
The group B are called the transition elements
These are called the inner transition elements and they belong here
Group 1A are the alkali metals Group 2A are the alkaline earth metals
Group 7A is called the Halogens Group 8A are the noble gases
Review
Rewrite the periodic law in your own words What group are the Alkali Metals in? Halogens?
Why do families have the same properties?
The part of the atom another atom “sees” is the electron cloud.
More specifically the outside (valence) orbitals. A “family’s” orbitals fill up in a regular pattern. The outside orbital electron configuration
repeats. The properties of atoms repeat.
1s1
1s22s1
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p63s23p64s23d104p65s1
1s22s22p63s23p64s23d104p65s24d10 5p66s1
1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1
H1
Li3
Na11
K19
Rb37
Cs55
Fr87
He2
Ne10
Ar18
Kr36
Xe54
Rn86
1s2
1s22s22p6
1s22s22p63s23p6
1s22s22p63s23p64s23d104p6
1s22s22p63s23p64s23d104p65s24d105p6
1s22s22p63s23p64s23d104p65s24d10
5p66s24f145d106p6
Alkali metals (column 1) all end in s1
Alkaline earth metals (column 2) all end in s2
really have to include He but it fits better later.
He has the properties of the noble gases.
Demo- Ca and Mg in test tube (pg. 142)
s2s1
S- block: Reactive Metals
Transition Metals -d block
d1 d2 d3 d4d5 d6 d7 d8 d9
d10
The P-block:Main Group Elements
p1 p2 p3 p4 p5 p6
F - block
inner transition elements
f1 f5f2 f3 f4
f6
f7 f8 f9 f10 f11 f12 f14
f13
Each row (or period) is the energy level for s and p orbitals.
1
2
3
4
5
6
7
D orbitals fill up after previous energy level so first d is 3d even though it’s in row 4.
1
2
3
4
5
6
7
3d
f orbitals start filling at 4f
1
2
3
4
5
6
7 4f
5f
Review
Create an acronym for remembering the position of the s, p and d blocks of the table
The s-Block Elements
1 s electron 2 s electrons
Alkali metals Alkaline-earth metals
Extremely reactive A little less reactive
silvery vary
soft Stronger, harder
Not found in nature Not found in nature
Melt at lower temps Melt at higher temps
Group 1 Group 2
Quantum Formula: ns1,2
Example problems Sample Problem A on pg. 143 Without looking at the periodic table, identify the
group, period, and block that [Xe] 6s2 is located Answer: group 2, sixth period, s block Write the electron config. For the Group 1 element
in the third period. Will it be more reactive or less? Answer: Grp 1, third period= 1s22s22p63s1
Must be more reactive, because it’s in
group 1
Quantum Formula: ns1,2
The d-Block5 orbitals, 10 e’ total
Transition metals
Extremely reactive
Conduct electricity
High luster
Less reactive
Not always the same outer e’configuration
Groups 3-12
d block
Quantum Formula: ns0-2 (n-1)d1-10
Example Problem
Sample Problem B on pg. 146 For d-block problems, identify group using
this formula: d + s Without using the periodic table, identify the
period, block and group of [Kr] 4d55s1
Answer: 5 period, d-block, group 6 Molybedenum
Quantum Formula: ns0-2(n-1)d1-10
p- Block
Group number – 10 for electrons
Main group elements
Properties vary
Nonmetals, metalloids, halogens
Not always the same outer e’configuration
Groups 13-18
Quantum Formula: ns2np1-6
Example
Sample Problem C on pg. 148 Write the outer electron configuration for the Group
14 element in the second period, name it and identify it as a metal, nonmetal or metalloid
Group number is higher than 12- so it’s the P block, second period makes n=2, group number -10 = 14-10= 4
So, you have 2 left over for the p’s 2s22p2
Quantum Formula: ns2np1-6
Halogens (p-block) Group 17 Gases, mostly Most reactive of nonmetals Seven valence electrons
Metalloids (p-block) Groups btw metals and nonmetals solids Electrical conductivity
Metals (p-block) Harder and denser than alkaline-earth
metals, but softer than d-block Found only in compounds
Problem Examples Without the table, write the outer electron
config. For Group 14 in second period. Name it and classify it as a metal, nonmetal or metalloid
Answer: 14= p block, 14-10= 4 so 2s22p2, carbon= nonmetal
Quantum Formula: ns2np1-6
f- Block
Sixth and seventh periods La-Hf = Cerium-Lutetium Ac-Rf = Thorium-Lawrencium Mostly lab made (sythetic)
f orbitals start filling at 4f
1
2
3
4
5
6
7 4f
5f
More Example Problems Name the block and group for each and identify as metal,
nonmetal or metalloid [Xe]4f145d96s1
Answer: d-block, group 10, Pt, metal (period 6) [Ne]3s23p6
Answer: p-block, group 18, Ar, nonmetal (period 3)
Group #
Group Config
Block Comments
1, 2 ns1,2 s One or two electrons in ns
3-12 ns0-1(n-1)d1-10 d Sum of electrons in ns and (n-1)d equals group number
13-18 ns2np1-6 p Number of electrons in np sublevel equals group number +/- 10
Nobel gasesHalogens
NonmetalsMetalloids
Metals
Alkali Alkali-Earth Transition
Driving Force of Atoms
Full Energy Levels are very stable Noble Gases have full orbitals. Atoms behave in ways to achieve noble gas
configuration.
