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Anomalous Electron Configurations Anomalous electron configuration
• A few exceptions to the Aufbau of some elements!!!
principles exist. Stable Element Atomic Expected Actual
number configuration Configuration
Cr 24 3d4 4s2 3d5 4s1
– half-filled d shell: Cu 29 3d9 4s2 3d10 4s1
• Cr has [Ar]4s13d5; Mo 42 4d4 5s2 4d5 5s1
• Mo has [Kr] 5s14d5 *Pd 46 4d8 5s2 4d10 5s0
– filled d subshell: Ag 47 4d9 4s2 4d10 5s1
*Pt 78 5d8 6s2 5d9 6s1
• Cu has [Ar]4s13d10
Au 79 5d9 6s2 5d10 6s1
• Ag has [Kr]5s14d10.
• Au has [Xe]6s14f145d10 The explanation for this deviation lies in the superior stability
of completely filled or all half-filled orbitals than nearly filled
• Exceptions occur with larger of nearly half-filled orbitals.
****They are exception to this rule also. So, remember.
elements where orbital energies are Some other of this kinds are Nb(41), Ru(44), W(74),
An application of Electronic Cont’d
Configuration of Atom
§But at d & s orbital of outer shell, if there are 8/9/10 e
Locating elements in The Periodic Table then it is in the Group VIII
• In The Periodic Table elements are organized on the basis Ex: Fe (26) - 1s22s22p63s23p63d64s2 : Group VIII
of their electronic configuration
§If you have more then 10 e at d & s of outer shell
• The highest ‘n’ value in electronic configuration
then only ‘s’ orbital’s e will give the Group
determines the period of that element.
• No of outer shell e determines the Group Ex: Cu (29) - 1s22s22p63s23p63d104s1 : Group I
§ If at s & p orbital of the outer shell have 1 to 7 e (s1 to s2 • ‘A’ Sub Group: if at the outer shell you do not have ‘d’
p5) then they are elements of Group I to VII orbital or ‘d’ is filled only then it in ‘A’ Sub
Ex: Na (11) – 1s22s22p63s1 : Group I Ex: Cl (17) – 1s22s22p63s23p5 : Group VIIA
Cl (17) – 1s22s22p63s23p5 : Group VII Ga (31) - 1s22s22p63s23p63d104s24p1 : Group IIIA
§ If at s & p orbital of outer shell have s2 p6 then the • ‘B’ Sub Group: if you have e in ‘d’ orbital at outer
element is in group O. Ex: Ne (10) shell (d1–d10) then it is in ‘B’ Sub Group
§ Again at d & s orbital of outer shell determines the Group Ex: Sc (21) - 1s22s22p63s23p63d14s2 : Group IIIB
Ex: Sc (21) - 1s22s22p63s23p63d14s2 : Group III
Zn (30) - 1s22s22p63s23p63d104s2 : Group IIB
When an atom or molecule gain or loses an e- it becomes an
Valence Electrons ion
• A cation has lost e- and therefore has positive charge
• An anion has gained an e- and therefore has a negative
• The electrons in all the sub-shells with the charge.
highest principal energy shell are called the
• e in lower energy shells are called core Electron Configuration of Anions
• Anions are formed when atoms gain enough
• Starting with one valence electron for the first
element in a period, the number of electrons electrons to have 8 valence electrons
increases as you move from left to right across – filling the s and p sublevels of the valence shell
a period. • The sulfur atom has 6 valence electrons
• chemists have observed that one of the most S atom = 1s22s22p63s23p4
important factors in the way an atom behaves,
both chemically and physically, is the number • In order to have 8 valence electrons, it must gain
of valence electrons 2 more
S2- anion = 1s22s22p63s23p6
Electron Configuration of Cations The Periodic Table of Elements
• cations are formed when an atom loses all its valence Week -7
Lecture 13 &14
– resulting in a new lower energy level valence shell
– however the process is always endothermic Development of the Periodic Table
• the magnesium atom has 2 valence electrons • There were 114 elements known by 1999.
Mg atom = 1s22s22p63s2 • The majority of the elements were discovered
• when it forms a cation, it loses its valence electrons between 1735 and 1843.
Mg2+ cation = 1s22s22p6 • How do we organize 114 different elements in a
• for transition metals electrons, may be removed meaningful way that will allow us to make
from the sublevel closest to the valence shell
Al atom = 1s22s22p63s23p1
predictions about undiscovered elements?
