Introduction to the Periodic Table I am Dmitri Mendeleev! I made the PERIODIC TABLE !
The how and why History Dmitri Mendeleev u Russian scientist Dmitri Mendeleev taught chemistry in...
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Transcript of The how and why History Dmitri Mendeleev u Russian scientist Dmitri Mendeleev taught chemistry in...
History Russian scientist Dmitri MendeleevDmitri Mendeleev
taught chemistry in terms of properties.
Mid 1800 - molar masses of elements were known.
Wrote down the elements in order of increasing mass (this is wrong!!!!)
Found a pattern of repeating properties.
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.
Found some gaps. Must be undiscovered elements. Predicted their properties before they
were found.
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.
Vertical columns are called Groups.
Elements are placed in columns by similar properties.
Also called families
The group B are called the transition elements
These are called the inner transition elements and they belong here
Periodicity Explained Valence electron cloud Outside orbitals The orbitals fill up in a regular
pattern The outside orbital electron
configuration repeats The properties of atoms repeat when
placed in order of Atomic Number
1s1
1s22s1
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p63s23p64s23d104p65s1
1s22s22p63s23p64s23d104p65s24d10 5p66s1
1s22s22p63s23p64s23d104p65s24d105p66s2
4f145d106p67s1
H1
Li3
Na11
K19
Rb37
Cs55
Fr87
He2
Ne10
Ar18
Kr36
Xe54
Rn86
1s2
1s22s22p6
1s22s22p63s23p6
1s22s22p63s23p64s23d104p6
1s22s22p63s23p64s23d104p65s24d105p6
1s22s22p63s23p64s23d104p65s24d10
5p66s24f145d106p6
Alkali metals all end in s1
Alkaline earth metals all end in s2
really have to include He but it fits better later.
He has the properties of the noble gases.
s2s1 s- block
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.
Trends in Atomic Size Influenced by two factors. Energy Level Higher energy level - electrons
are further away from nucleus. Charge on nucleus – Zeff
More charge (# of protons) pulls electrons in closer.
Group trends As we go down a
group Each atom has
another energy level
So the atoms get bigger.
HLi
Na
K
Rb
Periodic Trends As you go across a period the radius
gets smaller. Same energy level. More nuclear charge = more protons. Outermost electrons are closer.
Na Mg Al Si P S Cl Ar
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.
IE1<IE2<IE3
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
What determines IE nuclear charge = IE distance from nucleus = IE Filled and half filled orbitals have
lower energy, so achieving them is easier, lower IE.
Shielding
Periodic trends: Ionization Energy (IE)
All the atoms in the same period have the same energy level.
Same energy level = same shielding, but each atom gains a proton, therefore….
Increasing nuclear charge …helps pull e- in tighter, therefore it is harder to remove.
So IE generally increases from left to right. Exceptions at full and 1/2 fill orbitals.
e-
e-
Periodic trends: Ionization Energy (IE)
As you go down a group first IE decreases because….
The electron is further away…and There are energy levels between
the nucleus and the e- thus…. Shielding the e- from + nucleus.
e-
e-
Shielding The electron on the
outside energy level has to look through all the other energy levels to see the nucleus
Shielding The 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.
Effective Nuclear Charge, Zeff
AtomZeff 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 Increase in Z* across a Z* across a periodperiod
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
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
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
Electron Affinity The energy change associated with
adding an electron to a gaseous atom.
Easiest to add to group 7A. Gets them to full energy level. Increase from left to right atoms
become smaller, with greater nuclear charge.
Decrease as we go down a group.
Periodic trends: Electron Affinity(EA)e-
e-
• From left to right atoms become smaller, with greater nuclear charge.
e - are attracted by the + charged nucleus. •therefore EA will increase from left to right
Periodic trends: Electron Affinity(EA)e-
e-
• Down a group atoms become larger, and have greater nuclear charge. • e - are attracted by the increased +charge but shielding also increases …which has a greater influence!! • therefore EA will decrease as you go down a group or family
• The energy change associated with adding an electron to a gaseous atom.
• e - are attracted by a + charge.– Therefore more protons more attraction– But remember “shielding”
• Easiest to add to group 7A.• Gets them to full energy level.• Increase from left to right atoms become smaller,
with greater nuclear charge.• Decrease as we go down a group.
Periodic trends: Electron Affinity(EA)
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 all have 10 electrons all have the configuration 1s12s22p6
Size of Isoelectronic ions Positvie ions have more protons so
they are smaller.
Al+3
Mg+2
Na+1 Ne F-1 O-2 N-3
• Na2O(s) + H2O(l) 2 NaOH(aq)
• CaO(s) + H2O(l) Ca(OH)2 (aq)
• MgO(s) + 2HCl(aq) MgCl2(aq) + H2O (l)
• NiO(s) + H2SO4 (aq) NiSO4 (aq) + H2O (l)
Metal Oxides are BASIC