Periodicity Flashcards
Possible questions?
Who noticed the patterns of periodicity?
Who arranged the elements?
In the mid-1800’s, Dmitri Mendeleev, a Russian chemist, noticed a repeating pattern of chemical properties in the elements that were known at the time. Mendeleev arranged the elements in order of increasing atomic mass. Later the elements were rearranged in order of increasing atomic number.
The repeating pattern of chemical properties in the periodic table allows us to predict and learn the properties of families of elements rather than learning the properties of all 109 elements!
Why is the s-block and p-block given its name?
Signifies what element belongs to each orbital
S = sigma
P = pi
Which side of the periodic table do we catabolise the metals and which the non-metals?
Non-metals on the right
Metals are generally found on the left
Where is the dividing line between the metals and non-metals? What are these elements called?
In P block
Metalloids
Periods (rows)
All elements within a period;
Have the same number of electron shells. But widely differing characteristics
Groups (columns)
All elements within a group;
Have similar chemical properties and they have the same number of electrons in their valence shell (outer shell)
What’s the general valence electronic configurations for groups 1-8?
Group 1 - ns ^1
Group 2. ns2
Group 3. ns 2 np1
Group 4. ns2 np3
Group 5. ns2 np3
Group 6. ns2 np4
Group 7. ns2 np5
Group 8. ns2 np6
Elemental types
Element & description and elemental bonding
Group 8 elements; noble gases -
Atomic elements; weak Van der Waals
Group 1,2,3 -
Metallic element; metallic bonding
Group 4; carbon and silicon -
Giant covalent three dimentional lattice; strong covalent bond between atoms
Groups 5,6 and 7 -
Simple molecular; covalent bonding between atoms, Van der Waals between molecules
General periodic trends:
Across a row; ionisation energy and electronegativity
Ionisation energy and electronegativity generally increase across a row. This is because of the progressive increase in nuclear charge which outweighs any shielding effect.
I.e. additional electrons enter the same shell and the increase in nuclear charge tends to pull them in closer.
General periodic trends:
Down a group; ionisation energy and electronegativity
Ionisation energy and electronegativity decreases down a group. This is because the number of electronic shells increases which means an increase in shielding that outweighs any increase in nuclear charge. Therefore the effective nuclear charge decreases.
General periodic trends:
Across a row; melting point
Row 2 increase or decrease between what elements?
Row 3 same question?
Melting point is related to the strength of bonding
Row 2: The melting increases from Lithium to Carbon then suddenly decreases to Nitrogen. Nitrogen (N2) oxygen (02) fluorine (Fa) and Neon (Ne) all have similar low melting points.
Row 3; The melting point increases from Na to Si then suddenly drops to phosphorus. The melting points of phosphorus (Pa), sulphur (Se), chlorine (Clz) and argon (Ar) are all low however the melting points decrease in the order:
Sulphur -> phosphorus -> chlorine -> argon
Why does the melting point decrease form carbon to nitrogen?
Carbon is a giant covalent structure while nitrogen is a simple milecular with a weak Van der Waals forces that are easier to break
Why are melting temperatures of N2, O2 and F2 all low?
They all have weak Van der Waals forces that are easy to break and take less energy to break it
