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History and Trends of the Periodic Table
Year of Discovery
Ch.8:2,4,5,8,19,
34,35,37,51,56,
61,68
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Element Grouping
Johann Dobereiner – Noticed several “triads” of similarly
behaving elements (1829)
His practice of grouping elements led to the concept of
periodicity (The repeating of chemical properties on a
sequential basis)
Element Organization
John Newlands – Noticed periodicity based on elemental
mass and sorted elements by mass (1864)
• Law of Octaves – every 8th (known) element was similar
• His contribution wasn’t recognized till years later
Dmitri Mendeleev (1869) “Father of the periodic table”
• Also arranged elements by mass
and property, but included empty
space for irregular jumps in mass
• These accounted for many
undiscovered elements
successfully predicted later
Modern Periodic Law
• Problems still arose as new elements
were discovered and didn’t always fit
nicely when ordered by mass
Henry Moseley (Rutherford’s assistant)
• Discovered a method for counting each element’s
protons with x-rays
•Found that if elements were organized by atomic
number, and not mass, the problems disappeared. (1912)
Periodic law – Properties vary with their atomic numbers
in a periodic way. (Current basis of P.T. organization)
Periodic
Elements
Atomic Number - Number of protons
• Basis of Periodic Table organization.
• Distinguishing attribute of each element (always an
integer)
• Name - full name of element
• Symbol - 1,2, or 3 letter representative symbol
• Atomic Mass - weighted average mass of relative isotopes
• Atomic mass units (amu) or grams/mole
Groups (families) - Vertical columns
• Share chemical traits (properties)
• IUPAC - numbered groups 1 - 18
• NACPT notation - 1A - 8A / 1B - 8B
Periods (series) - Horizontal rows
• Indicates highest energy electron level n (1-7)
Other Divisions:
– Representative elements - (1A - 8A) (s and p block)
• Consistently "periodic" based on valence electrons
– Transition metals - (1B - 8B) (d and f block)
• Many exceptions exist for electron configurations
• Characteristics are less periodic
• f-block (inner-transition metals) (Lanthanide and
Actinide series)
IUPAC: 1-18
NACPT: 1A-8A
1B-3B
Classification of the Elements
*Noble gases are part of the
Representative elements
Valence electrons are the outer shell electrons. The
valence electrons are the electrons that take part in
chemical bonding.
Group
e- configuration
# of valence e-
1A
ns1
1
2A
ns2
2
3A
ns2np1
3
4A
ns2np2
4
5A
ns2np3
5
6A
ns2np4
6
7A
ns2np5
7
Core electrons: all non-valence electrons
Octet (8) Rule
• Many chemical properties are determined by an
atom’s valence electrons
• Atoms seek to obtain an electron configuration like
that of a Noble Gas
Neon
1s2 2s2
2p6
• A full outer electron shell has 8 valence electrons
“oct”-et (of representative elements)
• Helium reaches full valence with only 2 electrons
-1
-2
-3
+3
+2
+1
Cations and Anions Of Representative Elements
Metals (~75% of elements)
• Left and middle of periodic table
• Solid at RT (except Hg)
• Lustrous (shiny)
• Malleable (flattened out)
• Ductile (drawn to strips)
King of Random: The Metal Melter
www.youtube.com/watch?v=GCrqLlz8Ee0
• High melting points (~950 – 3,700 °C)
• Good conductors of heat and electricity
Nonmetals (~18% of elements)
• Top Right corner of periodic table + Hydrogen
• Typically gaseous or soft, crumbly solids at room
temperature (low melting temps).
• Exceptions: Bromine(l) and Cdiamond (s)
• Poor electrical conductors (Carbon is a thermal conductor)
• Combine with other non-metals to form molecules.
• Combine with metals to form Ionic compounds
Discovery of the 6 Noble Gases (1898 – 1900)
• Difficult to detect due to inert nature
• Ramsay reacted N2(g) with Mg(s) to form Mg3N2(s) along with
a volume of unknown gas that would not react.
• Using a discharge tube, they noted the
emission spectrum was unique.
• It was called Argon “the lazy one”
Sir William Ramsay
Metalloids (~7% of elements)
• Adjacent the Stair-step line that separates metals
from non-metals
• Have shared properties of both metals and non
metals
• Metallic luster
• Brittle or crumbly
• Fair conductors
Trans-Uranium Elements
– Only elements up to atomic 92 (Uranium) have been
found naturally occurring (trace amounts of Pu & Np)
– Nuclear reactions have produced the remaining
elements synthetically
Radioactive Elements
All elements after #83, starting
with Polonium, are radioactive
(element decays rapidly)
Crash Course: The Periodic Table
www.youtube.com/watch?v=0RRVV4Diomg
Periodic trend summary
• Effective Nuclear charge – nuclear pull “felt” by electrons
• Atomic radii - Distance between nucleus and outer e-
• Ionic radii - same as atomic radii, but distance for
each atom's common ion
• Ionization energy - energy for atom to lose an electron (lower
means more likely to lose an electron)
• Electron affinity - energy released when an atom gains an electron
(higher means more likely to gain electron)
• Electronegativity - measure of each atom's attraction
towards bonding electrons in a molecule
Effective nuclear charge (Zeff) is the attractive force
from the nucleus felt by an electron with all forces
taken into account.
