Module 3 (chapter 7 and 8) - periodicity, group 2 and 17 Flashcards

1
Q

what did Mendeleev do?

A
  • arranged the 60 elements known at the time into the ‘periodic table’ and left gaps for undiscovered elements. (e.g. predicted ekasilicon)
  • arranged them in columns of similar chemical properties
  • if an element appeared to be in the wrong place due to its atomic weight he moved it to where it fitted with the pattern he had discovered (arranged by mass number)
  • discovered the Nobel gases in 1890s which fitted in his final group
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2
Q

what did Moseley do?

A
  • arranged the periodic table by atomic number
  • using an X-ray gun fired at elements he measured the wavelength of X-rays given, using it to calculate the frequency
  • when square rooted and plotted against atomic number, the graph showed a perfect straight line
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3
Q

groups

A
  • the vertical columns

- each element in a group has the same number of outer shell electrons and therefore similar chemical properties

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4
Q

periods

A
  • horizontal rows
  • the number of the highest energy electron shells in an element’s atom (e.g. period 2, the 2s sub shell fills with two electrons followed by the 2p sub shell with six electrons)
  • for each period, the S- and P-s sub shells are filled in the same ways
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5
Q

blocks

A
  • the elements in the periodic table can be divided into blocks corresponding to their highest energy sub shell
  • this gives four distinct blocks (s, p, d, f)
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6
Q

ionisation energy

A

measures how easy an atom loses an electron to form positive ions

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7
Q

first ionisation energy

A

the energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions
-(e.g. Na (g) –> Na+ (g) + e-)

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8
Q

what factors affect ionisation energy

A
  • atomic radius = the greater the distance between the nucleus and the outer electrons, the less the nuclear attraction. the force of attraction falls of sharply with increasing distance so atomic radius has a large effect
  • nuclear charge = the more protons there are in the nucleus of an atom, the larger the attractive force on the outer electrons
  • electron shielding = electrons are negatively charged and so inner shell electrons repel outer shell electrons. this repulsion is called the shielding effect, reducing the attraction between the nucleus and outer shell electrons
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9
Q

what are successive ionisation energies?

A
  • an element has many ionisation energies as there are electrons
  • second ionisation energy is the energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions
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10
Q

what do successive ionisation energies prove?

A
  • evidence for the different electron energy level in an atom (shells) as by analysing jumps scientists can deduce:
  • they allow predictions to be made about the number of electrons in the outer shell
  • the groups of the element in the periodic table
  • the identity of an element
  • there is a general increase in ionisation energy across each period but a sharp decrease between the end of one period and the start of another.
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11
Q

what is the general trend in successive ionisation energies?

A
  • as each electron is removed there is less repulsion between the remaining electrons and so each shell will be drawn in slightly closer to the nucleus
  • the positive nuclear charge will outweigh the negative charge every time an electron is removed
  • as the distance of each electron from the nucleus decreases slightly, the nuclear attraction increases and so more energy is needed to remove an electron.
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12
Q

trend in first ionisation energy down a group

A
  • atomic radius increases
  • more inner shells so shielding increases
  • nuclear attraction on outer shell electrons decreases
  • first ionisation energy decreases
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13
Q

trend in first ionisation energy across a period

A
  • nuclear charge increases due to more protons
  • same shell so similar levels of shielding
  • nuclear attraction increases
  • atomic radius decreases
  • first ionisation energy increases
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14
Q

on an ionisation graph, what do the small drops in energy represent?

A
  • indicate a new sub shell which is slightly further away form the nucleus (e.g. boron less than beryllium)
  • the even smaller drops reflect that despite an increased nuclear charge there are two electrons in the same orbital compared to an unpaired electron. these two repel each other rathe than spinning at right angles where there is equal repulsion in paired, making it easier to remove (e.g. oxygen lower than nitrogen)
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15
Q

what are ionisation energies used in

A

plasma displays where gas is electrically ionised to form a mix of positive ions.

