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Ch. 5 “The Periodic Table”
Why is the Periodic Table important to me?
• The periodic table is the most useful tool to a chemist.
• It organizes lots of information about all the known elements.
• You get to use it on every test.
Pre-Periodic Table Chemistry …
• …was a mess!!!• No organization of
elements.• Imagine going to a
grocery store with no organization!!
• Difficult to find what you need.
• Chemistry didn’t make sense.
(CHAOS)
During the nineteenth century, chemists began to categorize
the elements according to similarities in their physical and chemical properties. The end result of these studies was our
modern periodic table.
How did chemists begin to organize the known elements?
Chemists used the properties of elements to sort them into groups Ex. Chlorine, bromine, and iodine have very similar chemical properties.
As the number of elements increased, chemists inevitably began to find patterns in their properties.
1780 - 1849
Model of TriadsIn 1817, Johann Dobereiner classified some elements into groups of three, which he called triads.The elements in a triad had similar chemical and physical properties.
Law of Octaves
1838 - 1898
In 1865, John Newlands suggested that elements be arranged in “octaves” because he noticed when he arranged the elements in order of increasing atomic mass certain properties repeated every 8th element.
LiBeBCNOF
NaMgAlSiPS
Cl KCa??
AsSeBr
He called this the “Law of Octaves”
because of its similarity to
musical octaves
Lightest to heaviest.LiBeBCNOF
NaMgAlSiPS
Cl
KCa??
AsSeBr
ABCDEFGABCDEFGABCDEFG
John Newlands Law of Octaves
Newlands' claim to see a repeating pattern was met with savage ridicule on its announcement. His classification of the elements, he was told, was as arbitrary as putting them in alphabetical order and his paper was rejected for publication by the Chemical Society.
The Modern Periodic Table
1834 - 1907
In 1869 Dmitri Mendeleev published a table of the elements organized by increasing atomic mass.He was trying to organize elements so his students could learn them more easily!
A. Mendeleev and Chemical Periodicity
• Mendeleev placed known information of elements on cards (atomic mass, density, etc…). He arranged them in order of increasing atomic masses, certain similarities in their chemical properties appeared at regular intervals. Such a repeating pattern is referred to as periodic.
1830 - 1895
At the same time, Lothar Meyer published his own table of the elements organized by increasing atomic mass.
• Both Mendeleev and Meyer arranged the elements in order of increasing atomic mass.
• Both left vacant spaces where unknown elements should fit.
So why is Mendeleev called the “Father of the Periodic Table” and not Meyer, or both?
Mendeleev published first!Could it be his dashing good looks?!
• Mendeleev left blank spaces in his table when the properties of the elements above and below did not seem to match. The existence of unknown elements was predicted by Mendeleev on the basis of the blank spaces. When the unknown elements were discovered, it was found that Mendeleev had closely predicted the properties of the elements as well as their discovery.
and the elusive element 32…
Predicted Properties
Observed Properties
Atomic weight 72 72.61
Density 5.5 g/cm3 5.32 g/cm3
Melting point 825 C 938 C
Oxide formula RO2 GeO2
Density of oxide 4.7 g/cm3 4.70 g/cm3
Dates (predicted and found)
1871 1886
Color Dark gray Gray-white
Gallium
Germanium
After the discovery of these unknown elements between 1874 and 1885, and the fact that Mendeleev’s predictions were amazingly close to the actual values, his table was generally accepted.
The Father of the Periodic Table
However, in spite of Mendeleev’s great achievement, problems arose when new elements were discovered and more accurate atomic weights were determined.
By looking at our modern periodic table, can you identify what problems might have caused chemists a headache?
Ar and KCo and NiTe and I
Th and Pa
1887 - 1915
In 1913, through his work with X-rays, Henry Moseley determined the actual nuclear charge (atomic number) of the elements*. He rearranged the elements in order of increasing atomic number.*“There is in the atom a fundamental quantity which increases by regular steps as we pass from each element to the next. This quantity can only be the charge on the central positive nucleus.”
Remember This…?!
Henry Moseley
His research was halted when the British government sent him to serve as a foot soldier in WWI. He was killed in the fighting in Gallipoli by a sniper’s bullet, at the age of 28. Because of this loss, the British government later restricted its scientists to noncombatant duties during WWII.
Periodic LawWhen elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.
The Current Periodic Table
• Mendeleev wasn’t too far off.• Now the elements are put in rows by
increasing ATOMIC NUMBER!!
