PERIODIC TABLE TRENDS
1. Who compiled the periodic table? What is the periodic law?
2. Define periods (rows) and groups (families) in the periodic table. How are the elements
arranged in the periodic table?
3. Name the elements in each group – alkali metals, alkaline earth metals, metalloids, halogens
and noble gases. Name the elements in the period 1 and period 2.
4. From left to right and top to bottom, the directions of increase in the periodic table for
various properties is listed below:
Electronegativity
Electronegativity (EN) is the relative ability of a bonded atom to attract shared electrons. EN increases
across a period from left to right and from bottom to top of a group.
EN is maximum for Fluorine (4.0) and minimum for Francium (0.7). EN differences between atoms
account for the bond energy and polarity of ionic and covalent bonds in molecules.
Atoms arranged with increasing EN values: O > Al > Co > Sr > Fr
ΔEN 0.0 0.4 1.7 3.3
Nonpolar Mostly Polar Mostly
covalent covalent covalent ionic
How to indicate polarity of a bond:
F N
First, determine the EN values of the atoms and point the arrow towards the more negative atom with
the higher EN. F N
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Non-metallic properties
Metallic properties
, Non-metals are on the right side (p –group elements) of the periodic table, while metals are on the left
side (s group elements) and the transition metals (d group elements) in the middle. The metalloids (B,
Al, Si, Ge, As, Sb, Te, Po,At) divide the metals from the non-metals. The increase in metallic behavior is
consistent with an increase in atomic size and decrease in IE.
Metal atoms bond by delocalizing their valence electrons forming an orderly array of nuclei and
electrons in the electron-sea model. These metallic bonds are strong and determine metallic properties
like boiling points (very high) and melting points (moderately high), good electrical and thermal
cconductivity (due to mobile electrons). Metallic properties increase down a group and decrease across
a period from left to right.
Nonmetals usually form covalent compounds (example CO, CO2, O2, O3) or form ionic bonds with metals
(Group 7) and do not possess properties characteristic of metals due to localized electron pairs in
covalent and ionic bonds. Their properties vary based on electron configuration, EN and type of bonds
they form.
Transition metals form paramagnetic substances when unpaired electrons are present in their
configuration. N is a gaseous nonmetal, while Bi is a metal.
For most elements, the A-Group number is the highest oxidation number (positive). For metals and
metalloids, the A-group number minus 8 is the lowest oxidation number (negative).
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Acid –forming properties
Base-forming properties
Cations from atoms (Group I and II) form bases with water (example NaOH, Ca(OH)2), while anions from
atoms (Group VII) or polyatomic anions that accept H+ form acids (examples HCl, HNO3)
Within Groups 1 and 2, basic properties of oxides of metals increase down a group and decrease across
a period.
Within Group 7, HF is the weakest acid while HI is the strongest. This is because bond energy is highest
for HF with the highest ΔEN making it more difficult to dissociate the H+ ions.
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1. Who compiled the periodic table? What is the periodic law?
2. Define periods (rows) and groups (families) in the periodic table. How are the elements
arranged in the periodic table?
3. Name the elements in each group – alkali metals, alkaline earth metals, metalloids, halogens
and noble gases. Name the elements in the period 1 and period 2.
4. From left to right and top to bottom, the directions of increase in the periodic table for
various properties is listed below:
Electronegativity
Electronegativity (EN) is the relative ability of a bonded atom to attract shared electrons. EN increases
across a period from left to right and from bottom to top of a group.
EN is maximum for Fluorine (4.0) and minimum for Francium (0.7). EN differences between atoms
account for the bond energy and polarity of ionic and covalent bonds in molecules.
Atoms arranged with increasing EN values: O > Al > Co > Sr > Fr
ΔEN 0.0 0.4 1.7 3.3
Nonpolar Mostly Polar Mostly
covalent covalent covalent ionic
How to indicate polarity of a bond:
F N
First, determine the EN values of the atoms and point the arrow towards the more negative atom with
the higher EN. F N
____________________________________________________________________________________
Non-metallic properties
Metallic properties
, Non-metals are on the right side (p –group elements) of the periodic table, while metals are on the left
side (s group elements) and the transition metals (d group elements) in the middle. The metalloids (B,
Al, Si, Ge, As, Sb, Te, Po,At) divide the metals from the non-metals. The increase in metallic behavior is
consistent with an increase in atomic size and decrease in IE.
Metal atoms bond by delocalizing their valence electrons forming an orderly array of nuclei and
electrons in the electron-sea model. These metallic bonds are strong and determine metallic properties
like boiling points (very high) and melting points (moderately high), good electrical and thermal
cconductivity (due to mobile electrons). Metallic properties increase down a group and decrease across
a period from left to right.
Nonmetals usually form covalent compounds (example CO, CO2, O2, O3) or form ionic bonds with metals
(Group 7) and do not possess properties characteristic of metals due to localized electron pairs in
covalent and ionic bonds. Their properties vary based on electron configuration, EN and type of bonds
they form.
Transition metals form paramagnetic substances when unpaired electrons are present in their
configuration. N is a gaseous nonmetal, while Bi is a metal.
For most elements, the A-Group number is the highest oxidation number (positive). For metals and
metalloids, the A-group number minus 8 is the lowest oxidation number (negative).
_____________________________________________________________________________________
Acid –forming properties
Base-forming properties
Cations from atoms (Group I and II) form bases with water (example NaOH, Ca(OH)2), while anions from
atoms (Group VII) or polyatomic anions that accept H+ form acids (examples HCl, HNO3)
Within Groups 1 and 2, basic properties of oxides of metals increase down a group and decrease across
a period.
Within Group 7, HF is the weakest acid while HI is the strongest. This is because bond energy is highest
for HF with the highest ΔEN making it more difficult to dissociate the H+ ions.
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