Post on 30-Dec-2015
description
The How and Why
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HistoryElements known to ancients: C, Cu, Au, Fe,
Pb, Hg, Ag, S, Sn
Added before 1700: As, Sb, Bi, P, Zn
Dobereiner, Johann (1780-1849): arranged elements in triads (Ca, Sr, Ba; Cl, Br, I)
John Newlands (1837-1898): arranged elements in group of eight
Properties repeat every 8th elements:
Li, Be, B, C, N, O, F, Na
Na, Mg, Al, Si, P, S, Cl, K
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History
Dmitri Mendeleev (1834-1907): used the masses of elements as most of the masses were determined in XIX century (1869)
Arranged elements in order of increasing atomic masses
Found a pattern of repeating properties
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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 ( Te, I).
Found gaps in the trends- maybe undiscovered elements.
Predicted their properties before they were found
( eka boron – Sc; eka aluminum- Ga; eka silicon: Ge).
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The Modern Table
Elements are still grouped by properties.Similar properties are in the same column.Order is in increasing atomic number
(Moseley, 1914).Added a column of elements Mendeleev
didn’t know about (Noble Gases).The noble gases weren’t found because they
didn’t react with anything.
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Periodic Law
Mendeleev (1869): Properties of elements are a function of the atomic masses of the elements.
Modern periodic Law (Mosley, 1914) properties of elements are a periodic function of their atomic numbers ( # of protons in the nucleus). Explains Mendeelev’s irregularities.
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Horizontal rows are called periods There are 7 periods1
2
3
4
5
6
7
8
Vertical columns are called groups or families
Elements are placed in columns by similar properties.
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1A
2A 3A 4A 5A 6A7A
8A0
The elements in the A groups are called the representative elements (also numbered 1-18)
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The group B are called the transition elements
Inner Transition elements
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Group 1A(1) are the alkali metals
Group 2A(2) are the alkaline earth metals
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Group 7A (17) is called the Halogens
Group 8A (18) are the noble or inert gases
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Why the similarities in Properties?
The part of the atom another atom sees is the electron cloud.These are the outside or valence
orbitals. The orbitals fill up in a regular pattern.
The outside orbital electron configuration repeats.
The properties of atoms repeat.
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1s1
1s22s1
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p63s23p64s23d104p65s1
1s22s22p63s23p64s23d104p65s24d10 5p66s1
1s22s22p63s23p64s23d104p65s24d105p66s2
4f145d106p67s1
H1
Li3
Na11
K19
Rb37
Cs55
Fr87
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He2
Ne10
Ar18
Kr36
Xe54
Rn86
1s2
1s22s22p6
1s22s22p63s23p6
1s22s22p63s23p64s23d104p6
1s22s22p63s23p64s23d104p65s24d105p6
1s22s22p63s23p64s23d104p65s24d10
5p66s24f145d106p6
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Alkali metals all end in s1
Alkaline earth metals all end in s2
Have to include He but it fits better later.
He has the properties of the noble gases.
s2s1 S- block
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Transition Metals -d block
d1 d2 d3s1
d5 d5 d6 d7 d8s1
d10 d10
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The P-block p1 p2 p3 p4 p5 p6
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F - block inner transition elements
f1 f5f2 f3 f4
f6 f7 f8 f9 f10 f11 f12 f14
f13
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Each row (or period) is the energy level for s and p orbitals.
1
2
3
4
5
6
7
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d orbitals fill up after previous energy level … first d is 3d even though it’s in row 4.
1
2
3
4
5
6
7
3d
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f orbitals start filling at 4f
1
2
3
4
5
6
7 4f
5f
Writing Electron configurations the easy way
Review NotesReview Notes
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Electron Configurations Repeat The shape of the periodic table is a
representation of this repetition of electron configurations.
When we get to the end of the column the outermost energy level is full.
This is the basis for our shorthand.
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Trends and Properties of Elements in the Periodic Table
The following properties will be examined:
• Radius of the atom
• Ionization energy
• Electron affinity
• Electron negativity
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Effective Nuclear Charge Effective nuclear charge is experienced
by an outer electron at the outer edge of an atom
Zeff = Z – S Z is the atomic numberS is the number of core electrons.See blackboard for examples.
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Atomic Size (Radius)First problem where do you start measuring.
The electron cloud doesn’t have a definite edge.
Determined by measuring more than 1 atom at a time.
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Atomic Size
Atomic Radius = half the distance between two nuclei of a diatomic molecule.
}Radius
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Trends in Atomic Size
Influenced by two factors.
Energy Level
Higher energy level is further away from nucleus
Charge on nucleus ( Zeffective)
More charge pulls electrons in closer.
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Group Trends
As we go down a group
Each atom has another energy
level
So the atoms get bigger.
HLi
Na
K
Rb
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Periodic TrendsAs you go across a period the radius
gets smaller.Filling the same energy level.
***More nuclear charge (higher Zeff).Outermost electrons are pulled closer.
Na Mg Al Si P S Cl Ar
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Atomic Radii in the PT
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Overall
Atomic Number
Ato
mic
Rad
ius
(nm
)
H
Li
Ne
Ar
10
Na
K
Kr
Rb
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Ionization Energy (IE)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 e- is called the first ionization energy.
A(g) + IE → A(g) +1
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Ionization EnergyThe 2nd ionization energy is the energy required to remove the second electron.
2nd IE is always greater than 1st IE.
The 3rd IE is the energy required toremove a third electron.
