IB Chemistry: Topic 3/13  Periodicity
3.1: The Periodic Table
*The relationship between electronic structure and chemical properties is one of the key
concepts in chemistry.
*Periodicity is the repeating patterns of ______________ and ______________
properties that are seen at regular intervals in the periodic table as a result of similar
__________ electron configurations. Periodicity is evident when the elements in the
periodic table are arranged in order of increasing __________ number.
* period = _______________ row
* group = ___________ column (“family” of elements e.g. group 1 = alkali metals
share/display similar chemical and physical properties)
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Group 1: ____________ metals
Group 2 :_______________ earth metals
Group 6/16: _______________
Group 7/17: _______________
Group 0/8 or 18: ___________ gases
 valence electrons  number of electrons in _____________ shell (How determine?)
3.2: Periodicity  Physical Properties
PERIODIC TRENDS
1) ATOMIC RADIUS
 impossible to measure directly (e“cloud” model= no sharp boundary.)
 diatomics are used to determine radius then react
with others and determine the radius of others
 radius of atom extends from the nucleus to
approximate outer boundary of electron cloud
The Trends:
a) DOWN A GROUP
Atomic radii ____down a group.
*Explanation:
 The number of energy levels ____ for each successive atom as
we move ________ a group.
b) ACROSS A PERIOD (from left to right)
Atomic radii ____ across a period.
*Explanation:
 The ______________ nuclear charge, Zeff, (the more protons
for the ________ number of energy levels) _____ as we move from
_______ to _________ across the periodic table.*This _________
the electron cloud until the point at which electron _____________
overcome the nuclear attraction and stop the contraction of electron
shells.
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2) IONIC RADII
a) CATIONS – The size of an atom always ____when
converted into a cation since the nucleus is now attracting
___________electrons.
Also, a whole ___________ shell of electrons has been lost.
Furthermore, if Mg is converted into Mg+, for example, the
electron-electron repulsion between Mg’s _____ valance
electrons _____, which results in a ____ in size.
b) ANIONS– The size of an atom always_____ when it is
converted into an anion since the nucleus is now attracting
________ electrons than there are protons AND there is
enhanced electron-electron ______________.
*Almost every common cation is smaller than any anion.
*isoelectronic – entities with the same electron
________________________(e.g. Na+, F-, Ne)
3) FIRST IONIZATION ENERGY(IE)
*Recall: The IE of an atom is the minimum amount of energy
required to remove a mole of ___________ from a mole of
_____________ atoms to from a mole of gaseous ______.
Q(g) + IE  ______(g) + ______
 Removing each subsequent electron requires _______ energy
(e.g. second IE, third IE, etc.) mainly due to a ____in e-/erepulsions.
A significant energy price is paid if the subsequent removal
of an electron is from another principal energy level.
The Trends:
A. DOWN A GROUP
*IE _____down a group.
*Explanation:
 i) As we move down a group, atoms get _________ – each successive atom has an additional _________ level. As a result, the
valence electrons are located progressively further and further away from the _____________. The attractive force of the nucleus for
the valence electrons ____as this distance increases. Therefore, the energy required to remove an electron (i.e. the ionization energy)
_____ as we go down a group.
 ii) ____________ effect: _________ electrons block the attraction of the __________ for the outer _____________ electrons.
Once again, this means that it requires ______ energy to remove an electron as we move down group.
B. ACROSS A PERIOD (from left to right)
*IE____across a period.
*Explanation:
 The increase is a result of increasing effective ___________ charge (Zeff). (i.e. The attraction between the nucleus and the valence
electrons becomes progressively ____________ as we move across a period. As a result, the electrons become more _____________
to remove.
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4) ELECTRON AFFINITY (Eea)
Definition #1  the energy required to detach an __________________ from a singly charged
_________________ ion in the gas phase: X-(g)  X(g) + e*Definition #2 (more common)  the energy released when 1 ________ of ________________ is attached
to 1 mol of _____________________ atoms or molecules in the gas phase: X(g) + e-  X-(g)
*The more negative the electron affinity, the ________________ the attraction of the atom for the electron.
Trends (not well highlighted)
A. Down a group  vary (e.g. alkali metals: becomes less negative down a group)
B. Across period  becomes more _______________ across a period, in general.
5) ELECTRONEGATIVITY (EN)
 EN = a relative (no units) measure of how strongly an atom attracts electrons in a _______________ bond
 Developed by Linus Pauling. (Won Nobel Prize for Chemistry AND Peace)
The Trends: (same as ____ trends)
A. DOWN A GROUP
*EN ____down a group.
*Explanation:
In general, as atomic size, EN ____. As you move down a group, valence electrons are _______ strongly attracted to the nucleus.
In a compound, this means that the nucleus attracts bonding pairs less strongly…which translates into a ____ in EN.
B. ACROSS A PERIOD
*EN ____across a period.
*Explanation:
Once again, as atomic size , EN___. Smaller atoms attract a bonding pair ______ strongly, because the bonding pair is closer to the
____________.
 Why is F the most EN?...highest ______ and ________ size – therefore, the nucleus is closest to the “action”
 Why is Fr the least EN?…lowest ______ and ________size – therefore, the nucleus is farthest from the “action”
6) Periodic Trends in Metallic and Non-Metallic Character
*Metal character ______________________ across a period and ___________________ down a group.
*Metals tend to lose electrons (i.e. get __________________) during chemical reactions.
*Nonmetals tend to gain electrons (i.e. get ___________________) during chemical reactions.
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