Topic 4: inorganic chemistry and the periodic table Flashcards
State and explain the trend in MP going down group 2 metals
MP decreases. themetallic bond weaks as the atomic size increases. electrostatic attraction between the positive ions and the delocalised electrons decreases therefore less energy needed to break the bonds.
State and explain the trend in 1st ionisation energy going down group 2
1st ionisation energy decreases. atoms are getting larger and therefore the outermost electrons are further from the nucleus and more shielded, decreasing the attraction between the nucleus and the electrons therefore less energy needed to remove the outer electrons.
state the trend in reactivity going down group 2 metals
reactivity increases down the group as its easier to lose the electrons
magnesium + oxygen
2Mg + O2 -> 2MgO
MgO is a white solid with a high mp due to its ionic bonding
Mg burns with a bright white flame
when using magnesium ribbon, why do you need to clean it before use
magnesium reacts slowly with oxygen without a flame. therefore the magnesium ribbon will have a thin layer of magnesium oxide on it.
if testing for reaction rates with Mg and acid, the Mg and MgO would react at different rates, giving false results
MgO + HCl-> MgCl2 + H2O
magnesium + steam
Mg + H2O(g) -> MgO + H2
bright white flame
magnesium + warm water
Mg + 2H2O -> Mg(OH)2 + H2
much slower reaction than with steam and no flame
group 2 metals + cold water products
froms the hydroxide (aq) + H2 (g)
observations when group 2 metals react with hydroxides and trends in the reaction going down the group
fizzing (morevigorous down group)
the metal dissolves(faster down the group)
solution heats up (more down the group)
calcium + cold water observation
a white precipitate appears (less precipitate forms down the group)
group 2 oxides are ionic/basic?
basic as the oxide ions accept protons to become hydroxide ions (when reacting with water)
groups 2 oxides + water products
just the hydroxide (s)
reaction with Mg gives pH9, with calcium gives pH12. This is because magnesium is only slightly soluble in water so fewer free OH- produced
group 2 oxides + acid
-> salt (aq) + water
groups 2 hydroxides + acids
-> salt (aq) + water
Eg
Mg(OH)2 (aq) + 2HNO3 (aq) -> Mg(NO3)2 (aq) + 2H2O(l)
solubility of group 2 hydroxides
become more soluble down the group
all group 2 hydroxides when not soluble appear as __
white precipitates
how is magnesium hydroxide used in medicine
in suspension as milk of magnesia to neutralise excess acid in the stomach and to treat constipation.
it is safe to use because it is so weakly alkaline. more preferable to use than calcium carbonate as it will not produce CO2
calcium hydroxide use in agriculture
neutralise acidic soils
how can an aqueous solution of calcium hydroxide be used to test for CO2
aqueous solution of calcium hydroxide is limewater.
limewater turns cloudy as the white calcium carbonate is produced.
Ca(OH)2 (aq) + CO2 -> CaCO3 (s) + H2O (l)
barium hydroxide + water ionic equation
Ba(OH)2 (s) + aq -> Ba2+ (aq) + 2OH- (aq)
solubility of group 2 sulfates
group 2 sulfates become less soluble going down the group
why does barium metal react slowly with sulfuric acid
insoluble barium sulfate produced will cover the surface of the metal and act as a barrier to further attack
same effect does not happen with other acids like HCL or HNO3 as they form soluble group 2 salts
thermal decomposition of group 2 carbonates
group 2 carbonates become more thermally stable going down the group. as cations get bigger they have less of a polarising effect and distort the carbonate less therefore the C-O bond is weakened and more energy needed to break it.