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
Atomic Radius = half the distance between two nuclei of a diatomic molecule.
}Radius
What’s A Trend?
Name some fashion trends Color trends? Behavior trends?
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.
Periodic Trends As you go across a period the radius gets
smaller. They have the same energy level, though. More nuclear charge. Outermost electrons are closer.
Na Mg Al Si P S Cl Ar
Group trends
As we go down a group
Each atom has another energy level,
So the atoms get bigger.
HLi
Na
K
Rb
Ionic Size
Cations form by losing electrons. Form positive ions
Groups 1-3 Cations are smaller than the atom they come
from. Metals form cations
Ionic size
Anions form by gaining electrons. Form negative ions
Anions are bigger than the atom they come from.
Nonmetals form anions.
Ionization Energy- Pg. 153
The amount of energy required to remove an electron from an atom (only deals with losing an e’)
Removing one electron makes a +1 ion. The energy required is called the first ionization
energy. Measured in kilojoules per mole
If an atom has a low IE, it will release an electron easier than one with a high IE (making them more reactive) So, which group would have the highest IE?
Nobel gases
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 of 2nd IE.
Symbol First Second ThirdHHeLiBeBCNO F Ne
1312 2731 520 900 800 1086 1402 1314 1681 2080
5247 7297 1757 2430 2352 2857 3391 3375 3963
11810 14840 3569 4619 4577 5301 6045 6276
Symbol First Second ThirdHHeLiBeBCNO F Ne
1312 2731 520 900 800 1086 1402 1314 1681 2080
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 because The electron is further away. More shielding occurs
What’s shielding?
Shielding
The electron on the outside energy level has to look through all the other energy levels to see the nucleus
Inner shell electrons “shield” nuclear charge from outer shell electrons
Period trends
All the atoms in the same period have the same energy level.
Same shielding. Increasing nuclear charge So IE generally increases from left to right
A higher charge will more strongly attract electrons, holding them “hostage”
Firs
t Ion
izat
ion
ener
gy
Atomic number
He
He has a greater IE than H.
same shielding greater nuclear charge H
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li has lower IE than H
more shielding further away outweighs greater
nuclear charge
Li
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Be has higher IE than Li
same shielding greater nuclear
charge
Li
Be
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He B has lower IE than
Be same shielding greater nuclear
charge By removing an
electron we make s orbital half filled Li
Be
B
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
N
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
N
O
Breaks the pattern because removing an electron gets to 1/2 filled p orbital
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
N
O
F
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
N
O
F
Ne Ne has a lower IE
than He Both are full, Ne has more
shielding Greater distance
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
N
O
F
Ne Na has a lower
IE than Li Both are s1
Na has more shielding
Greater distance
Na
Firs
t Ion
izat
ion
ener
gy
Atomic number
Electron Affinity- Pg 157
The energy associated with adding an electron to an atom (only deals with gaining e’) Easiest to add to group 7A. A highly negative number = a high EA, that means
the atom will gain electrons easily Trends:
EA increases from left to right because atoms become smaller, with greater nuclear charge.
EA decrease as we go down a group. Also measured in kJ/mol
Electronegativity- Pg. 161
The tendency for an atom to attract electrons to itself when it is chemically combined with another element. How fair it shares the electron with the atom its bonding
with
Big electronegativity means it pulls the electron toward it- it’s a bully!
Atoms with large negative electron affinity have larger electronegativity.
Flouine is the boss!
Na Cl
Valence Electron
Electronegativity= 0.9Electronegativity= 3.0
So, Chlorine isn’t sharing the electron fairly with sodium, because it has such a large electronegativity
Na Cl
Valence Electron
Electronegativity= 0.9Electronegativity= 3.0
So, Chlorine isn’t sharing the electron fairly with sodium, because it has such a large electronegativity
Group Trend
The further down a group, the farther the electron is away from the nucleus because there are more energy levels Therefore, atoms are more willing to share these
electrons. Low electronegativity.
Periodic Trend
Cations let their electrons go easily Low electronegativity
Anions want more electrons- try to steal them High electronegativity
Ionization energy, electronegativity
Electron affinity INCREASE
Atomic size increases, shielding constant
Ionic size increases