Al+3 ion = 1s22s22p6 • Arrange elements to reflect the trends in chemical
Fe atom = 1s22s22p63s23p64s23d6 and physical properties.
Fe+2 ion = 1s22s22p63s23p63d6
Fe+3 ion = 1s22s22p63s23p63d5 • First attempt (Mendeleev and Meyer) arranged
Cu atom = 1s22s22p63s23p64s13d10 the elements in order of increasing atomic weight.
Cu+1 ion = 1s22s22p63s23p63d10
The Father of the Periodic Table —Dimitri Mendeleev
Introduction of The Periodic Table
The elements were first arranged in this way by Dmitri
Mendeleev, a professor at St. Petersburg University, in 1869. His
arrangement was based on atomic mass. – The periodic table is made up of rows of
When Mendeleev was setting out the table, only elements and columns.
63 elements had been discovered. His big idea – An element is identified by its chemical
was to leave gaps for yet to be discovered symbol.
elements. He was able to predict the properties of
some of these elements, including silicon and
– The number above the symbol is the atomic
boron. When his predictions were shown to be number
accurate his table became accepted, and it is the – The number below the symbol is the atomic
basis of the one we use today. weight of the element.
• Moseley later discovered that the periodic nature of the – A row is called a period
elements was associated with atomic number, not atomic – A column is called a family or group
– Elements are arranged left to right and top to
• The Periodic Law bottom in order of increasing atomic number
– When elements are arranged in order of increasing
atomic number, there is a periodic pattern in their – This order usually coincides with increasing
physical and chemical properties. atomic mass
Atoms & Elements
• Periodic Patterns
– The chemical behavior of elements is determined by
consist of its electron configuration
Metals or Non metals – The first three periods contain just A families. Each
period begins with a single electron in a new outer
that have that have electron shell.
– Each period ends with a completely filled outer
Chemical symbols shell that has the maximum number of electrons for
arranged in the Protons
– The outer shell electrons are responsible for
Make up the
are in chemical reactions. Elements in the same family
have the same number of outer shell electrons; so
they will have similar chemical properties.
Outer shell electrons
– Group A elements are called representative
has a elements
Mass Number Periodic
– Group B elements are called transition elements.
law Group number
• Chemical Families Metals, Nonmetals, and Metalloids
– IA are called alkali metals because the react with
water to from an alkaline solution, very soft metals.
•Metallic character refers to the properties of metals
(except H2) – Shiny or lustrous
– Group IIA are called the alkali earth metals because – Malleable (can be hammered into shape)
mostly we found them in soils as salts/minerals & – Ductile (can be drawn out into wires)
they are also reactive, but not as reactive as Group – All except mercury are solids at room temperature
IA. – They are sonorous (make a ringing sound when hit)
• They are also soft metals, though not as soft as alkali – In solution lose electrons in reactions - oxidized
metals – Most oxides are basic and ionic
Ex: Metal oxide + water → metal hydroxide
– Group VIIA are the halogens Na2O(s) + H2O(l) → 2NaOH(aq)
• These need only one electron to fill their outer shell – Tends to form cation in aqueous solution
• They are very reactive.( disinfectants, bleach) • Metallic character increases down a group.
• Metallic character decreases across a period.
– Group VIIIA are the noble gases as they have
completely filled outer shells
• They are almost non reactive. Only a few metals are magnetic.
Magnetism is not a property of most metals!
Metals, Nonmetals, and Metalloids Properties of Non-Metals
Metals § They are poor conductors of
• When metals are oxidized they tend to form electrical energy
characteristics cations. § They are poor conductors of thermal
Both a diamond and a pencil ‘lead’ are
made of the same element – carbon.
• All group 1A metals form M+ ions. energy
• All group 2A metals form M2+ ions. § Many of them are gases
• Most transition metals have variable charges. § They are brittle if they are solid
§ Form anions
§ Most oxides are acidic
Ex: nonmetal oxide + water → acid
P4O10(s) + H2O(l) → 4H3PO4(aq)
§ Gain electrons in reactions –
§ When nonmetals react with metals,
nonmetals tend to gain electrons:
metal + nonmetal → salt
2Al(s) + 3Br2(l) → 2AlBr3(s)
Metals, Nonmetals, and Metalloids The Groups of the Periodic Table
Metalloids • Group 1: The Alkali Metals
Metalloids have properties that are intermediate – Most reactive metals on the PT
between metals and nonmetals. – Rarely found free in nature
Example: Si (shown here) has a metallic luster but
– Charge of +1 = 1 valence electron
it is brittle.
