Chapter 7 Periodicity Flashcards
How did Mendeleev organise his periodic table?
arranged using atomic mass and grouped elements that had similar properties.
–> If grouped properties did not fit, Mendeleev swapped elements around and left gaps, assuming the atomic mass measurements were incorrect
How is the modern Periodic table organised?
114 elements organised into 7 horizontal periods and 18 vertical groups
–> position is linked to physical and chemical properties
–> arranged in increasing atomic number
Groups and periods
vertical columns: groups which all have atoms with the same number of outer-shell electrons and similar properties
horizontal rows: periods
–> no on period gives number of highest energy electron shell in an element’s atom
Periodicity & properties
repeating trends in properties of elements
–> electron configuration
–> ionisation energy
–> structure
–> melting point
Trends across a period
Across period 2, 2s sub-shell fills with two electrons, then 2p sub-shell with 6 electrons e.g. Be is (He)2s^2 and Ne is (He)2s^2,2p^6
Across period 3 same pattern of filling for 3s and 3p sub-shells
Across period 4, although 3d sub-shell is involved, highest shell number=4 so from n=4, only 4s and 4p sub-shells are occupied
Trends down a group
Elements have the same number of electrons in each sub-shell (can be divided into 4 distinct blocks)
–> Group 1&2 are s-block
–> Group 3-12 are d block
–>Group 13-18 are p-block
Names and numbers of groups
1- Alkali metals
2- alkaline earth metals
3-12 Transition elements
15 pnictogens (N,P,As,Sb,Bi)
16 chalcogens (O,S,Se,Te,Po)
17 Halogens
18 Noble gases
Ionisation energy
How easily an atom loses electrons to form positive ions
First ionisation energy
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
Na(g) –> Na+(g) + e- //. X(g) –> X+(g) + e-
Successive ionisation of Chlorine
Cl(g) –> Cl+(g) +e- It is not diatomic because ionisation energy is from ONE atom at a time
IE2: Cl+(g) –> Cl2+(g) + e- (only one electron is removed per ionisation energy)
Factors that affect ionisation energy
Atomic radius, nuclear charge, electron shielding
Atomic radius IE
The greater the distance between the nucleus and outer electrons, the less nuclear attraction
–> force of attraction falls of sharply with increasing distance so atomic radius has a large effect
Nuclear Charge IE
More protons in nucleus= greater attraction between nucleus and outer electrons
Electron Shielding IE
Inner-shell electrons repel outer-shell electrons and reduces the attraction between the nucleus and outer electrons
Why is the second ionisation energy greater than the first IE?
Once an electron is lost, the single electron is pulled in closer as there are now more protons than electrons.
–> causes a stronger attraction on the remaining electron
Second ionisation energy definition
the 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
What does a large difference in ionisation energy mean?
This means a different shell is now losing electrons
–> much closer to the nucleus so stronger attraction
–> more energy is required to lose the electrons)
Trend in first ionisation energies
As you go across each period, there is a general increase in first IE
–> higher nuclear charge= more protons in nucleus= more attraction between nucleus and outermost electron
A sharp decrease between end of one period and start of another ( He->Li/ Ne->Na/ Ar->K)
As you go down a group, why does first IE decrease?
atomic radius increases
more inner shells in total so shielding increases (More repulsion)
nuclear attraction on outer electrons decreases
first ionisation energy decreases
Trends in First IE across a period
Across a period, the nuclear charge increase as there are more protons
–> same shell means similar shielding levels
–> nuclear attraction increases and atomic radius decreases
–> first ionisation energy increases
Why does Beryllium have a higher ionisation energy than Boron
Beryllium has no 2p subshell but Boron has one electron in the 2p sub-shell.
2p sub-shell in boron has a higher energy than 2s sub-shell in beryllium.
–> Therefore, in boron, the 2p electron is easier to remove than one of the 2s electrons in beryllium (lower ionisation IE)
Why does oxygen have a lower ionisation energy than nitrogen
Due to electron pairing in p-orbitals of 2p sub-shell
–> in N & O, the highest energy electron are in a 2p sub-shell
–> in oxygen, paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electron from an oxygen atom than a nitrogen atom
–> ionisation energy of oxygen is less than nitrogen
Nitrogen and Oxygen electron configuration
3 2p electrons: 2px1 2py1 2pz1
–> one electron in each 2p orbital
–> spins are at right angles - equal repulsion as far apart as possible
4 2p electrons: 2px2 2py1 2pz1
–> two electrons in one 2p orbital
–> 2p electrons start to pair
–> paired electrons repel
Metallic bonding and structure
in solid metal structure, each atom has donated its negative outer-shell electrons to a shared sea of electrons. (delocalised)
–> positive ions ( cations) left behind consist of nucleus and inner electron shells of metal atoms
Metabolic atoms: strong electrostatic forces of attraction between cations and delocalised electrons
–> fixed cation with mobile delocalised electrons
(giant metallic lattice)
Properties of metals
strong metallic bonds
high electrical conductivity
high melting and boiling points
Electrical conductivity
metals conduct electricity in solid and liquid states
–> voltage applied across metal= delocalised electrons can move through the structure, carrying charge
–> ionic compounds cannot carry charge when solid
Melting and boiling points
depends on strength of metallic bonds in a giant metallic lattice
–> most require high temps to provide large amount of energy needed to overcome strong electrostatic attraction between cations and electrons
–> strong attraction results in most metals having high melting and boiling points
Solubility
metals do not dissolve (can interact in a polar solvent)
Giant covalent structures
strong network of covalent bonds
–> carbon and silicon both have 4 electrons
–> can form 4 covalent bonds to other carbon or silicon atoms (results in tetrahedral shape)
–> 109.5 bond angle
Properties of giant covalent structures
melting and boiling points: high due to strong covalent bonds (need lots of energy to break)
solubility: insoluble as covalent bonds are way too strong to be broken by interaction with solvents
electrical conductivity: non conductors, except graphene and graphite
- in diamond all 4 electrons are used
periodic trend in melting points
- Melting point increase from group 1-14 (giant structures)
- sharp decrease in melting point from group 14-15 (simple molecules)
- melting points are comparatively low from group 15-18
Bond strength
Giant covalent> Giant metallic> simple molecular
simple molecules: weak London forces between molecules