• The vertical columns are called groups or families and are labeled from 1 to 18 (modern)
• or in A & B Groups (with Roman numerals)
Groups…Here’s Where the Periodic Table Gets Useful!!
• Elements in the same group have similar chemical and physical properties!!
• (Mendeleev did that on purpose.)
Why??• They have the
same number of valence electrons.
• They will form the same kinds of ions.
Groups in the Periodic Table
Elements in groups react in similar ways!
Periods in the Periodic Table
11
1
2
3
4
5
6
7
The horizontal rows are called periods and are labeled from 1 to 7.All elements in a period have the same number of energy levels (= to period #)
Energy LevelsEnergy Levels
n = 1
n = 2
n = 3
n = 4
In addition to Group Labels, many of the groups have Family Names
Group 1A: Alkali Metals
Group 1A: Alkali Metals
Cutting sodium metal
lithium
potassium
Alkali Metals • They are the most reactive metals.
• They react violently with water.
• Alkali metals are never found as free elements in nature - they are always in compounds with other elements.
• Only 1 valence electron
• Soft metals• Must be stored under
mineral oil, etc.
Group 2A: Alkaline Earth Metals
Only 2 valence electronsToo reactive to be uncombined in nature.
calcium
strontium
barium
Group 7A: The Halogens
7 valence electronsAll non-metalsVery reactiveAll physical states representedColored gases (always poisonous!)Occur as diatomic molecules when pure
fluorine
F2 Cl2Br2
chlorine bromine
Iodine I2
The Noble Gases
Noble Gases
• Noble Gases are colorless gases that are extremely un-reactive.(inert)
• They are inactive because their outermost energy level is full. (8 valence electrons – except He which has 2)
• Having 8 valence electrons is low in energy and, therefore, very stable.
Hydrogen• The hydrogen square sits atop
Family IA, but it is not a member of that family. Hydrogen is in a class of its own. (An orphan?)
• Like the Alkali metals, it only needs to lose one electron to be stable. (but it is not a metal!)
• Sometimes it’s shown above 7A.• Like the Halogens, it only needs to
gainone electron to have the stable Noble Gas electron configuration. (but it is not a Halogen!)
Hey Cameron, why are those elements by themselves on the bottom of the Periodic Table?!
I’ll handle this one, Cam! If they weren’t put on the bottom, the Periodic Table wouldn’t fit very nicely on a page! In fact, the table would look like this.
In fact, we have Glen Seaborg to thank for the fact that my Periodic Table doesn’t stick out of my notebook in a truly tasteless manner!
Glenn T. SeaborgAfter co-discovering 10 new elements, in 1944 he moved 14 elements out of the main body of the periodic table to their current location below the Lanthanide series. These became knownas the Actinide series.
1912 - 1999
“I was warned at the time that it was professional suicide to promote this idea, which has since been called one of the most significant changes in the periodic table since Mendeleev’s 19th century design. Luckily, I stuck to my guns and have seen the actinide concept become the foundation for many significant discoveries in heavy element research.”
Seaborgium
Glenn T. SeaborgHe is the only person to have an element named after him while still alive.
1912 - 1999
"This is the greatest honor ever bestowed upon me - even better, I think, thanwinning the Nobel Prize." 106
SgSeaborgium
271
There are many ways that we can break the PeriodicTable up into sections!
MetalsMetals
Metals
Metals are good conductors of heat. That's why a branding iron is made from metal. The heat transfers quickly to the animal's hide.
Metals also conduct electricity.
Notice that the Tesla coil sparks seek out metallic objects because they conduct electricity better than the
nonmetallic materials such as wood or soil.Metals are also malleable and can be bent or hammered
into various shapes.
Metals• What comes to mind?• Most elements are metals• Loosely held valence e-’s• Properties of metals:
1. Good conductors of heat and electricity (p)
2. High density (p)3. High melting points (p)4. Luster (p)5. Malleable (p)6. Ductile (p)7. 1, 2, or 3 valence electrons
Nonmetals
Nonmetals do not conduct heat well. The insulating tiles from the Space Shuttle are made from fibers of silicon and oxygen (silica=sand).
