Periodic table(tests) Flashcards
define allotrope
different forms of the same element in the same state
describe features of diamond
-very high melting point, due to 4 strong c-c bonds -extremely hard, c-c bonds act in c-c -non-conductive, all electrons held in tight covalent bonds -good thermal conductor -insoluble, with any solvent
describe features of graphite
-arranged in sheets of flat hexagons -sheets are bonded together by london forces -sheets can slide over each other -delocalised electrons can move so is electrical conductor -strong covalent bonds on hexagonal sheets means graphite has a very high melting point insoluble in any solvent
describe features of graphene
-one 2D layer of graphite -delocalised electrons(same as graphite) allow electrical conduction. without layers electrons can move faster=best electrical conductor known -delocalised electrons also strengthen covalent bonds making graphene extremely strong -single layer of graphene is transparent and extremely light
define metallic bond
the strong electrostatic attraction between the positive ions and the negatively charged delocalised electrons
why are metals good thermal conductors
metallic bonding means the delocalised electrons can pass kinetic energy to each other
why are metals good electrical conductors
metallic bonding means the delocalised electrons can carry current
why do metals have high melting and boiling points what are the 3 things that affect them
1) metallic bonds are strong. the more delocalised electrons the stronger the bonds. Mg2+>Na.
2) size of the metal ion, smaller ionic radius will hold electrons closer to the nuclei
3) lattice structure also affect the melting and boiling points
why do metals have the ability to be shaped
no bonds holding specific ions together so metal ions can slide over eachother when structure is pulled.
this makes metals malleable and ductile
explain the solubility of metals
- insoluble, except in liquid metals
- due to strength of metallic bonds
define periodicity
trends in physical and chemical properties of elements as you go across the periodic table
elements in a period have the
same number of shells
elements in a group have the
same number of outer electrons
what are the elements that touch the step ladder line
metalloids
define ionisation
removal of one or more electrons from an atom or molecule, resulting in an ion forming
define second ionisation energy
energy needed to change 1 mole of gaseous 1+ ions into 1 mole of gaseous 2+ ions
define first ionisation energy
energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms
write the general equation for second ionisation energy
X+(g) -> X2+(g) + e-
write the general equation for first ionisation energy
X(g) -> X+(g) + e-
what 3 factors affect ionisation energy
1) atomic radius (affected by 2+3)
2) nuclear charge
3) shielding
whats the trend in first ionisation energy across a period
increase
whats the trend in first ionisation energy down an group
decrease
what affects melting/boiling points when going across a period
type of bond
1) metal-meta bonds = high, increase across
2) giant covalent structures(B,C and Si) = highest, increases
3) simple molecular structures= london forces weak
group 2 metals react with water to give
write equation
a metal hydroxide and hydrogen
M(s) + 2H2O(l) -> M(OH)2(aq) + H2(g)
group 2 elements react more readily
down the group because ionisation energies decrease
when group 2 metals burn in oxygen you get
write general equation
solid white oxides
2M(s) + O2(g) -> 2MO(s)
when group 2 metals react react with dilute HCl you get
write general equation
a metal chloride and hydrogen
M(s) + 2HCl(aq) -> MCl2(aq) + H2(g)
group 2 acids are know as
commonly used for
alkaline earth metals
neutralising acids
boiling/melting points of halogens do what down the group and why
what happens to volatility
increase
strength of london forces increase as number of electrons increase
decreases down group
what reactions do halogens do
write general equation
reduction to halide ions
X + e- -> X-
are halogen oxidsing/reducing agents
oxidising
whats the trend in reacitivity of halogens down the group
decreases
define displacement reaction
a reaction where a more reactive element pushes out a less reactive element from an ionic solution
when testing for halogens organic solvent can be added to form a solvent layer which has very different colours. what are the colours for iodine, bromine, chloring
iodine = violet/pink
bromine = orange/red
chlorine = very pale yelow/green
- when testing for halides what test can be used
- hwo should this test be performed
- what are the results
silver nitrate test
1) dissolve in distilled water if solid
2) add each unknown substance to a test tube with dilte HNO3
3) add AgNO3 (silver nitrate)
observations: Cl-=white prcpt
Br-= cream prcpt
I-=yelow prcpt
if halides cant be identified via silver nitrate test what test can be used to further distinguish between halide precipitates
- add dilute ammonia, see if prcpt dissloves
- add conc ammonia, see if prcpt dissolves
observ. : Cl- : dissolves in dilue NH3
Br-= dissolves in conc. but not dilute NH3
I-= insouble
define disproportionation
when an element is both oxidised and reduced in a single chemical reaction
halogens undergo disproportionation reactions when they react with cold dilute alkali solutions
write the full and ionic equation for a halogen reacting with NaOH
full: X2 + 2NaOH -> NaXO + NaX + H2O
ionic: X2 + 2OH- -> XO- + X- + H2O
when testing for ions what order should the 3 tests be in
carbonates -> sulfates -> halides
describe how you would carry out the 3 tests for carbonates -> sulfates -> halides
carbonates: 1)dilute strong acid -> CO2 + H2O
2) CO2 released
Sulphates: 1)add barium nitrate Ba(NO3)2(aq)
2)white precipitate forms
Halides: 1)add nitric acid, then silver nitrate solution
2)precipitate forms colours = AgCl-white, AgBr-cream, AgI-yellow
how do you test for ammonium ions
Ammonia gas: 1)add NaOH and warm
1) hold damp red litmus paper over test tube
2) turns blue
(also smells)