3rd IE > 2nd IE > 1st IE
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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
810 14840 3569 4619 4577 5301 6045 6276
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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
810 14840 3569 4619 4577 5301 6045 6276
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What Determines IE
1. The greater the nuclear charge the greater IE.
2. Distance form nucleus influences IE.
3. Filled and half filled orbitals have lower energy, so achieving them is easier, lower IE.
4. Shielding effect
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ShieldingThe electron on the outside
energy level has to look through all the other energy levels to see the nucleus. It is shielded from the nucleus by all the inner electrons
A second electron in the same energy level has the same shielding.
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Group Trends
As you go down a group first IE decreases because
The electron is further away
More shielding by inner electrons as there are
more energy levels.
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Periodic trendsAll the atoms in the same period have the
same energy level.
Same shielding.
Increasing nuclear charge increases the force of attraction between the nucleus and the electrons.
IE generally increases from left to right.
Exceptions at full and 1/2 fill orbitals.
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Firs
t Ion
izat
ion
ener
gy
Atomic number
He
He >IE than H
same shielding
greater nuclear charge H
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Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li < IE than H
more shielding
further away (>n)
outweighs greater nuclear charge
Li
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Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Be > IE than Li
same shielding
greater nuclear charge
Li
Be
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Firs
t Ion
izat
ion
ener
gy
Atomic number
H
HeB < IE than Be
same shielding
greater nuclear charge
By removing an electron we make s-orbital half filled
Li
Be
B
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Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
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Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
N
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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
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Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
N
O
F
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Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
N
O
F
NeNe < IE than He
Both are full,
Ne has more shielding
Greater distance (>n)
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Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
N
O
F
NeNa < IE than Li
Both are s1
Na has more shielding
Greater distance
(>n)
Na
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Firs
t Ion
izat
ion
ener
gy
Atomic number
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Driving Force for Ionization
Full Energy Levels are very low energy.
Noble Gases have full orbitals.
Atoms behave in ways to achieve noble gas configuration.
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2nd Ionization Energy For elements that reach a filled or
half filled orbital by removing 2 electrons 2nd IE is lower than expected.
True for s2
Alkali earth metals form +2 ions.
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3rd IE Using the same logic s2p1
atoms have a low 3rd IE.
Atoms in the aluminum family form + 3 ions.
2nd IE and 3rd IE are always higher than 1st IE!!!
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Electron Affinity (EA)The energy change associated with adding an
electron to a gaseous atom.
Easiest to add to group 7A.Filled energy level.
EA increases from left to right atoms become smaller, with greater
nuclear charge.
EA decreases as go down a group.
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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.
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Ionic size
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.
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Ionic Radii
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Configuration of IonsIons always have noble gas configuration.
Na is 1s12s22p63s1
Forms a +1 ion - 1s12s22p6
Same configuration as neon.
Metals form ions with the configuration of the noble gas before
them - they lose electrons.
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Configuration of IonsNon-metals form ions by gaining
electrons to achieve noble gas configuration.
They form the configuration of the noble gas after
them.
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Group trends
Adding energy level
Ions get bigger as you go down a column
Li+1
Na+1
K+1
Rb+1
Cs+1
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Periodic TrendsAcross the period nuclear charge
increases so ions get smaller.
Energy level changes between anions and cations.
Li+1
Be+2
B+3
C+4
N-3O-2 F-1
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Size of Isoelectronic ions
Iso - same
Iso-electronic ions have the same # of electrons
Al+3 Mg+2 Na+1 Ne F-1 O-2 and N-3
10 electrons
configuration: 1s12s22p6
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Size of Atoms and Ions
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Size of Isoelectronic ionsPositive ions have more protons so
they are smaller.
Al+3
Mg+2
Na+1 Ne F-1 O-2 N-3
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Atomic Radii and Ionic Radii Compared
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Comparison af Atomic and Ionic Radii
Electronegativity
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Electronegativity The tendency for an atom to attract
electrons to itself when it is chemically combined with another element.
How fair it shares. Big electronegativity means it pulls the
electron toward it. Atoms with large negative electron
affinity have larger electronegativity. Scale designed by Linus Pauling
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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.
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Periodic Trend Metals are at the left end. They lose electrons easily
Low electronegativity
At the right end are the nonmetals. They gain electrons.
High electronegativity.
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Ionization energy, electronegativity
Electron affinity INCREASE
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Atomic size increases, shielding constant
Ionic size increases
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Summary ot trends in the Periodic Table
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Oxidation States
Check Blackboard
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Metal, nonmetals, metalloids, & noble gases
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Properties of Metals Elements with 3 or less electrons in outer
level Electrons loosely bound Become positive ions Good electric and heat conductors Malleable, ductile Most are solids at room temperature
(exception : mercury) Grayish, silvery in color (exceptions: Cu, Au)
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Properties of Nonmetals Elements with 5 or more electrons in outer
level Gain electrons easily to become negative
ions Brittle Can be solid, liquid, or gas at room
temperature Good insulators Many have colors in naturals state (sulfur –
yellow, iodine: purple)
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Transition Elements Located in B group elements (or groups
3-12) Many have multiple oxidation states
(because of the d-orbitals) Many form colorful compounds: Ni+2
green; Cu+2 – blue; Mn+7, purple All are metals, some with the highest
known melting points (Tungsten)
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Properties of Metalloids Located on the zig-zag line Behave both as metals and
nonmetals with some exhibiting stronger metallic character (aluminum)
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Noble Gases Located in group 8A(or 18) Most are inert (exceptions: Xe –
some compounds with oxygen and fluorine are known)
High ionization energies. Have octet of electrons (exception:
He with only 2 electrons)