Nonmetals
• Opposite of metals• Properties of nonmetals:
1. Dull (no luster)
2. Do not conduct heat/elec.
3. Not ductile
4. Not malleable
5. All phases
6. Have 5, 6, or 7 valence electrons
• Form many compounds with metals
Metalloids(Semi-Metals)
• Means “metal-like”• Dividing line
between metals and nonmetals
• Al is the exception• Properties of both
metals and nonmetals
Four Main Categories of Elements
• Noble Gases- group 18 or 0 or 8A– s & p sublevels filled– 8 valence __s2…__p6
– Inert- not reactive- because of full outer shell of electrons
• Representative Elements also called main group elements- Groups 1A-7A– s & p partly filled– Includes alkali metals, alkaline earth
metals, and halogens
1A
2A 3A 4A 5A 6A7A
8A0
The elements in the A groups are called the representative elements
Four Main Categories of Elements
• Transition Metals – – Unfilled inner shells – outermost s & inner d sublevels contain
electrons – Hard & brittle
• Inner transition metals- – outermost s & nearby f sublevel contain
electrons – Lanthanides (4f) and actinides (5f)
RepresentativeRepresentative
Inner Transition Elements
Noble Gases
Using the Diagonal Rule is just so bothersome! I wish there was an easier way to figure out electron configurations!
Oh, but thereis! Watch this!
1s1
1s22s1
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p63s23p63d104s24p65s1
1s22s22p63s23p63d104s24p64d105s2 5p66s1
1s22s22p63s23p63d104s24p64d104f145s25p65d106s26p67s1
H1
Li3
Na11
K19
Rb37
Cs55
Fr87
1
2
3
4
5
6
7
Group 2A
Be
Mg
Ca
Sr
1s22s2
1s22s22p63s2
1s22s22p63s23p64s2
1s22s22p63s23p64s23d104p65s2
1H
2He
2 3Li
4Be
3 11Na
12Mg
4 19K
20Ca
5 37Rb
38Sr
6 55Cs
56Ba
7 87Fr
88Ra
s-block
Always for row you are on!
B
Al
Ga
Group 3A
1s22s22p1
1s22s22p63s23p1
1s22s22p63s23p64s23d104p1
The P-block p1 p2 p3 p4 p5 p6
Always for row you are on!
Electron Configurations in Groups• In atoms of the Group 1A elements below, there
is only one electron in the highest occupied energy level.
• In atoms of the Group 4A elements below, there are four electrons in the highest occupied energy level.
6.2
It’s always s1 forthe row it’s on!
Always s2p2 for the row they’re on!
Electron Configurations in Groups
– The Noble Gases• In atoms of the Group 8A elements below, there
are eight electrons in the highest occupied energy level.
6.2
Except for He, always s2p6 for the row they are on!
Chemical elements in d-block
Group →
3 4 5 6 7 8 9 10 11 12
↓ Period
4 21Sc
22Ti
23V
24Cr
25Mn
26Fe
27Co
28Ni
29Cu
30Zn
5 39Y
40Zr
41Nb
42Mo
43Tc
44Ru
45Rh
46Pd
47Ag
48Cd
6 71Lu
72Hf
73Ta
74W
75Re
76Os
77Ir
78Pt
79Au
80Hg
7 103Lr
104Rf
105Db
106Sg
107Bh
108Hs
109Mt
110Ds
111Rg
112Uub
Always 1 level in (back) from the row you’re on!
3d
4d
5d
6d
F - block inner transition elements
f1 f5f2 f3 f4
f6 f7 f8 f9 f10 f11 f12 f14
f13
Always go back 2 energy levels from the row you’re on!
4f
5f
s-block
Blocks of Elements
d-block p-block
f-block
Periodic Table e- configuration from the periodic
periodic table(To be covered in future chapters)
•B•2p1
1IA
18VIIIA
12
IIA13
IIIA14
IVA15VA
16VIA
17VIIA
2
33
IIIB4
IVB5
VB6
VIB7
VIIB8 9
VIIIB10 11
IB12IIB
4
5
6
7
H1s1
Li2s1
Na3s1
K4s1
Rb5s1
Cs6s1
Fr7s1
Be2s2
Mg3s2
Ca4s2
Sr5s2
Ba6s2
Ra7s2
Sc3d1
Ti3d2
V3d3
Cr4s13d5
Mn3d5
Fe3d6
Co3d7
Ni3d8
Zn3d10
Cu4s13d10
B2p1
C2p2
N2p3
O2p4
F2p5
Ne2p6
He1s2
Al3p1
Ga4p1
In5p1
Tl6p1
Si3p2
Ge4p2
Sn5p2
Pb6p2
P3p3
As4p3
Sb5p3
Bi6p3
S3p4
Se4p4
Te5p4
Po6p4
Cl3p5
Be4p5
I5p5
At6p5
Ar3p6
Kr4p6
Xe5p6
Rn6p6
Y4d1
La5d1
Ac6d1
Cd4d10
Hg5d10
Ag5s14d10
Au6s15d10
Zr4d2
Hf5d2
Rf6d2
Nb4d3
Ta5d3
Db6d3
Mo5s14d5
W6s15d5
Sg7s16d5
Tc4d5
Re5d5
Bh6d5
Ru4d6
Os5d6
Hs6d6
Rh4d7
Ir5d7
Mt6d7
Ni4d8
Ni5d8
Periodicity
When one looks at the chemical properties of elements, one notices a repeating pattern of properties when the elements are in order of
increasing atomic number.
