3.1 The periodic table Flashcards
periodic table arrangement
increasing atomic (proton) number
in periods showing repeating trends in physical and chemical properties (periodicity)
in groups having similar chemical properties
periodicity definition
repeating trend in properties in each element across a period
periodicity in electron configuration
electron added to highest energy level
s sub-shell filled in group 1-2, p sub-shell filled onwards
same in each period (periodic pattern)
electron configuration trend down a group
same number of election in outer shell
gives elements in same group similar chemical properties
first ionisation energy definition
energy required to remove one electron from each atom in 1 mole of gaseous atom to form 1 mole of gaseous 1+ ions
factors affection ionisation energy
atomic radius (greater distance between nucleus and outer electrons = less nuclear attraction, large effect) nuclear charge (more protons in nucleus = greater attraction between nuclear and outer electrons) electron shielding (inner shell electrons repel outer shell electrons due to negative charges, reduces attraction between nucleus and outer electrons - shielding effect)
successive ionisation energies
greater than previous ionisations
first electron is lost
same number of protons, less electrons so greater attraction to each electron
pulls electron closer to nucleus
nuclear attraction increases in outer shell electron so more energy required to ionise successively
how to work out number of electrons in shells with successive ionisation energies
sharp increase = new shell
number of ionisations in line = number of electrons in shell
can work out element, group of element, number of electrons in outer shell
first ionisation energy trend down a group
atomic radius increases
more inner shells, shielding increases
nuclear attraction on outer electrons decrease
first ionisation decreases
periodicity in first ionisation energy (general trend)
nuclear charge increases same number of shells, similar shielding nuclear attraction increases atomic radius decreases first ionisation energy increases
why first ionisation drops from beryllium to boron
only 2s sub-shell filled in beryllium
2p sub-shell also filled in boron
2p on higher energy level than 2s so easier to remove
first ionisation lower in boron than beryllium despite higher nuclear charge
why first ionisation decreases from nitrogen to oxygen
highest energy level elections in 2p sub-shell
oxygen has paired electrons in one of 2p orbitals
repel one another, nuclear attraction decreases
first ionisation energy less in oxygen than nitrogen
metallic bonding definition
strong electrostatic attraction between cations and delocalised electrons
giant metallic structure lattice
atoms donate outer shell electrons to form sea of delocalised electrons
cations fixed in position
delocalised electrons mobile and able to move throughout structure
charges of cations and delocalised electrons balance
electrons spread out and shared between all cations
forms giant metallic lattice
properties of metals
strong metallic bonds
high electrical conductivity
high melting and boiling points
why metals have good electrical conductivity
in solid and liquid states
delocalised electrons can move as carry charge (charge carriers) in structure when voltage applied
why melting and boiling points of metals are high
strong electrostatic attraction between cations and delocalised electrons
high temperatures necessary to provide large amount of energy to overcome metallic bonding
solubility of metals
don’t dissolve
any interaction with water results in a reaction
giant covalent lattices
billions of atoms held together by network of strong covalent bonds
forms giant covalent lattice
carbon and silicon form tetrahedral structure (group 4, electron-pair repulsion)
properties of giant covalent lattices
high MP/BP
insoluble in almost all solvents
non-conductors (apart from graphene and graphite)
why giant covalent lattices are generally insoluble
covalent bonds between each atom too strong to be broken by interactions with solvent
why giant covalent lattices have high MP/BP
covalent bonds are strong
require lots of heat energy to break covalent bond
electrical conductivity in giant covalent lattices
diamond and silicon can’t as all four outer shell electrons are used in covalent bonding
graphite and graphene can (not all electrons used in covalent bonding)
why MP/BP increases across period (for metals)
more delocalised electrons
increasing charge density (charge increases, ionic radius decreases)
increasing strength of metallic bond