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16
Q

why does atomic radius increase down a group but decrease across a period

A
  • increases across a group as there are more shells so shielding increases. despite the nucleus getting mote protons the shielding negates this increased nuclear charge
  • decreases across a period due to more protons and therefore a higher nuclear charge. there is the same number of electron shells and therefor shielding so there is a greater pull of electrons to the nucleus
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17
Q

where is the metal non-metal divide

A
  • top of group 13 to bottom of group 17

- elements near this divide (e.g. Si) are often called semi metals or metalloids

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18
Q

metallic bonding

A
  • strong electrostatic force of attraction between positive metal ions and delocalised electrons
  • regularly arranged rows of metal cations held together by a sea of delocalised electrons
  • the delocalised electrons are there because each metal atom donates it outer shell electrons to a shared pool of electrons which are delocalised throughout the whole structure
  • for 12 cation each with 1+ charge there are 12 electrons each with a 1- charge.
  • for metals containing 2+ cations, twice as many electrons are present to balance the charge
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19
Q

why do metals conduct electricity?

A
  • in solid and liquid states
  • when a potential differences applied, electrons are delocalised and can therefor remove through the whole structure, carrying charge
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20
Q

why do metals have high melting and boiling points?

A
  • strong electrostatic forces of attraction between cations and electrons
  • these require a large amount of energy to overcome
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21
Q

why aren’t metals soluble?

A
  • it is expected that there would be some interaction between polar solvents and the charges in the metallic lattice
  • but any interactions would lead to a reaction rather than dissolving
  • giant lattice structures have which are far too strong
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22
Q

what are other properties of metals?

A
  • ductile

- malleable as electrons can move and be bent into shapes without shattering the structure and breaking bonds

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23
Q

why does the boiling point of metals increase across a period?

A
  • more delocalised electrons so stronger electrostatic force of attraction between oppositely charged ions
  • there is also a greater charge density (ionic charge is greater and ionic radius decreases) so electrons are more attracted to the nucleus
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24
Q

periodicity

A

a repeating pattern in either chemical or physical properties across different periods

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25
Q

why does boiling point decrease form S-Ar?

A
  • exist as (s6, Cl2 and Ar)
  • there are less electrons within the molecule meaning the induced dipole dipole interactions are weaker and so less energy is needed to overcome them
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26
Q

why is there a sharp decrease in melting and boiling point between group 14 and 15?

A
  • marks a change from giant to simple molecular structures
  • simple molecular have weak intermolecular forces to overcome in comparison to strong covalent bonds so have lower boiling points
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27
Q

facts about group 2 elements

A
  • alkaline earth metals (due to alkaline properties of their metal hydroxides)
  • the elements are reactive metals and not appear naturally in their elemental form
  • the formation of 2+ ions form gaseous atoms requires the input of two ionisation energies
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28
Q

most common reaction of group 2 elements

A
  • redox reactions
  • each metal atom is oxidised losing two electrons to form a 2+ ion with the electron configuration of the corresponding Nobel gas
  • they will all react with oxygen to form a metal oxide with the general formula MO made up of M(2+) and O(2-)
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29
Q

group 2 elements reactions with water and acid

A
  • react to form an alkaline hydroxide solution with the general formula M(OH)2 and H2 gas
  • this reactions is very slow but becomes more vigorous further down the group due to increasing reactivity
  • redox with dilute acids to form a salt and hydrogen gas.
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30
Q

why does ionisation decrease down group 2

A

the attraction between the nucleus and outer shell electrons decreases due to increased atomic radius and shielding

  • although other energy changes take place when they react, the first and second ionisation energies make up the most of the energy input
  • the total energy input to form 2+ ions decreases down the group
31
Q

group 2 oxides and hydroxides

A
  • the oxides of group 2 elements react with water releasing hydroxide ions
  • only slightly soluble in water as when the solution become saturated, any further metal and hydroxide ions will form a solid precipitate
  • the solubility of hydroxides in water increases down the group so the solution contains more OH- ions and are more alkaline
32
Q

test to show the trend in reactivity of group 2 metal oxides with water

A
  • add a spatula of each group 2 oxide to water in a test tube
  • shake the mixture. on this scale their is insufficient water yo dissolve all of the metal hydroxides that from so you will have some white solid undissolved at the bottom of the test tube
  • measure the PH of each solution, the alkalinity will be seen to increase down the group
33
Q