1. Nuclear charge- the number of protons in the nucleus.
More protons = increased nuclear charge so increased attraction between the nucleus and electrons.
Think of the nucleus as a magnet – each extra proton makes the magnet more powerful at attracting electrons & holding them tight!
Explaining Periodic Trends
Why a property is higher/lower, bigger/smaller, etc.!
2. Shielding- lessens the attractive force of the nucleus for the valence electrons– caused by electrons in energy levels
between the nucleus and the valence electrons
Shielding increases as you go down a group
because there are more energy levels (more core electrons). Shielding stays the same as you move across a period because the number of energy levels is
staying the same.Which atom has more shielding? (A) K or Ca (B) Na or KWhich atom is smaller? (A) N or P (B) Li or K
Metallic Character
Atomic Radius
• Atomic Radius- half the distance between the nuclei of two atoms of the same element in a diatomic molecule
Atomic Radius
• Trend for atomic size- – Down a group, size increases
• Occurs because # of energy levels increases
*Makes a BIG difference in size!!• shielding also increases.
– Across a period, size decreases• # of protons increases (nuclear
charge increases), pulling electrons closershielding doesn’t change because electrons are added to the same energy level
Atomic Radius
Ionization Energy• Ionization Energy- energy needed to
remove an electron from an atom. • Outer shell electrons are easier to remove than
‘core’ electrons so it takes less energy to remove them!
Highest toward upper right corner of PT since
these atoms are smaller & their valence
electrons are closer to the nucleus
-so held more tightly
Trends in Ionization Energy
Periodic Trends
• Ionization energy– Down a group- decreases – because electrons are held more loosely
due to increased # of energy levels & increased shielding
– Across a period- increases because electrons are held more tightly due to increased nuclear charge (increased # of protons in the nucleus)
Ionization Energy
Ionization Energy
• There are big jumps in ionization energy whenever you try to remove an electron from an inner energy level!
Ionization Energy
Electron Affinity
• Electron affinity of an element is the energy given off when an atom (in the gas phase) gains an electron to form an ion
– Example: F(g) + e- F-(g) – Ho (ENERGY) = -328.0 kJ/mol
Trends in Electron Affinity
• It decreases down a group, because electron shielding blocks some of the attraction from the nucleus
• It increases across a period, because nuclear charge increases, attracting electrons more strongly.
Periodic Trends
• Electronegativity- tendency for the atoms of the element to attract electrons when the atoms are part of a compound
• Noble gases- no electronegativity values- don’t form compounds
• In general, metals have low EN and nonmetals have high EN.
• The actual amount of EN an atom has is indicated by a number on the Pauling Electronegativity Scale that goes from 0 to 4.
• Dr. Linus Pauling set up this scale and gave the element having the greatest EN an arbitrary number of 4, and he assigned numbers to the others relative to this element.
• Flourine is the most electronegative element at 4. (3.98) and Francium is the least electronegative at 0.7.
Periodic Trends
• Electronegativity Trends- – Down a group – decreases- electron shielding
results in less attraction for electrons by the nucleus
– Across a period- increases- higher atomic number and consistent electron shielding result in more attraction for electrons
• Electronegativity allows you to predict bond type: covalent- polar vs. nonpolar and ionic
Ionization Energy (IE)
Electron Affinity (EA) Electronegativity ( )c
Atomic & Ionic Radii
General Trends in the Periodic Table:
Decreasing
Increasing
IncreasingIncreasing
IE, EA, and c are useful concepts used to characterize different types of bonding and estimate bond energies.
