Chapter 7 IMPORANT Flashcards
Mendeleev periodic table then
Arranged in terms of atomic mass (didnt know about suabtomic particles then)
Lined up elements with similar properties
Swapped elements and lef gaps for undiscovered
Predicted properties of missing elements from group trends
Periodic table now
Arranged in periods and groups - first point of reference for chemists everywhere
Arranged in order of increasing atomic number
Groups - atom eith the same number of outer shell electrons and similar properties
Elemenst arranged in periods - number of highest energy electron shell in an elements atom
Periodicity?
Repeating trend in properties across each period - trend from metals to non-metals
1) Electron configuration
2) Ionisation energy
3) Sturcture
4) Melting points
Trend across a period - electron config
Each period starts with an electeon in highest energy subshell
Across period 2 - 2s sub shell fills with 2 electrons first followed by 2p with 6
Same across period 3 but with 3s and 3p sub-shells
Across period 4 - although 3d sub shell involved, only 4s and 4p sub shells are occupied (n=4 is highest shell number)
Trend down a group electron config
Atoms of elements in same group have same number of electrons in each outer sub-shell - this is what gives them their similar chemistry
Blocks?
Elements can be divided corresponding to their higest energy sub-shell to give 4 distinct blocks s, p, d and f
S block
First 2 groups (include helium)
D block
Transition metals (10 groups)
P block
6 groups non-metal
F block
14 groups below at bottom - DISREGARD LANTHANIDES ETC THEY ARE PART OF D BLOCK AS WITH TRANSITION METALS
Two ways of numbering groups
Old numbers - groups 1-7 and then 0 ; based pn s and p blocks - the advantage of the pld numbering is that the group number matches the number of electrons in outer shell
IUPAC - 1-18 groups - s, d and p blocks sequentially ; approved in 1988 but it takes time for old practices to change
Group 1
Alkali metals
Group 2
Alkaline earth metals
Groups 3-12
Transition elements
Group 7
Halogens
Group 0
Noble gases
3 factors affecting ionisation energy
Atomic radius
Nuclear charge
Electron shielding
What is 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
Exacmple of Na first and second ionisation energy
Na (g) -> Na+ (g) + e-
Na+ (g) -> Na2+ (g) + e-
Change from metal to non-metal?
Staircase from top of group 3 down to the bottom of group 7 ; elements newr rhe divide (boron, silicon, germanium, arsenic, antimony) can show in-between properties and are called metalloids
Change from non-metal to metal?
Going down the group - clearest in group 4 with carbon. 92 metals vs 22 non-metals ; nonxmetals are especially importsnt and in particular elements like carbon hydrogen oxygen and nitrogen
At room temperature metals?
All metals except mercury are solids - wide range of properties
Constant property of all metals
ABILITY TO CONDUCT ELECTRICITY - CHARGE MUST BE ABLE TO MOVE WITHIN A RIGID STRUCTURE FOR CONDUCTION TO RAKE PLACE
Metallic bonding
In solid metal structure, each metal arom has turned into cations by donating it’s outer electrons to a sea of delocalised electrons ; metallic bonding is strong electrostatic attraction between nucleus and these electrons. CATIONS FIXED - MAINTAINING SHAPE
ELECTRONS CAN MOVE - MOBILE CARRIERES OF CHARGE
How are they held together?
In a giant metallic lattice
Properties of most metals (giant structure + bonding)
Strong metallic bonds
High electrical conductivity
High melting and boiling points
How do metals conduct electricity?
In solid and liquid states - potential difference is applied then delocalised electrons can move through the structure carrying a charge (towards positive terminal) ; contrast eith ionic compounds which have no mobile charge carriers in solid state
Highest melting point
Tungsten - used in filaments of halogen lamps
Lowest melting point
Mercury + group 1 elements in peridoic table
What does melting point depend on?
