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A-Level Chemistry > 3.1 The periodic table > Flashcards

3.1 The periodic table Flashcards

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1
Q

periodic table arrangement

A

increasing atomic (proton) number
in periods showing repeating trends in physical and chemical properties (periodicity)
in groups having similar chemical properties

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2
Q

periodicity definition

A

repeating trend in properties in each element across a period

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3
Q

periodicity in electron configuration

A

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)

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4
Q

electron configuration trend down a group

A

same number of election in outer shell

gives elements in same group similar chemical properties

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5
Q

first ionisation energy definition

A

energy required to remove one electron from each atom in 1 mole of gaseous atom to form 1 mole of gaseous 1+ ions

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6
Q

factors affection ionisation energy

A 
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)
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7
Q

successive ionisation energies

A

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

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8
Q

how to work out number of electrons in shells with successive ionisation energies

A

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

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9
Q

first ionisation energy trend down a group

A

atomic radius increases
more inner shells, shielding increases
nuclear attraction on outer electrons decrease
first ionisation decreases

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10
Q

periodicity in first ionisation energy (general trend)

A 
nuclear charge increases
same number of shells, similar shielding
nuclear attraction increases
atomic radius decreases
first ionisation energy increases
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11
Q

why first ionisation drops from beryllium to boron

A

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

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12
Q

why first ionisation decreases from nitrogen to oxygen

A

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

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13
Q

metallic bonding definition

A

strong electrostatic attraction between cations and delocalised electrons

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14
Q

giant metallic structure lattice

A

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

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15
Q

properties of metals

A

strong metallic bonds
high electrical conductivity
high melting and boiling points

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16
Q

why metals have good electrical conductivity

A

in solid and liquid states

delocalised electrons can move as carry charge (charge carriers) in structure when voltage applied

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17
Q

why melting and boiling points of metals are high

A

strong electrostatic attraction between cations and delocalised electrons
high temperatures necessary to provide large amount of energy to overcome metallic bonding

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18
Q

solubility of metals

A

don’t dissolve

any interaction with water results in a reaction

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19
Q

giant covalent lattices

A

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)

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20
Q

properties of giant covalent lattices

A

high MP/BP
insoluble in almost all solvents
non-conductors (apart from graphene and graphite)

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21
Q

why giant covalent lattices are generally insoluble

A

covalent bonds between each atom too strong to be broken by interactions with solvent

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22
Q

why giant covalent lattices have high MP/BP

A

covalent bonds are strong

require lots of heat energy to break covalent bond

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23
Q

electrical conductivity in giant covalent lattices

A

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)

24
Q

why MP/BP increases across period (for metals)

A

more delocalised electrons
increasing charge density (charge increases, ionic radius decreases)
increasing strength of metallic bond
more heat energy required to overcome stronger bonds

25
Q

periodicity in melting points

A

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)

26
Q

electron configuration of group 2 elements

A

2 outer shell electrons in s sub-shell

27
Q

group 2 elements in redox reactions

A

loses 2 outer shell electrons to form 2+ ions

reducing agent

28
Q

oxidation definition

A

loss of electrons

increasing oxidation number

29
Q

reduction definition

A

gaining electrons

decreasing reduction number

30
Q

why reactivity increases down group 2

A

attraction between nucleus and outer shell electrons decrease
increased atomic radius and shielding
ionisation energies decrease down group

31
Q

group 2 oxides reaction with water

A

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)

32
Q

solubility and alkalinity of hydroxides down group 2

A

increases (more OH- ions disassociate, increasing pH)

33
Q

solubility and alkalinity of magnesium and barium hydroxide

A

Mg(OH)2: very slightly soluble, ~10 pH

Ba(OH)2: much more soluble, ~13 pH

34
Q

uses of group 2 compounds in agriculture

A

Ca(OH)2 (lime) increases pH in acidic soil

Ca(OH)2(s) + 2H+(aq) -> Ca2+(aq) + 2H2O(l)

35
Q

group 2 compounds in medicine

A

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

36
Q

carbonate test

A

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

37
Q

reaction of CO2 with limewater

A

CO2(g) + Ca(OH)2(aq) -> CaCO3(s) + H2O(l)

38
Q

sulfate test

A

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

39
Q

ionic equation for sulfate test

A

Ba^2+ (aq) + SO4^2-(aq) -> BaSO4(s)

40
Q

halide/measuring rate to hydrolysis method

A

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)

41
Q

correct order of tests and why

A

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)

42
Q

test for ammonia

A

add sodium hydroxide and agitate
heat gently with Bunsen while holding damp red litmus paper in test tube
paper turns blue = ammonium ions

43
Q

trend of boiling point down halogen

A

more electrons, larger molecule
stronger London forces
more energy required to overcome IM forces
boiling point increases down group

44
Q

fluorine in RTP

A

pale yellow gas

45
Q

chlorine in RTP

A

pale green gas

46
Q

bromine in RTP

A

red-brown liquid

47
Q

iodine in RTP

A

shiny greasy-black solid

48
Q

astatine in RTP

A

never been seen

49
Q

halogens in redox reactions

A

forms 1- anions

oxidising agents

50
Q

how to differentiate between bromine and iodine in solution

A

add cyclohexane

iodine is deep violet

51
Q

trend in reactivity down halogens

A

atomic radius increases
more inner shells (more shielding)
less nuclear attraction to capture electron from another species
reactivity decreases

52
Q

disproportionation definition

A

reaction when an element is oxidised and reduced simultaneously

53
Q

chlorine + water

A

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)

54
Q

chlorine + sodium hydroxide solution

A

Cl2(aq) + 2NaOH(aq) -> NaClO(aq) + NaCl(aq) + H2O(l)

more chlorate ions dissociate than with just water (from NaClO)

55
Q

benefits of chlorine use

A

makes sure water is drinkable

prevents outbreaks from waterborne diseases e.g. cholera and typhoid

56
Q

risks of chlorine use

A

chlorine gas is respiratory irritant and toxic

chlorinated hydrocarbons formed by chlorine + organic hydrocarbons suspected of causing cancer

57
Q

transition element definition

A

an element that forms a stable ion with a partially filled d orbital

A-Level Chemistry (12 decks)

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