Summary of Trends
6.3
Atomic Size Increases
Incr
ease
s
Decreases
Dec
reas
es
Size of cationsShieldingNuclear ChargeElectronegativityIonization energySize of anionsIonic size Constant
IONS • Remember – Atoms are neutral
• But…atoms can gain or lose electrons
(*# of protons NEVER changes during
reactions!)
• IONS are atoms or groups of atoms with
a charge.
• To tell the difference between an atom and an ion, look to see if there is a charge in the superscript!
• Examples: Na Ca I O• Na+ Ca+2 I- O-2
• When an atom loses an electron it gets a
positive charge (because it now has more
protons than electrons)
Mg --> Mg+2 + 2 e-
When an atom gains an electron it forms a negative ion (because it now has more electrons than protons)
F + e- --> F-
Atom versus Ion
Forming Cations & AnionsForming Cations & AnionsA CATION forms when an atom loses one or more electrons.
An ANION forms when an atom gains one or more electrons
Mg --> Mg2+ + 2 e-F + e- --> F-
Now has 12 protons & 10 electrons
Now has 9 protons & 10 electrons
–Metals have 1, 2, or 3 valence electrons so tend to lose electrons (to get an octet)- forming cations. (+ charge)
–Non-Metals have 5, 6, or 7 valence electrons so tend to gain electrons (to get an octet) - forming anions. (- charge)
Periodic Trends
• Ionic Radii Trends– Cations- smaller than neutral atom because
fewer electrons result in greater attraction by nuclei
– Anions- larger than neutral atom because more electrons result in less attraction by nuclei
– Within period- size decreases – Down a group – size increases
Forming cations
Forming anions
What would the charge be on a sodium ion?
EXAMPLE
Since sodium in in Group IA it has 1 valence e-
and so it would LOSE an electron
So it gets a charge of +1Remember an electron is negatively charged. When an atom loses electrons it forms positively charged ions.
When electrons are gained negatively charged ions form
It goes from 11 protons & 11 electrons to11 protons & 10 electrons
How would you write the symbol for the sodium CATION?
EXAMPLE
Na+1
How many outer electrons does sodium have before it loses one?
It has 1…remember the group number!
5
– 1. Which of the following sequences is correct for atomic size?• Mg > Al > S• Li > Na > K• F > N > B• F > Cl > Br
6.3 Section Quiz
6.3 Section Quiz
– 2. Metals tend to• gain electrons to form cations.• gain electrons to form anions.• lose electrons to form anions.• lose electrons to form cations.
6.3 Section Quiz
– 3. Which of the following is the most electronegative?• Cl• Se• Na• I
The Periodic Table
Summary of Trend• Periodic Table and Periodic Trends• 1. Electron Configuration
2. Atomic Radius: Largest toward SW corner of PT
3. Ionization Energy: Largest toward NE of PT4. Electron Affinity: Most favorable NE of PT
Periodic Table: electron behavior
• The periodic table can be classified by the behavior of their electrons1IA
18VIIIA
12
IIA13
IIIA14
IVA15VA
16VIA
17VIIA
2
33
IIIB4
IVB5
VB6
VIB7
VIIB8 9
VIIIB10 11
IB12IIB
4
5
6
7
West (South) Mid-plains East (North)METALS
AlkaliAlkaline
Transition
METALLOID NON-METALSNoble gasHalogensCalcogens
These elementstend to give up
e- and formCATIONS
These elementswill give up e- or
accept e-
These elementstend to accept
e- and formANIONS
ELEMENTS THAT EXIST AS DIATOMIC MOLECULES
ELEMENTS THAT EXIST AS DIATOMIC MOLECULES
Remember:
BrINClHOFThese elements only exist as
PAIRS. Note that when they
combine to make compounds, they
are no longer elements so they are no longer in
pairs!
1 2 3 4 5 6 7 8
Valence electrons
Select an element
= Internet link( )
Ions of representative elements have noble gas configuration Na is 1s22s22p63s1 Forms a 1+ ion - 1s22s22p6 Same configuration as neon Metals form ions with the configuration of the noble gas before them - they lose electrons
Configuration of Ions
This can explain why metals are shiny. This is the surface of copper at a ridiculously high magnification. The surface shows a lake of electrons along with ripples. The two islands are imperfections on the surface. Most likely a couple of atoms that aren't copper.