Chapter 7 - Periodicity Flashcards
who came up with the periodic table and what did they do?
- Mendeleev - arranged in increasing atomic mass
- lined up elements in groups w/ similar properties
- if properties didn’t fit he would swap them e.g. I & Te
- left gaps for elements not discovered yet
how is the periodic table arranged
- increasing atomic no.
- metals vs non-metals - staircase under boron
- groups - similar chemical properties and same no. of electrons in outer shell
- horizontal rows - periods
metalloids
- B
- Si
- Ge
- As
- Sb
- Te
- At
periodicity
repeating trend in properties of elements across periods
periodicity examples
- electron configuration
- ionisation energy
- structure
- melting points
trends across periods (electron config.)
s-sub shell fills then p sub-shell
blocks of periodic table
- corresponds to highest energy shell:
s-block –> grp 1&2
p-block –> grp 13-18
d-block –> grp 3-12
f-block –> lanthanides & acintides
what’s ionisation energy
- measures how easily an atom loses it’s electrons to form positive ions
first ionisation energy definition
energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous +1 ions
factors affecting ionisation energy
- atomic radius - as it increases. attraction decreases
- nuclear charge - more protons they are, the greater the attraction
- electron shielding - inner-shells repel - larger the atom, the more shielding, the less attraction
second ionisation energy
energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous +2 ions
trend in successive ionisation energy
- it increases each time as each time an electron is removed, nuclear charge attracted to less electrons
- large increase suggests it’s moved to inner shell - decrease in radius
trend in ionisation energy down a group
decreases, radius and shielding increases, attraction decreases so ionisation energy does
trend in ionisation energy across a period
nuclear charge increases as well as attraction, atomic radius decreases so ionisation energy increases
why is there a fall in ionisation energy between Be & B
Be fills the 2s sub-shell, B fills 2s & one 2p, this is more unstable so it loses an electron more easily, so lower ionisation energy
why is there a fall in ionisation energy between N & O
N has one electron in each 2p sub-shell
O has one 2p sub-shell w/ 2 electrons, the other 2 having one electron, this is more unstable so it loses an electron more easily, so lower ionisation energy
Metallic bonding & structure
electrostatic attraction between the cations and delocalised electrons
- cations in fixed position, delocalised electrons are free to move around - forms giant metallic lattice
properties of metals
- electrical conductivity - due to delocalised electrons
- strong metallic bonds - more delocalised electrons, stronger the bonds
- high m.p/b.p - depends on strength of bonds, high due to lots of energy needed to overcome strong electrostatic attraction between electrons and cations
- are not soluble
Giant covalent structures
many atoms hold together by a network of strong covalent bonds
properties of giant covalent structures
- high m.p/b.p - lots of energy needed to overcome strong covalent bonds
- insoluble as bonds are too strong to be broken by interactions with lattices
- don’t usually conduct electricity - no delocalised electrons
graphite
- parallel layers of hexagonally attracted C atoms, have 3 bonds, so delocalised electrons
- weak intermolecular forces between layers
- can conduct electricity
graphene
- one layer of graphite
- can also conduct electricity
melting point trends across periods
- increases as metallic bonds gets stronger
- continues to increase as it reaches giant covalent structures
- then has a large decrease when gets to simple molecules
structure trends across periods
- first giant metallic
- then giant covalent
- then simple molecular
