C7 Periodicity Flashcards
Define periodicity
A repeating trend in chemical and physical properties of the elements across each period of the periodic table.
How is the periodic table arranged now?
Arranged in increasing atomic number
Groups (vertical columns)
Periods (horizontal rows)
What’s the trend in electron configuration across period 2?
2s sub-shell fill with 2 electrons
Followed by 2p sub-shell with 6 electrons
How was the periodic table arranged back then and by who?
Mendeleev
Order of atomic mass (didn’t know about sub atomic particles)
Lined up elements in groups with similar properties
Left gaps (and predicted their properties)
What’s the trend in electron configuration across period 3
3s sub shell fill with 2 electrons
Then 2p sub shell with 6 electrons
What’s the history of the atoms?
early 1800s, Dalton, tiny hard spheres and couldn’t be divided or splitted
end 1800s, Thomson (discovered tiny neg charged particle ‘electron’) plum pudding model, electrons embedded in ball of pos charge.
10 yrs later, Geiger & Marsden, experiment with alpha particles fired at thin piece of gold (some alpha particles repel, must be tiny spot of pos charge in centre of atom), Rutherford - nuclear model (electrons orbit nucleus which contains protons)
Bohr, suggested electrons orbit nucleus in shells
1932, Chadwick, discovered neutrons (uncharged particles) in nucleus
Why is the 1st ionisation energy of oxygen less than 1st ionisation energy of nitrogen
because of electron pairing)
in N & O highest energy e-s in 2p sub shell
in O, paired e-s in 1 of 2p orbitals repel each other, so its easier to remove e- form O than N
What’s the trend in 1st ionisation energy down a group?
- atomic radius increases
- more inner shells (shielding increases)
- nuclear attraction on outer electrons decreases
- 1st IE decrease
What’s the trend in 1st ionisation energy across a period?
- nuclear charge increases
- similar shielding (same shell)
- Nuclear attraction increases
- atomic radius decreases
- 1st IE increases
why is the 1st IE of boron less than the 1st IE of berylium?
(because of 2s and 2p sub shells)
- fall in 1st IE from Be to B marks filling of 2p sub shell
- 2p sub shell in B has higher energy than 2s sub shell in Be
- in B 2p e- easier to remove than 1 of 2s e-s in Be
so 1st IE of B less than Be
define first ionisation energy
energy required to remove 1 electron from each atom in 1 mole of gaseous atoms of an element to form 1 mole of gaseous 1+ ions.
Na(g) -> Na+ (g) + e-
What factors affect ionisation energy?
Atomic radius
Nuclear charge
Electron shielding
Define 2nd ionisation energy
energy required to remove 1 electron form each ion in 1 mole of gaseous 1+ ions of an element to form 1 mole of gaseous 2+ ions.
He+(g) -> He 2+(g) + e-
why is the 2nd ionisation energy of helium greater than the 1st?
- 2 protons attract 2 electrons in 1s sub shell
- 1st electron lost, 2nd electron pulled closer to nucleus
- greater nuclear attraction on 2nd electron increases
- more IE needed to remove 2nd electron
Define metallic bonding
Strong electrostatic attraction between positive ions and delocalised electrons
why can metals conduct electricity in solid and liquid states?
delocalised electrons can move through giant metallic lattice and carry charge
what is electrical conductivity of giant covalent lattices?
non conductors of electricity
except graphene and graphite (forms of C)
what is the structure and bonding of B, C and Si?
Giant covalent lattice
many billions of atoms are held together by a network of strong covalent bonds
what is the solubility of Giant Covalent Lattices?
- insoluble in nearly all solvents
- CBs holding atoms together in lattice are too strong to be broken by interaction with solvents
Do giant covalent lattices have high/low MPs & BPs?
high MPs & BPs
high temperatures required to provide large quantity of energy required to break strong CBs
why do metals have high melting and boiling points?
high temperatures are necessary to provide the large quantity of energy required to overcome strong metallic bonds (electrostatic attraction between cations and delocalised electrons)
what is the solubility of metals?
metals do not dissolve
( there maybe some interaction between polar solvents and charges in a metallic lattice )
Carbon and Silicon
group 14 (4)
4 electrons in their outer shell
C (in diamond form) and Si use these 4 electrons to form CBs to other C/Si atoms
what is the periodic trend in MPs across period 2 & 3?
MP:
- increase from group 1 to 4
- sharp decrease from group 4 to 5 marks change from giant to simple molecular structures
- are comparatively low from group 5 to 0 giant
metallic structure => giant covalent structure => simple molecular structure
Graphene and graphite
- 3 of 4 outer shell e-s used in CB
- remaining e- released into pool of delocalised e-s shred by all atoms in the structure
- good electrical conductors
- giant covalent structures of C
- planar hexagonal layers
- bond angle = 120* by electron-pair repulsiongraphene =
- single layer of graphite
- hexagonally arranged C atoms linked by strong CBs
graphite =
- parallel layers of hexagonally arranged C atoms
- layers bonded by weak LFs
- 3 of 4 outer shell e-s used in CB
- remaining e- released into pool of delocalised e-s shred by all atoms in the structure
Metallic bonding and structure
electrons delocalise and can move
cations are fixed in place and maintain structure and shape of the metal
what is the name of the structure that contains metallic bonding?