more heat energy required to overcome stronger bonds
periodicity in melting points
increasing from group 1 to 3 strong metallic bonds)
increase from group 3 to 4 (giant metallic lattices to giant covalent lattices)
sharp decrease from group 4 to 5 (giant covalent lattices to simple covalent structures)
generally low onwards (simple covalent structures)
electron configuration of group 2 elements
2 outer shell electrons in s sub-shell
group 2 elements in redox reactions
loses 2 outer shell electrons to form 2+ ions
reducing agent
oxidation definition
loss of electrons
increasing oxidation number
reduction definition
gaining electrons
decreasing reduction number
why reactivity increases down group 2
attraction between nucleus and outer shell electrons decrease
increased atomic radius and shielding
ionisation energies decrease down group
group 2 oxides reaction with water
metal oxide(s) + water (l) -> metal hydroxide (s) (after solution becomes saturated)
metal oxide(s) + water(l) -> Ca2+(aq) + 2OH-(aq) (before solution is fully saturated)
solubility and alkalinity of hydroxides down group 2
increases (more OH- ions disassociate, increasing pH)
solubility and alkalinity of magnesium and barium hydroxide
Mg(OH)2: very slightly soluble, ~10 pH
Ba(OH)2: much more soluble, ~13 pH
uses of group 2 compounds in agriculture
Ca(OH)2 (lime) increases pH in acidic soil
Ca(OH)2(s) + 2H+(aq) -> Ca2+(aq) + 2H2O(l)
group 2 compounds in medicine
antacids (milk of magnesia, indigestion tablets)
Mg(OH)2(s) + 2HCl(aq) -> MgCl2(aq) + 2H2O(l)
CaCO3(s) + 2HCl(aq) -> CaCl2(aq) + H2O(l) + CO2(g)
or magnesium carbonate
carbonate test
add dilute nitric acid to test solution
collect gas produced using clean pipette
bubble gas through lime water (dissolved Ca(OH)2)
place bung over test tube
if cloudy, gas was CO2 and carbonate is present in solution
reaction of CO2 with limewater
CO2(g) + Ca(OH)2(aq) -> CaCO3(s) + H2O(l)
sulfate test
add equal volume of HNO3 to unknown solution and mix
add equal volume BaNO3(aq)
if dense white precipitate formed (barium sulfate), sulfate is present in unknown solution
ionic equation for sulfate test
Ba^2+ (aq) + SO4^2-(aq) -> BaSO4(s)
halide/measuring rate to hydrolysis method
ethanol acts as a solvent
add dilute nitric acid to solution
add 2cm^3 of AgNO3(aq)
heat in a water bath at the same temperature
white precipitate = chloride
cream precipitate = bromide
yellow precipitate = iodide
add dilute ammonia (clear = chloride, not clear = bromide/iodide)
add concentrated ammonia under fume cupboard (clear = chloride/bromide, not clear = iodide)
correct order of tests and why
carbonate
sulfate
halide
sulfate and halide ions don’t produce gases (if gases produced carbonate has to be present)
BaCO3 is also white insoluble precipitate like BaSO4 (carbonate test must be done before sulfate test)
AgCO3 and AgSO4 are both insoluble precipitates like silver halide (other two tests must be done before)
test for ammonia
add sodium hydroxide and agitate
heat gently with Bunsen while holding damp red litmus paper in test tube
paper turns blue = ammonium ions
trend of boiling point down halogen
more electrons, larger molecule
stronger London forces
more energy required to overcome IM forces
boiling point increases down group
fluorine in RTP
pale yellow gas
chlorine in RTP
pale green gas
bromine in RTP
red-brown liquid
iodine in RTP
shiny greasy-black solid
astatine in RTP
never been seen
halogens in redox reactions
forms 1- anions
oxidising agents
how to differentiate between bromine and iodine in solution
add cyclohexane
iodine is deep violet
trend in reactivity down halogens
atomic radius increases
more inner shells (more shielding)
less nuclear attraction to capture electron from another species
reactivity decreases
disproportionation definition
reaction when an element is oxidised and reduced simultaneously
chlorine + water
Cl2(aq) + H2O(l) -> HClO(aq) + HCl(aq)
chlorine is oxidised and reduced(disproportionation)
hydrochloric acid and chloric acid kill microbes
when indicator is added, turns red then white (chloric acid acts as bleach)
chlorine + sodium hydroxide solution
Cl2(aq) + 2NaOH(aq) -> NaClO(aq) + NaCl(aq) + H2O(l)
more chlorate ions dissociate than with just water (from NaClO)
benefits of chlorine use
makes sure water is drinkable
prevents outbreaks from waterborne diseases e.g. cholera and typhoid
risks of chlorine use
chlorine gas is respiratory irritant and toxic
chlorinated hydrocarbons formed by chlorine + organic hydrocarbons suspected of causing cancer
transition element definition
an element that forms a stable ion with a partially filled d orbital