Strength of the metallic bonds holding together the lattice - most elements, a lot of energy is needed to overcome these string electrostatic forces of attraction therefore high melting point
Solubility of metals
They do not dissolve - instead interactions lead to reactions takinf place
Simple molecules?
Many non-metals exist as simple covalently bonded molecules ; a simple molecular lattice structure held together by weak intermolecular forces therefore low boiling point
Carbon and solicon and boron?
Billions of atoms held together by a network of strong covalent bonds to form a giant covalent lattice
What is so special about carbon and silicon?
4 electrons in outer shell (group 4) - use this to form 4 covalent bonds to other carbon/silicon atoms resulting in a large complex tetrahedral structure with 109.5 degrees electron psir repulsion
Properties of giant covalent structures?
Insoluble - covalent bonds are too strong to be broken down by any interactions with solvents
High melting and boiling points - high temeprature necesary to beeak the many covalent bonds
Non-conductors of electricity (mostly) - ONLY EXCEPTION IS GRAPHENE AND GRAPHITE ; in carbon and silicon all electrons are used in bonding therefore none available to carry any charge
Graphene and graphite
Form of carbon where only 3 take part in bonding with each atom releasing 1 into a pool of delocalised elecyrons - bond angles of 120 degrees trigonal planar. Hexagonal layers that slide lver each other
Periodic trends in melting points?
Across Period 2 and 3 - melting point increases from 1 to 4 (culminating with giant covalent lattices) then there is a sharp decrease in melting point between group 14 and group 15 with the melting points even lower from 15 to 18
Where is this trend repeated?
Across peripd 4 in s and p block too - sharp decrease in melting point makes a change from giang structures to simple structures ; strong forces to weak forces
Atomic radius
Greater distance between nucleus and outer electrons = less attraction
Nuclear charge
More protons = grewter attraction between psootvie nucleus and outer electrons
Electron shielding
Inner shell electrons repel outer shell ; this repulsion reduces attraction between the nucleus and outer electrons
How many ionisation energies?
An element has as many ionisation energies as electrons
Why is second ionisation energy higher?
Fewer electrons pulled towards the same number of protons therefore larger attraction therefore more ionisation energy will be needed to remove the electron
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
Number of the ionisation energy?
Same as the charge on the ion produced ; second ionisation energy PRODUCES 2+ ion from 1+ ion
What do successive ionisation energies allow us to deduct?
Number of electrons in the outer shell and the group of the element in the peridoic table and thus the identity of an element
Trend in first ionisation energy down a group?
Decreases as nuclesr charge increases but this is outweighed by atomic radius and shielding increasing therefore less attraction thus essier to remove outer electron
Trend in ionisation energy across a period
General increase in first ionisation energy - nuclear charge increased with similar shielding snd a decrease in atomic readius therefore nuclear attraction increases and thus first ionisation energy increases
Sub-shell trends in ionisation energy ; periodicity
Across both period 2 and period 3 the first ionisation energy does fall in two places ; drops occuring in the same place across each period suggesting that there might be a periodic cause - LINKED WITH EXISTENCE OF SUB-SHELLS, ENERGIES AND HOW ORBITALS FILL
Where are two falls across period 2
1) Beryllium to boron
2) Nitrogen to oxygen
Comparing beryllium and boron
Fall in first ionisation energy marks the start of filling the 2p sub-shell. 2p sub+shell has a higher energy level than 2s sub-shell therefore in boron the 2p electron is easier to remove that one of the 2s electrons in beryllium.
Comparing nitrogen and oxygen
Fall is because in oxygen, the elctrons beginnpairing in the 2p sub-shell ; they repel one another making ti easier to remove an electron from an oxygen than a nitrogen atom. Thus first ionisation energy of oxygen is less than the first ionisation energy of nitrogen
Electrons in 2p nitrogen
Equally spaced apart and spins are at right angles with equal repulsions
Electrons in 2p in oxygen
2p electrons start to pair together so the paired electrons repel
