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Chemistry Year 12 > 3.1 > Flashcards

3.1 Flashcards

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1
Q

brief periodic table history

A

döbereiner: groups of three elements (triads) ordered by atomic weight had similar properties
newlands: arranged in atomic weight and suggested law of octaves every eighth had similar properties
mendeleev: still in atomic weight but left gaps for undiscovered elements and predicted their properties also switched elements if the atomic weight order didn’t match the properties
moseley: arranged by atomic number not weight

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2
Q

what is periodicity

A

the trend in properties that is repeated across each period

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3
Q

what are the three things to mention when explaining a trend in periodicity

A

electron shielding
atomic radius
nuclear charge

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4
Q

what is the first ionisation energy

A

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

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5
Q

write the first ionisation equation for chlorine

A

Cl (g) —> Cl+(g) + e-

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6
Q

what are the rules when creating a first ionisation equation

A

must be gaseous
only 1+ because it’s the first electron
never diatomic
never minus because an electron is being removed

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7
Q

what are successive ionisation energies

A

a measure of the amount of energy required to remove each electron in turn

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8
Q

what happens to electron attraction when an electron is removed

A

there is less repulsion between the remaining electrons and each shell is drawn in slightly closer to the nucleus

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9
Q

what’s the trend in ionisation energy across a period

A

increases across the period (because smaller atomic radi and higher nuclear charger) but decreases between group 2 and 3 because it goes from s subshell to p subshell (p orbital have higher energy than s orbitals) also a decrease from group 5 to 6 (p3 to p4) because teh orbital pairs up causing a repulsion which destabilises the atom and makes it easier to lose an electron

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10
Q

what happens to ionisation energy down a group

A

decreases

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11
Q

what is second ionisation energy

A

the energy required to remove one electron from each 1+ ion in 1 mole of the gaseous element to form 1 mole of 2+ ions

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12
Q

structure of metallic bonding

A

positive ions in a fixed giant lattice
outer shell electrons are delocalised in a sea of delocalised electrons that is spread throughout the structure and can move within. the metal is held together by the attraction between all the positive ions and all the negative electrons

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13
Q

why do metals have a high melting and boiling point

A

the attraction between the positive ions and negative atoms is very strong
high energy is needed to dislodge the ions from their rigid positions in the lattice

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14
Q

what does metallic bonding strength rely upon

A

number of outer electrons donated to the electron cloud (from group number) and the size of the metal ion

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15
Q

metallic bonding strength trends across the period and down the group

A

increases across period (more electron delocalised)
decreases down group (lower electron density)

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16
Q

why can metals conduct electricity

A

they have mobile delocalised electrons that can flow through the structure electrons attracted to the positive end are replaced by those entering from the negative end

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17
Q

why are metals ductile and malleable

A

because the layers can slide over eachother due to the delocalised electrons

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18
Q

melting point trends across a period

A

increasing from group 1-3 metals, then at group 4 (or three aswell in boron’s case) increases further because of giant covalent then dramatically decreases because from group 5 onwards they are simple molecularly bonded (just depends on number of electrons cos london forces)

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19
Q

what’s the reactivity trend down group 2

A

increases down the grouo

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20
Q

physical properties of group 2

A

reasonably high melting and boiling points
light metals with low densities
form colourless compounds

21
Q

what do group two metals form when they react with oxygen

A

reacts vigorously to form a metal oxide

2M + O2 —> 2MO

22
Q

what is formed when group 2 elements react with water and what is the exception

A

they all except from beryllium form hydroxides with ther general formula M(OH)2 hydrogen gas is also formed

23
Q

what’s the trend down group 2 with reactivity with water

A

Be - doesn’t
Mg - very slowly
increases down the grouo

24
Q

how do group 2 react with dilute acids

A

all except Be react to form a salt and hydrogen gas - more vigorous down the group

25
Q

how are the properties of group two metal hydroxides different down the group

A

they get more soluable down the group causing the pH to rise

26
Q

what are group 2 compounds used for

A

neutralising souls and indigestion tablets

27
Q

what’s the typical pH of a group 2 hydroxide in solution

A

10-12

28
Q

solubility of group two hydroxides

A

increases down the group, but beryllium doesn’t form a hydroxide and BeO isn’t soluble in water

29
Q

what’s an issue with using group two elements in buildings

A

they readily react with acids (eg rainwater is slightly acidic)

30
Q

what happens to melting point down group 7 and why

A

increases because there are strong london forces between them

31
Q

what’s the reactivity trend down group seven

A

reactivity decreases

32
Q

what colour is chlorine in water and cyclohexane

A

pale green

33
Q

what colour is bromine in water and cyclohexane

A

orange

34
Q

what colour is iodine in water and cyclohexane

A

brown in water violet in cyclohexane

35
Q

what is disproportionation

A

the oxidation and reduction of the same element in a redox reaction

36
Q

what is a displacement reaction

A

an element in a compound is replaced with a more reactive element to form a new compound

37
Q

what’s chlorines reaction with water

A

Cl2 + H2O <—> HCl + HOCl
0 -1 +1

38
Q

what’s chlorines reaction with aqueous sodium hydroxide to make bleach and what conditions and oxidation numbers

A

cold dilute aqueous sodium hydroxide
2NaOH + Cl2 —> NaClO + H2O + NaCl
0 +1 -1
bleach

39
Q

what order should you do the negative ion tests in and why

A

carbonate —> sulphate —> halide

becuase they can give false positives

40
Q

carbonate test

A

add dilute strong acid
collect gas formed
bubble through limewater
cloudy = carbonate

41
Q

ionic equation for the carbonate test

A

CO3 ^2- (aq) + 2H^+ (aq) —> H2O (l) + CO2 (g)

42
Q

sulphate test

A

add dilute acid and barium chloride
white precipitate = sulphate

43
Q

sulphate test ionic equation

A

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

44
Q

halide test

A

dissolve in water
acidify with dilute nitric acid
add few drops of silver nitrate solution
white precipitate = chloride
cream precipitate = bromide
yellow precipitate = iodide

45
Q

how do you confirm which halogen it is in halide test

A

dissolves in dilute ammonia = chloride
dissolved in concentrated ammonia = bromide
doesn’t dissolve in any ammonia = iodide

46
Q

halide test ionic equation

A

Ag^+ (aq) + X^- (aq) —> AgX (s)

47
Q

ammonium ion test

A

add sodium hydroxide solution and GENTLY WARM. will stink if ammonia but test with damp red litmus paper and it will turn blue

48
Q

ionic equation for ammonium test

A

NH4^+(aq) + OH^-(aq) —> NH3(aq)+H2O(l)

Chemistry Year 12 (7 decks)

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