Shielding effect: Repulsive forces from other electrons that
lessen the net force felt from the positive nucleus
Harder to remove, b/c no shielding
effects (e-/e- repulsions)
Effective Nuclear Charge (Zeff  Total – Core electrons)
Z
Core (e-) Zeff
Radius (pm)
11Na
11
10
1
186
12Mg
12
10
2
160
13Al
13
10
3
143
decreasing Zeff
Increasing Zeff
Additional P+ have
greater pull than
additional valence eGreater Zeff leads to
smaller atomic radius
Trends in Atomic Radii
Atomic radii increases going down group (adding n)
5
Period:
6
4
3
2
Atomic radii get smaller as atoms
get more massive across a period
Decreasing Atomic radii
Only representative elements shown
Zeff = 1.28
Zeff = 3.14
Zeff = 5.76
Per n, adding Protons has
greater effect than adding
valence e-
(133 pm)
(78 pm)
(152 pm)
(72 pm)
Cation is always smaller than atom from which it is formed.
Same # of protons (+) pulling in less electrons (-)
Anion is always larger than atom from which it is formed.
More electrons (-) being held by Same # of protons (+)
Comparison of Atomic Radii with Ionic Radii
Cations get smaller
Anions get larger
25
The Radii (in pm) of Ions of Familiar Elements
1 m = 1 x 1012 pm
Ionization energy is the minimum energy (kJ/mol) required to
remove an electron from a gaseous atom in its ground state. (Cation
forming)
Filled n=1 shell
Filled n=2 shell
Noble gases have full valence orbitals
= very stable = unreactive
Filled n=3 shell
Filled n=4 shell
Filled n=5 shell
Alkali metals have 1 valence e- and easily give it up
General Trends in First Ionization Energies
Decreasing First Ionization Energy
Increasing First Ionization Energy
Very Likely to
give away electrons
Very unlikely to
give away electrons
Higher energy levels electrons can more readily leave
(further from nucleus, more core electrons shielding ~ lower Zeff)
I1 < I2 < I3 < …
“Typical” last
Ionization:
Note the large jump
in energy afterward
Electron affinity is the negative of the energy
change that occurs when an electron is accepted
by an atom in the gaseous state to form an anion.
X (g) + e-
X-(g)
F (g) + e-
F-(g)
EA = +328 kJ/mol
O (g) + e-
O-(g)
EA = +141 kJ/mol
Larger release of energy indicates
more stable anion (more likely to gain electron)
Why would Alkali metals have a higher affinity than Alkaline Earth metals?
Why is Nitrogen 0? Half-full p-subshell is more stable than 4/6
Halogens (17/7A) most likely to gain one electron
Electronegativity - attraction of shared electron cloud
Linus Pauling mathematically
determined the measure of attraction
between an atom's nucleus and its
valence electrons by analyzing bond
strength of various compounds.
No values because
they
don't form
compounds
Periodic trend summary
• Effective Nuclear charge – nuclear pull “felt” by electrons
• Atomic radii - Distance between nucleus and outer e-
• Ionic radii - same as atomic radii, but distance for
each atom's common ion
• Ionization energy - energy for atom to lose an electron (lower
means more likely to lose an electron)
• Electron affinity - energy released when an atom gains an
electron (higher means more likely to gain electron)
• Electronegativity - measure of each atom's attraction towards
bonding electrons in a molecule
Practice Questions
1. a) Place the neutral atoms in order from smallest to largest
Radius: N, Na, P
b) Explain what is causing the change across a period (L to R).
2. a) Place the neutral atoms in order from smallest to largest
Ionization energy: K, Se, Rb
b) If an atom has a low value for Ionization energy, what action
is this specifically referring to and how likely will it happen?
3. The first six Ionization energies (kJ/mol) are listed sequentially
for Aluminum. Make two observations from the data.
Practice Questions
4. a) Place the atoms in order from smallest to largest electron
affinity energy: Ca, Se, Kr
b) If an atom has a low value for electron affinity, what action
is this specifically referring to and how likely will it happen?