Giant metallic lattice
what is the name given to elements near to the metal/non-metal divide on the periodic table?
semi metals or metalloids
e.g. B, Si, Ge, As, and Sb
what are the properties of most metals?
- strong MBs
- high electrical conductivity
- high MPs & BPs
why can’t silicon and carbon (diamond) conduct electricity?
in both structures, all 4 outer shell electrons are involved in CB, so none are available for conducting electricity.
what do many non-metallic elements exist as?
simple covalently bonded molecules
in solid state, they form simple molecular lattice structures which are held together by weak intermolecular forces - theses structures have low MPs and BPs
what is the bond angle and structure of (carbon) Diamond?
- tetrahedral arrangement
- bond angle = 109.5*
because of electron pair repulsion
What change happens across each period
Elements change from metals to non-metals
How can the electron configuration be written in short
The noble gas before the element is used to abbreviate
E.g. Li —> 1s^2 2s^2 ; Li —> [He]2s^1
Write an equation for the 1st ionisation energy of magnesium
Mg(g) —> Mg^+ (g) + e-
What factors affect ionisation energy
Atomic radius
Nuclear charge
Electron shielding
Why does 1st ionisation energy decrease between group 2 & 3
In group 3 the outermost electrons are in p-orbitals whereas in group 2 they’re in s-orbital, so the electrons are easier to be removed
Why does 1st ionisation energy decrease between group 5 to 6
Due to group 5 electrons in p-orbital are single electrons and in group 6 the outermost electrons are spin paired, with some repulsion.
Therefore, electrons are slightly easier to remove
Why does 1st ionisation energy decrease between the end of one period and the start of next
There’s an increase in atomic radius
Increases in electron shielding
How are the elements arranged in a periodic table
They’re arranged in order of increasing atomic numbers
What is a period on a periodic table
The horizontal rows in the periodic table
What is a group on a periodic table
The vertical columns
What is meant by periodicity
The repeating trends in chemical and physical properties
What change happens across each period
Elements change from metals to non-metals
How can the electron configuration be written in short
The noble gas before the element is used to abbreviate
E.g. Li —> 1s^2 2s^1 ; Li —> [He]2s^1
Define first ionisation energy
The energy required to remove 1 electron from each atom in 1 mole of the gaseous elements to from 1 mole of gaseous 1+ ions
Write an equation for the first ionisation energy of magnesium
Mg(g) —> Mg+(g) + e-
What are the factors that affect ionisation energy
Atomic radius
Nuclear charge
Electron shielding or screening
Why does 1st ionisation energy decrease between group 2 & 3
Decrease between 2 to 3 because in group 3 the outermost electrons are in p orbitals whereas in group 2 they are in s orbital so the electrons are easier to be removed
Why does 1st ionisation energy decrease between group 5 to 6
The decrease between 5 to 6 is due to the group 5 electrons in p orbitals whereas are single electrons and in group 6 the outermost electrons are spin paired, with some repulsion.
Therefore, the electrons are slightly easier to remove
Does 1st ionisation increase or decrease between the end of 1 period and the start of next? Why?
Decrease
There is increase in atomic radius
Increase in electron shielding
Does 1st ionisation increase of decrease down a group? Why?
Decrease
Sheilding increases —> weaker attraction
Atomic radius increases —> distance between outer electrons and nucleus increases —> weaker attraction
Increase in number of protons is outweighed by increase in distance and sheilding
What are the properties of giant metallic lattices
High MP & BP
Good electrical conductors
Malleability
Ductility
What is a ductile metal
The metal can be made stretched
E.g. can be made into wires
What is a malleable metal
The metal can be shaped into different forms
Describe the structure, forces and boning In every element across period 2
Li & Be —> giant metallic lattice, strong attraction between pos ions and delocalised electrons, metallic bonding
B & C —> giant covalent lattice, strong forces between atoms, covalent bonding
N2, O2, F2, Ne —> simple molecular, weak IMFs between molecules, covalent bonding within molecules and IMFs between molecules
Describe the structure, forces and bonding in every element across period 3
Na, Mg, Al —> giant metallic, strong attraction between pos ions and delocalised electrons, metallic bonding
Si —> giant covalent, strong forces between atoms, covalent
P4, S8, Cl2, Ar —> simple molecular, weak IMFs between molecules, covalent bonding within molecules and IMFs between molecules