what are uses of group 2 bases in agriculture

A
  • calcium hydroxide
  • used in agriculture- is added to fields as lime by farmers to increase the PH of the acidic soils
  • it neutralises the acid in soil forming neutrals water
34
Q

uses of group 2 bases in medicines

A
  • uses as antiacids for treating indigestion causes by excess acid in the stomach
  • magnesium and calcium carbonate
  • this allows neutralisation to happen in your stomach
  • magnesium hydroxide is used as milk of magnesia
35
Q

uses of group 2 bases for chimneys

A
  • calcium carbonate is sprayed down power station chimneys to neutralise suffer dioxide and other acidic gases
  • this is a scrubber helping to combat acid rain
  • also used a flogulent in sewage
36
Q

why does alkalinity increase down group 2

A
  • become more soluble in water
  • this means the solution contains more OH- ions and so will be more alkaline
  • at the top of the group they are only slightly soluble and so become saturated quickly
  • this means any further metal and hydroxide ions will form a solid precipitate
37
Q

why do group 2 elements become more reactive down the group

A
  • number of electrons and therefor electron shells increases meaning the atomic radius is larger.
  • this means there is increased shielding form inner electrons and so the attraction between nucleus and outer shell electrons is weaker
  • makes it easier to lose to the metal atoms two outer shell electrons
38
Q

physical appearance of halogens

A
  • fluorine is a pale yellow gas (reacts with almost any substance it comes into contact with)
  • chlorine is a pale green gas
  • bromine is a brown/orange liquid
  • iodine is a grey solid
  • astatine is a black solid (never actually been seen as is very radioactive so decays rapidly)
39
Q

why does the boiling point of halogens increase as you go down the group

A
  • more electrons and so electron shells
  • this not only means the molecules are bigger but it also means there are stronger induced dipole dipole interactions between molecules which require more energy to overcome
40
Q

why does reactivity decrease down halogen group

A
  • wants to gain one electron
  • further up the group there are fewer electrons and therefore outer shells meaning the outer shell electrons are closer to the nucleus (smaller atomic radius with less shielding)
  • the outweighs the increase nuclear charge down the group
  • therefore due to a greater attraction there it is easier to gain an electron and so the reactivity increases up the group
41
Q

benefits and risks of choline

A
  • a toil gas which acts as a respiratory irritant in small concentrations and in large concentrations can be fatal
  • in drinking water it can react with organic hydrocarbons forming chlorinated hydrocarbons which are carcinogenic and can cause cancer
  • yet kills bacteria and microorganisms making water potable
  • the high risks associated with not purifying water which would otherwise be diseased (cholera and typhoid) far outweigh the relatively low rises of cancer
42
Q

why are halogens oxidising agents

A

they themselves are reduced (gains an electron) causing another species to be oxidised

43
Q

halogen displacement reactions

A
  • a solution of each halogen is added to aqueous solutions of the other two halides
  • if the halogen is more reactive than the halide present a reaction takes place, stalking the halide from the solution, changing its colour
  • standard equation (Br2 + 2KCl- –> Br2 + 2KCl-)
44
Q

how can you tell iodine and bromine apart in a solution

A

-can appear a similar brown/orange solution depending on the concentration
-add an organic non-polar solvent like cyclo-hexane can be added and the mixture shaken
-the non-polar halogens dissolve more readily in this than water making their colours easier to tell apart
(iodine becomes violet)

45
Q

disproportionation reaction

A

a redox reaction where the same species is both oxidised and reduced

46
Q

how is the reaction of chlorine with water a disproportionation reaction?