• Group 2: The Alkaline Earth Metals
Metalloids have found fame in the semiconductor – Still quite reactive
industry. – Charge of +2 = 2 valence electrons
• Groups 3-12: Transition Metals
– Found freely and in compounds in nature
– Charge is usually +2 but can vary = usually 2 valence
• Group 13: Boron Family
– Charge is +3 = 3 valence electrons
The Groups of the Periodic Table Periodic Properties
• Group 14: The Carbon Family • Periodic law = elements arranged by
– Contains elements that can form unusual bonds atomic number gives physical and
(carbon and silicon) chemical properties varying periodically.
– Charge is +4 or -4 = contains 4 valence electrons • Various Elemental Properties change fairly
• Group 15: The Nitrogen Family smoothly going across a period or down a
– Charge is -3 = contains 5 valence electrons group.
• Group 16: The Oxygen Family
– Also known as the chalcogens
• We will study the following periodic
– Charge is -2 = 6 valence electrons trends:
• Group 17: The Halogens – Atomic radii
– Most reactive nonmetals
– charge is -1 = 7 valence electrons
– Ionization energy
• Group 18: The Noble Gases (The Inert Gases) – Electron affinity
– Nonreactive – Melting Points and Boiling Points
– Charge is 0 = 2 or 8 valence electrons
Electron Shells and the Sizes of Electron Shells and the Sizes of Atoms
• As the principle quantum number increases (i.e., we
Atomic Sizes move down a group), the distance of the outermost
• As a consequence of the ordering in the electron from the nucleus becomes larger. Hence, the
atomic radius increases.
periodic table, properties of elements vary • As we move across the periodic table, the number of
periodically. core electrons remains constant. However, the nuclear
• Atomic size varies consistently through the charge increases. Therefore, there is an increased
periodic table. attraction between the nucleus and the outermost
• As we move down a group, the atoms become electrons. This attraction causes the atomic radius to
Trends in Atomic Radius - Transition Metals
• As we move across a period, atoms become
smaller. • increase in size down the Group
• atomic radii of transition metals roughly the same
There are two factors at work: size across the d block
•principal quantum number, n, and – valence shell ns2, not the d electrons
•the effective nuclear charge, Zeff. – effective nuclear charge on the ns2 electrons approximately
Fig. 8.15 Atomic Radii for Main Group Elements
See Figure 8.16
Trends in Ionic Radius
• Ions in same group have same charge
• Ion size increases down the group
– higher valence shell, larger
• Cations smaller than neutral atom; Anions
bigger than neutral atom
• Cations smaller than anions
– except Rb+1 & Cs+1 bigger or same size as F-1 and O-2
• Larger positive charge = smaller cation
– for isoelectronic species
– isoelectronic = same electron configuration
• Larger negative charge = larger anion
– for isoelectronic series
Ionization Energy Ionization Energy
• minimum energy needed to remove an • The first ionization energy, I1, is the amount of
electron from an atom energy required to remove an electron from a
– gas state gaseous atom:
Na(g) → Na+(g) + e-.
– endothermic process
•The second ionization energy, I2, is the energy
– valence electron easiest to remove
required to remove an electron from a gaseous
– M(g) + IE1 → M1+(g) + 1 e- ion:
– M+1(g) + IE2 → M2+(g) + 1 e- Na+(g) → Na2+(g) + e-.
• first ionization energy = energy to remove electron
from neutral atom; 2nd IE = energy to remove from +1
The larger ionization energy, the more difficult it
is to remove the electron.
• IE increases (irregularly) as you move from There is a sharp increase in ionization energy
left to right across a period. when a core electron is removed.
