Chapter 7- Periodicity Flashcards
how did Mendeleev order the periodic table
by atomic mass - he left gaps where required so that elements could be in groups of similar chemical properties
what happens to atomic number moving left to right
it increases
what are groups
elements with the same chemical properties because they have the same number of outer electrons
what are periods
horizontal rows, period number relates to the highest energy level held by electrons in that atom
what does periodicity include
- electron configuration
- ionisation energy
- structure
- melting/boiling points
what are the blocks of the periodic table
+ s-block = groups 1 and 2 –> outermost electrons in the s-subshell
+ d-block = transition metals - outer electrons in d subshell
+ p-block = right side = outermost electrons in the P subshell
what is the definition of ionisation energy
“the energy required to remove one electron from each atom in one mole of gaseous atoms”
what are the 3 factors affecting ionisation energy
nuclear charge, atomic radius, shielding
MENTION ALL 3 IN ANSWERS
explanation of how nuclear charge affects ionisation energy
the greater the atomic number/nuclear charge, the more protons there are; this exerts a greater pull on the outer electrons / more electrostatic forces, so more energy is required to remove the outer electron
explanation of how atomic radius affects ionisation energy
a greater atomic radius leaves the outer electrons further from the nucleus, across a period atomic radius decreases due to increased nuclear charge, a greater distance gives less electrostatic force so less ionisation energy
explanation of how shielding affects ionisation energy
greater shielding reduces the attraction from the nucleus to the outer electrons, this decreases ionisation energy. so more electron shells gives a lower ionisation energy
what happens to ionisation energy across a period
- nuclear charge increases across a period so electrostatic attraction increases
- atomic radius decreases across a period due to greater nuclear charge. this means the outer electrons are closer to the nucleus so attraction increases
- shielding remains constant
- overall ionisation energy increases
what are the two exceptions to the general ionisation energy rule across a period
- where a higher energy subshell is being filled/electron being removed from; this is further away so ionisation energy decreases
- paired electrons in P subshell are easier to remove due to the repulsion between electrons (only for first paired p electron)
explanation of ionisation energy down a group
- shielding increases as you move down a group; this decreases attraction to outer electrons
- atomic radius increases down a group; this decreases electrostatic attraction to outer electrons
- nuclear charge increases down a group but this is not sufficient to overcome the changes due to the other factors
- overall ionisation energy decreases
how does energy change with successive ionisations
- with each successive ionisation, more energy is required as you are then removing an electron from a positive ion
- this increases the effective nuclear charge
what does a ‘big jump’ indicate in successive ionisation energies
that a new shell has been entered
what is metallic bonding
metallic bonding occurs when metal atoms ‘donate’ their outer electrons to form positive ions in a sea of delocalised electrons.
the cations are fixed, the electrons are not
what are the properties of metallic substances
most are:
- strong due to large electrostatic forces
- high electrical conductivity- the electrons are free to move a charge
- high mpt and bpt; ‘a large amount of energy is required to overcome the strong electrostatic forces in metallic bonding’
- insoluble
what is one factor of a metal’s bpt
- the ion that it forms, 1+, 2+ etc.
giant covalent structures examples
boron, carbon, silicon
what are some properties of giant covalent structures
- very high mpt and bpt because a large amount of energy is required to break covalent bonds
- almost completely insoluble
- don’t conduct electricity (other than graphite and graphene)
why do group 2 metals have a higher bpt and mpt than group 1 metals
they donate 2 electrons not 1 so there are greater electrostatic forces
mpt and bpt trends across periods 2 and 3
- increases G1 to G2 and aluminium is higher again
- bpts max at carbon and silicon because covalently bonded
- tends to decrease after this because only intermolecular forces need to be broken
what is an exception of bpts decreasing after carbon and silcon
sulphur has a higher bpt than phosphorus because it is S8 but only P4 so there are greater London forces
Why is the hydrogen bonding stronger in water than in ammonia (2)
- water has two lone pairs so it can form twice as many hydrogen bonds
- O-H is a greater dipole than N-H so it has a greater electrostatic attraction
if there is a large jump between the Xth and (X+1)th ionisation energies, which position in the period does the element sit
the Xth position
- if there is a large jump at 5–> 6 for example then the element must have 5 in its outer shell
ideally how should you write the electron configurations
in order of numbers so 3d before 4s even though 4s fills first, doesn’t really matter though
which things should always be mentioned when comparing the atomic radius of elements in the same period
- nuclear charge and therefore electrostatic attraction
- SAME NUMBER OF SHELLS