5. What happens to the relative atomic radius of neutral Potassium
atoms and Bromine atoms after they ionize and combine to
form KBr?
6. Place the elements in order of increasing electronegativity:
F, As, N, Ne
Descriptive Chemistry
• Study of the elements and the compounds
they form.
• Physical and Chemical Properties
• Similar for each group/family
Hydrogen
•
•
•
•
•
Lightest and most abundant element
Non metal (though displayed in 1A)
Considered a family of its own
Colorless, odorless, and tasteless gas
Has chemical properties similar to both alkali
metals (reactivity) and halogens (physically).
• Occurs "di-atomically" H2, not H
1s1
The Alkali Metals Group 1A Elements
1
(ns )
M
M+1 + 1e- (loss of 1 valence e-)
2M(s) + 2H2O(l)
2MOH(aq) + H2(g) (Hydrogen displacement)
4M(s) + O2(g)
2M2O(s)
(Combustion/Oxide formation)
•
•
•
•
•
Very chemically reactive
Form +1 charge cations
Conductive and Lustrous
Soft at room temperature
React with water
• Lithium batteries
• 133Cs in atomic clocks
• Na+ and K+ : mediate conduction
across membrane synapses for
nervous system
Cs in slow-mo
Increasing reactivity
The Alkali Metals: Group 1A Elements (ns1)
Alkaline-Earth Metals/Group 2A Elements (ns2)
M
M(s) + 2H2O(l)
M+2 + 2e-
M(OH)2(aq) + H2(g) (M = Mg, Ca, Sr, or Ba)
Luminescent & radioactive
Alkaline-Earth Metals
• Solid at room temperature
• Forms +2 charge cations
• Reactive at higher periods
• Milk of Magnesia (Mg(OH)2)
• Strontium: flares, red fireworks
• Barium: rat poison,
gastrointestinal x-ray, green in
fireworks
Boron Family: Group 3A Elements
2
1
(ns np )
Tend to form +3 ions
Boron: found in lab glassware
Aluminum: very common in alloys (light weight)
Periodic Videos:
Gallium
Carbon Family: Group 4A Elements
2
2
(ns np )
• Typically do not form ionic bonds, but covalent
• Carbon: organic chemistry: vitamins/drugs. Main
component of all biomolecules (protein, fat, sugar, DNA)
• Silicon: heavily used in electronics
Nitrogen Family: Group 5A Elements (ns2np3)
• Tend to form -3 ions
• Nitrogen: most abundant atmospheric gas, found in all
proteins and DNA, very inert as N2
• Phosphorous: Very reactive, match heads; present in
DNA, ATP, and lipids)
Periodic Videos:
Phosphorous
Bi oxide
Oxygen Family: Group 6A Elements (ns2np4)
• Tend to form -2 anions
• Almost all life is sustained by aerobic respiration which requires
oxygen (electron carrier)
• O2 is required for combustion (fire)
• O forms many important compounds
(oxides) (CO2, NO3, PO4, SO4)
• Sulfur: pure form has distinct smell
of rotten eggs, found in 2 amino acids
Burning sulfur
Barking Dog
Properties of Oxides Across a Period
Oxygen can form
compounds with elements of
various groups
basic
acidic
Group 17/7A Elements (ns2np5): Halogens
59
•
•
•
•
Halogens
Pure forms are diatomic (ex. F2, Cl2)
Form salts when they react with metals
Large electronegativities; Highest e- affinity
Very reactive
• Bleach contains chlorine compounds
• Fluoride: toothpaste
• Iodine used as disinfectant
Periodic Videos:
Chlorine
Increasing reactivity
Group 17/7A Elements
2
5
(ns np ):
Noble Gases: Group 18/8A Elements
2
6
(ns np )
Completely filled ns and np subshells.
Inert: Highest ionization energy/low e- affinities
Low tendency to lose/accept electrons.
Colorless, odorless, and tasteless gases
Low boiling/ freezing points
Compounds of the Noble Gases
PtF6
gas
A number of xenon compounds exist: XeF4, XeO3, XeO4.
A few krypton compounds have been prepared, such as KrF2.
Transition Metals
Found to be “less periodic”
– Difficult to purify and thus harder identify
– Did not fit into major groups, overall similar properties
– Contain the Coinage and Precious metals (Au, Ag, Pt)
Necessary in trace amounts in living organisms
Fe binds O2 in Hemoglobin
Zinc present in many DNA
unwinding proteins
Periodic Videos:
Zinc
Inner Transition Metals
Also called the Lanthanide and Actinide Series
• Paramagnetic
Neodymium
• Actinides are radioactive and
most are synthetically made
Nd
U
Uranium:
U-235 used in
nuclear reactors
Nd2Fe14B alloy