Chapter 7 periodicity Flashcards
Describe the periodic table Mendeleev created
- arranged in order of atomic mass
- He lines elements in groups with similar properties
- if the elements did not fit the similar properties he would switch them and leave gasp
Describe the modern periodic table
- from left to right elements are arranged in order of increasing atom number
- each successive element has atoms with one extra proton
What are groups
- vertical columns
- each element in group has atoms with the same number of outer shell electrons
- similar properties
What are periods
- horizontal rows
- the highest energy electron shell in an element’s atom
What is periodicity
the repeating trend in properties of the elements
What is the s block
highest energy electron in an s orbital
What is an element in the p block
highest energy electron in the p orbital
What is ionisation energy
- how easily an atom loses electrons to form a positive ion
What is the first ionisation energy
- the energy required to remove one electron from each atom in one mole of gasours atoms of an element to form one mole of gaseous 1+ ions
What factors affet ionisation energy
- atomic radius
- nuclear charge
- electron shielding
How does atomic radius affect ionisation energy
- the greater the distance between nucleus and outer electrons with the less nuclear attraction
- the force of attraction falls with increasing distance
How does nuclear charge affect ionisation energy
- the more protons there are in the nucleus of an atom the greater the attraction between nucleus and outer electrons
How does electron shielding affect ionisation energy
- electrons are negativley charges so inner shell electrons repel outer shell electrons
- the repulsion is shielding reducing the attraction betwene nucleus and outer electrons
How do you know how many successive ionisation energies an element has
- same number as electrons
Describe the second ionisation energy in helium
- greater than the first ionisation energy
- two protons attracting two electrons in te 1s sub shell
- after first electron lost single electron puller closer to the nucleus
- nuclear attraction incresases so more ionisation energy needed to remove this second electron
What is the second ionisation energy
- is 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
How do you predict ionisation energies from a table
- see where the largest jump is in ioisation eergies
- the amount of ionisation energies till the jump is the number of electrons in its outer shell
- therfore you can see the group of the element
Describe the successive ionisation energies in a graph
- large jumps are change in shells
- small jumps are change in sub shells
Describe the two key patterns in the first 20 key elements in ionisation energy in a periodic table
- general increase in first ionisation energy across each period (H-> He, LI-> Ne, Na-Ar)
- a sharp decrease in first ionidation energy between the end of one period and the start of the next period (He-> Li, Ne-Na, Ar->K)
Decribe the trend on ionisation energy down a group
- atomic raidus increases
- more inner shells so shieldong increases
- nucelar attraction on outer electrons decreases
- first ionisation energy decreases
Describe the trend of first ionisation energy across a period
- nuclear charge increases
- same shell similar shielding
- nuclear attraction increases
- atomic radius decreases
- first ionisation energy increases
Compare the ionisation energy between beyillum and boron
- the 2p sub shell in boron has a higher energy than the 2s sub shell in beryillium
- ttherefore boron the 2p electron is easieer to remove than one of the 2 electrons in beryllium
- the first ionisation energy of boron is less that the frist ionisation energy of beryllium
Compare the ionisation energy between nitrogen and oxygen
- in nitrogen and oxygen the highest energy electrons are in 2p sub shells
- in pxygen paired elecrons in one of the 2p orbitals repel one another making it easier to remove en electron from an oxygen atom than a nitrogen atom
- therefore the first ionisation energu of oxygen is less than the first ionisation energy of nitrogen
Describe the structure of metals
- in a solid structure each ato has donated its negative outer shell electrons to a shared poll of electrons delocalised throughout whole structure
- positvie ions left behind consist of nucleus and the inner electron shells of metal atoms
- cations are fixed in positions maintaining structure
- delicalised electrons are movile and are able to move throughout the structure
What is metallic bonding
- strong electrostatic attraction between cations and delocalised electrons
What is the main structure of metals
giant metallic lattice
Why can metals conduct electricity
- can conduct in solid and liquid stated
- delocalise electrons can move throught structue carrying charge
Descrube why metals have high melting and boiling points
- large amount of energy needed to overcome the strong electrostatic attraction between cations and electrons
- higher in those whose electrostatic charge is larger because they have a larger charge on atom which makes them smaller and more delocalised electrons
Describe the solubulity of metals
- they do not dissolve
Describe giant covalent structures and what elements have these structures
- boron ,carbon and silicon
- instead of small molecules and intemolecular forces billions of atoms held with strong covalent bonds to from a giant covalent lattice
- carbon and silicon bond to other carbon and silicon atoms forming a tetrahedral structure
Describe the structure of diamond
- tetrahedral arrangment
- bond angles are 109.5 by electron pair repulsion
What are the melting and boiling like of giant covalent structures
- high melting and boiling points
- covalent bonds are strong
- large quantity of energy to break strong covalent bonds
Describe the solubility of giant covalent structures
- insoluble
- covalent bonds holding together atoms in the lattice are far too strong to be broken by interactin with solvents
Describe the electrical conductivity of giant covalent structures
- non conductors of electricity
- except graphene and graphite - they only form 3 covalent bonds so one atom is delocalised which can carry a charge round structure
Describe the trends in metling and boiling point in period 2
- Li-Be - giant metallic lattice - strong metallic bonds between cations and delocalised electrons
- B-C - giant covalent structures -strong covalent bonds between atoms
- N2-Ne - simple molecular structure 0 weak london forces between molecules
Describe the trend in in melting and boiling points in period 3
- Na-Al- - giant metallic lattice - strong metallic bonds between cations and delocalised electrons
- Si- giant covalent structures -strong covalent bonds between atoms
- P4,S8,Cl2,Ar - simple molecular structure 0 weak london forces between molecules