what is the group 1 carbonate that is the exception and will decompose
lithium carbonate
lithium ion is small enough ot have a polarising effect on the carbonate ion. it has a big enough charge density
is it easier or harder to thermally decompose group 2 carbonates going down the group
going down the group it gets harderto thermally decompose them
define thermal decomposition
the use of heat to break down a compound into more than one product
MgCO3 (s) ->
MgO (s) + CO2 (g)
how would you investigate thermal stability of group 2 carbonates
heata known mass of carbonate in a side arm boiling tube
pass the gass produced through limewater
time for limewater to go cloudy
thermal decomposition of group 2 nitrates (V)
becomes harder to decompose going down the group
they decompose to produce group 2 oxides and nitrogen dioxide gas
(lithium nitrate decomposes the same as group 2 nitrates)
observation when decomposing group 2 nitrates
brown gas (NO2) whitenitrate solid is seen to melt to a colourless solution and then re solidify
equation for thermal decomposition of magnesium nitrate
2Mg(NO3)2 -> 2MgO + 4NO2 +O2
how do group 1 nitrates, with the exception of lithium nitrate decompose
decompose to give a nitrate (III) salt and ocygen
2NaNO3 -> 2NaNO2 + O2
flame test
nichrome wire
clean by dipping inot conc HCL
dip into solid and then into blue flame
observe colour change
explain flame tests
an excited electron is promoted to a higher energy level by the heat from the flame
when the electron drops to a lower energy level, energy is released in the from of visible light
why is there no flame colour with magnesium
energy emitted of a wavelength outside visible spectrum
appearance of halogens at room temp
fluorine- very pale yellow gas, highly reactive
chlorine- greenish, reactive gas, poisonous in high conc
bromine- red liquid, it gives off dense brown/orange poisonous fumes
iodine- dark grey solid, sublimes to purple gas
trend in MP and BP of halogens down the group
increases down the group
molecules are larger, more electrons, larger London forces, more energy needed to overcome the intermolecular forces
trend in electronegativity of halogens
going down the group, the EN decreases
atomic radii increases, nucleus is less able to attract the bonding pair of electrons
define electronegativity
electronegativity is the relative tendency of an atom in a molecule to attract electrons in a covalent bond to itself
oxidising strength of halogens down the group
oxidising strength decreases down the group- harder to accept electrons
potassium bromide + chlorine ionic eq
Cl2(aq) + 2Br- (aq) -> 2Cl- (aq) + Br2 (aq)
potassium iodide + chlorine ionic eq
Cl2 (aq) + 2I- (aq) -> 2Cl- (aq) + I2 (aq)
potassium iodide + bromine ionic eq
Br2 (aq) + 2I- (aq) -> 2Br- (aq) + I2 (aq)
halogen + metal reactions- what is oxidsed
the metals are oxidised
Br2 (l) + 2Na(s) -> 2NaBr (s)
reaction between chlorine/ bromine with iron (II)
iron (II) is oxidised to (III)
Cl2 (g) +2Fe2+ (aq) -> 2Cl- (aq) + 2Fe3+ (aq)
reaction between iodine and iron (II)
iodine is not a strong oxidising agent to oxidise iron (II) to iron (III). the reaction is reversed
2I- (aq) + 2Fe3+ (aq) -> I2 (aq) + 2Fe2+ (aq)
define disproportionation
a reaction where an element is simultaneously oxidised and reduced
chlorine + water
disproportionation
Cl2 (g) + h2o(l) -> HClO (aq) +HCl (aq)
observations with chlorine + water disproportionation reaction
adding universal indicator will first turn it red due to acidity of both reaction products. then turns colourless as HCLO bleaches the colour
useof chlorine
usedin water treatment. used to treat drinking water and water in swimming pools
reaction of aqueous halogens with cold dilute NaOH solution
Cl2 9aq) +2NaOH (aq) ->NaCl (aq) + NaClO (aq) +H2O (l)
mixture of NaCl and NaClO is used as bleach and to disinfect/kill bacteria
reaction of halogens with hot dilute NaOH solution
disproportionation occurs but the halogen is oxidised
3Cl2 (aq) +6NaOH(aq) -> 5NaCl (aq) +NaClO3 (aq) +3H2O (l)
explain trend in reducing power of HALIDES going down the group
increasing reducing power down the group
they have a greater tendency to donate electrons as ions get bigger so attraction between nucleus ans the outer electrons decreases
sulfuric acid+ fluoride/chloride
no redox reaction occurs as the sulfuric acid is not strong enough to oxidise the halides.
acid base reactions occur
NaF (s) + H2SO4 (l) -> NaHSO4 (s) +HF(g)
HF is white steamy fumes
sulfuric acid + bromide
after the intial acid bas step, bromide reduces the sulfur in H2SO4 from 6+ to 4+ in SO2
acid basestep: NaBr(s) + H2SO4(l) -> NaHSO4 (s) +HBr (g)
redox step: 2HBr + H2SO4 -> Br2 (g) +SO2(g) +2H2o (l)
in first step sulfuric acid acts as an acid but in 2nd step acts as an oxidising agent as itgets reduced
observations of sulfuric acid + bromide
in acid base step: white steamy fumes of HBr
in redox step: red fumes of Br2 (g) and a colourless, acidic gas SO2 (g)
iodide + sulfuric acid (4 steps)
NaI(s) + H2SO4 (l) -> NaHSO4 (s) +HI
2HI + H2SO4 -> I2 (s) + SO2(g) +2H20 (l)
6HI + H2SO4 -> 3I2 + S (s) + 4H2O(l)
8HI + H2SO4 -> 4I2 (s) + H2S (g) + 4H2O (l)
oxidation numbers of sulfur when reacting with iodide
+6 in H2SO4 to
+4 in SO2 to
0 in S to
-2 in H2S
observations with iodide and sulfuric acid
HI gives white steamy fumes black solid and purple fumes of I2 colourless acidic gas of SO2 yellow solid of sulfur bad egg smell gas H2S