• IE decreases (irregularly) as you move down a
group. 39 40
General Trends in 1st Ionization
• larger the effective nuclear charge on
the electron, the more energy it takes to
• the farther the most probable distance
the electron is from the nucleus, the less
energy it takes to remove it
• 1st IE decreases down the group
– valence electron farther from nucleus
• 1st IE generally increases across the
– effective nuclear charge increases
Tro, Chemistry: A Molecular
Example – Choose the Atom in Each Pair Irregularities in the Trend
with the Higher First Ionization Energy • Ionization Energy generally increases from left
1) Al or S to right across a Period
2) As or Sb • except from 2A to 3A, 5A to 6A
3) N or Si
4) O or Cl? opposing trends ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ ↑
1s 2s 2p 1s 2s 2p
↑↓ ↑↓ ↑ ↑↓ ↑↓ ↑↓ ↑ ↑
1s 2s 2p 1s 2s 2p
Which is easier to remove an electron
from B or Be? Why?
Which is easier to remove an electron
from N or O? Why?
Irregularities in the - First Ionization Energy Trends
Trends in Successive
↑↓ ↑↓ ↑↓ ↑
1s 2s 2p
1s 2s 2p
To ionize Be you must break up a full sublevel, cost extra energy
• removal of each successive
↑↓ ↑↓ ↑ ↑↓ ↑↓ electron costs more energy
1s 2s 2p 1s 2s 2p
– shrinkage in size due to having
When you ionize B you get a full sublevel, costs less energy more protons than electrons
B, Al, Ga, etc.: their ionization energies are slightly less than the ionization energy of the – outer electrons closer to the
element preceding them in their period.
• Before ionization ns 2np 1. nucleus, therefore harder to
• After ionization is ns 2. Higher energy ⇒ smaller radius.
↑↓ ↑↓ ↑ ↑ ↑ ↑↓ ↑↓ ↑ ↑
• regular increase in energy for
1s 2s 2p 1s 2s 2p
To ionize N you must break up a half-full sublevel, cost extra energy
each successive valence
↑↓ ↑↓ ↑↓ ↑ ↑ ↑↓ ↑↓ ↑ ↑ ↑
O O+ • large increase in energy
1s 2s 2p 1s 2s 2p
When you ionize O you get a half-full sublevel, costs less energy
when start removing core
Group 6A elements. electrons
• Before ionization ns2np4.
• After ionization ns2np3 where each p electron in different orbital (Hund’s
HIGHER IONIZATION ENERGIES Electron Affinities
• Electron affinity is the opposite of ionization
• Electron affinity is the energy change when a
gaseous atom gains an electron to form a
Cl(g) + e- → Cl-(g)
• Electron affinity can either be exothermic (as
the above example) or endothermic:
• more energy released (more -); the larger the
generally increases across period
becomes more negative from left to right
lowest EA in period = alkali earth metal or noble gas
highest EA in period = halogen
Magnetic Properties of
Transition Metal Atoms & Ions
• electron configurations that result in unpaired
electrons mean that the atom or ion will have a
net magnetic field – this is called
– will be attracted to a magnetic field
• electron configurations that result in all paired
electrons mean that the atom or ion will have
no magnetic field – this is called diamagnetism
– slightly repelled by a magnetic field
• both Zn atoms and Zn2+ ions are diamagnetic,
showing that the two 4s electrons are lost
before the 3d
– Zn atoms [Ar]4s23d10
The added electron in Cl is placed in the 3p orbital to form the stable 3p6 – Zn2+ ions [Ar]4s03d10
electron configuration. 54
Melting Points and Boiling Points
Example 8.6 – Write the Electron Configuration and •Trends in melting Points and boiling points can be
Determine whether the Fe atom and Fe3+ ion are used as a measure of the attractive forces between
Paramagnetic or Diamagnetic atoms or molecules.
•Within the halogens (group 17 or VIIA) melting
• Fe Z = 26 points and boiling points increase so that at room
temperature fluorine and chlorine are gases,
• previous noble gas = Ar bromine is a liquid, and iodine is a solid as you go
– 18 electrons down this periodic group.
•This indicates that the intermolecular forces
• Fe atom = [Ar]4s23d6 become stronger going down a group.
• unpaired electrons 4s 3d Melting Point Trends in Period
• paramagnetic •In the second period, melting points increase, going
from left to right across the period for the first four
• Fe3+ion = [Ar]4s03d5 elements.
• unpaired electrons • Melting points then decrease drastically for nitrogen,
• paramagnetic oxygen, and fluorine, which are all diatomic molecules.
• The lowest melting point is for neon, which is
Melting Points of Elements
Trends in Density
•Densities of elements increase in a group
as atomic number increases.
•In periods, going from left to right,
densities increase, then decrease.
•Elements with the greatest densities are
at the center of period 6.