Module 3 content Flashcards
(61 cards)
How is the periodic table arranged?
By increasing atomic/proton number
In periods showing trends in chemical/physical properties
Trends across the periods in table
Period 2,3= s and p fill
4 onwards= d shell fills,
Trends down the group
Same number of atoms in outer shell
Atoms with the same number of electrons in each sub shell
Blocks in the periodic table
1-2= s
3-12=d
13-18= p
What is first ionisation energy?
The energy required to remove one electron from each atom in one molecule of gaseous atoms of an element to form one mole of gaseous 1+ ions
Factors affecting ionisation energy
First electron lost= highest energy level/least attraction from nucleus
Atomic radius: Higher=less nuclear attraction=decrease ionisation energy
Nuclear charge: more protons in nucleus= higher nuclear attraction= higher ionisation energy
Electron shielding: inner shell electrons repel outer shell electrons, more electrons in middle=higher shielding effect= lower ionisation energy
What affects the values of successive ionisation energies?
Gets much larger when changes from a further away shell to a closer one
When 1 e- is lost the positive nucleus attracts fewer electrons so are pulled in more tightly
Trends of first ionisation down a group
Decrease down a group:
Atomic radius increases
More inner shells= more shielding
Nuclear attraction decreases
Nuclear charge increase is outweighed by radius
Trends in first ionisation energy across a period
Increase across 1st 3 periods:
Nuclear charge increases
Same shell= similar shielding
Nuclear attraction increases
Atomic radius decreases
Each new period down= another shell= lower energy
Sub shell trends in ionisation energy
Falls when has to start filling a new sub shell
as had higher energy level
2p sub shell has higher energy level so less energy is required to remove than 2s so ionisation energy is less
Easier to remove electrons when some are paired and repel= lower ionisation energy
Metallic bonding definition
Strong electrostatic attraction between cations and delocalised electrons
Structure of a giant metallic lattice
Cations are in a fixed position
Delocalised electrons are mobile and can move throughout structure
Properties of metals
High conductivity= delocalised electrons carry charge and move through substance
Bp= High as need large amounts of energy to overcome electrostatic attraction between cations and delocalised electrons
Not soluble
Giant covalent structures
A network of loads of atoms held together by string covalent bonds
Examples of giant covalent lattices
Graphene= 3/4 e- used in bonding, can conduct, planar hexagonal layers, 120 degrees
Graphite= parallel graphite layers bonded w/ London forces
Diamond/silicon= 4/4 e- used in bonds, tetrahedral, 109.5 degrees
Giant covalent lattice properties
High bp= strong covalent bonds
Insoluble= bonds too strong
Non conductors apart from graphene/graphite because delocalised electron
Periodic trends in melting points and explanation
Across period 2&3:
Increase from Group 1-14 (giant)
Decrease 14-15 (sharp)
Comparatively low from 15-18 (simple)(weak IMFs, London)
Group 2 elements characteristics
Each has 2 electrons in outer shell
In outer s sub shell
Redox–>oxidised for a 2+ ion
Group 2 element is a reducing agent (reduces other species)
Trends in reactivity of Group 2 down the group
+ water= alkaline hydroxide & H2
Reactivity increases down the group and become stronger reducing agents
Metal+acid–> salt+H2
- React by losing electrons
- Ionisation energies decrease down group because of more shielding/radius
Metal oxide reaction w/ water
Form alkaline solution of metal hydroxides
Only slight;y soluble (group 2) so when solution is saturated any more metal and OH- form precipitate
Hydroxides solubility trends
Increases down the group
So the solution contains more OH- and alkaline
pH then increases= higher alkalinity
Group 2 uses
Calcium hydroxide as lime is added to fields to increase pH
Bases are used at antacids (calcium carbonate) to neutralise acids and Mg(OH)2
Trends in bp of halogens
Exist as diatomic molecules at RTP
Boiling point increases down the group
due to more electrons, more London forces, and more energy required to break the IMFs
Halogen definition
Outer shell s2p5 and gains 1 electron in redox to form 1- ion
Is an oxidising agent because oxidises another species