A

-for each molecule of chlorine, one atom is oxidised and the other reduced
Cl2 (aq) + H20 (l) –> HClO (aq) + HCl (aq)
reduced (-1) to form HCl
oxidised (+1) to form HClO

47
Q

why is chlorine used as a disinfectant

A
  • kills bacteria
  • both products are acids when reacted with water
  • the bacteria are killed by chloric acid and chlorate ions rather than by the chlorine
  • add indicator and it will initially turn red due to the presence of an acid and then decolourise as the bleaching action if chloric acid takes affect
48
Q

what happens in the reaction of chlorine with aqueous NaOH

A
  • with cold, dilute NaOH
  • the reaction of chlorine with water is limited by the low solubility of chlorine in water
  • if the water contains dissolved NaOH much more chlorine is dissolved and another disproportionation reaction takes place
49
Q

reaction of chlorine with cold, dilute aqueous NaOH

A

Cl2 (aq) + 2NaOH (aq) –> NaClO (aq) + NaCl (aq) + H2O (l)
NaCl (-1) reduced
NaClO (+1) oxidation
-the resulting solution contains a large concentration of chlorate ClO- ions from the NaClO formed (this is used in bleach)

50
Q

carbonate test

A
  • add dilute nitric acid to a test tube
  • if you see bubbles form it is likely to gas Is carbon dioxide
  • to prove this bubble the gas through limewater (Ca(OH)2)
  • this should form a white precipitate of calcium carbonate turning the solution cloudy
51
Q

sulphate test

A
  • most sulphate’s are soluble in water but barium sulphate isn’t and forms a white precipitate - this is the basis of the test
  • add a few drops of nitric acid followed by barium nitrate solution
  • nitric acid removes contaminants like carbonate ions
52
Q

halide tests

A
  • most halides are soluble in water but silver halides are insoluble
  • aqueous silver ions react with aqueous halide ions to form precipitates of silver halides (Ag+ + X- –> AgX)
  • silver chloride is white precipitate
  • silver bromide is a cream precipitate
  • silver iodide is a yellow precipitate
53
Q

how to clarify halide tests

A

add aqueous ammonia to test the solubility of the precipitate

  • chlorine dissolves in dilute ammonia
  • bromine dissolves in concentrated ammonia
  • iodine doesn’t dissolve in either concentrated or dilute ammonia
54
Q

what should the sequence of anion test be?

A
  • carbonate
  • sulfate
  • halides
55
Q

why does the carbonate test come first

A
  • neither sulphate nor halide ions produce bubbles with dilute acid
  • this allows the test to be carried out without incorrect identification
  • barium carbonate is white and insoluble so sulphate test must come after a carbonate test a white precipitate will form giving the impression of a sulphate even if its a carbonate
56
Q

why does the halide gets come last

A
  • silver carbonate and silver sulphate are both insoluble in water and will so form precipitates
  • conduct it last to rule out these possibilities.
  • if you intent to test for halide or sulphate ions after carbonate use dilute HNO3 as sulphuric acid contains sulphate ions and HCl carbonate ions
  • use barium nitrate rather than chloride if after a sulphate test you wish to test for halide ions
57
Q

how to test for ammonium ions

A

-when heated together aqueous ammonium ions and aqueous hydroxide ions react to form ammonia gas
NH4+(aq) + OH- (aq)–> NH3(g) + H2O (l)

58
Q

actual test for ammonium ions

A
  • add sodium hydroxide to a solution of an ammonium ion
  • ammonium gas is produced (unlikely to see bubbles as soluble in water)
  • mixture is warmed and ammonia gas released
  • you are likely to be able to smell ammonia but if not test with moist PH indicator paper as ammonia is alkaline and will turn to blue
59
Q

quick fit apparatus

A
  • round bottom or pear shaped flask
  • receiver
  • screw tap adaptor
  • condenser
  • still head
60
Q

heating under reflux

A

common procedure used to prepare an organic liquid without boiling off the solvent, reactants or products
-heat under water bath or heating mantle to be carried out a fixed temperature and continually boiled preventing volatilise components from escaping. it also means there is no flame present adding an extra level of safety should apparatus crack

61
Q

first three stages of heating under reflux

A
  • before fitting the condenser add anti bumming granules to the flask so the content will boil smoothly. otherwise this may cause large bubbles to form and the glassware to vibrate violently
  • apply a thing layer of grease to the ground glass joint on the condenser
  • place the condenser carefully in the flask and gently rotate the condenser back and forth to provide a good seal and ensures apparatus comes apart easily at the end
62
Q

last three stages of heating under reflux

A

clamp the condenser gently as to not jack the outer glassware

  • don’t put a stopper on the top as this will create a closed system building pressure and causing air to expand
  • rubber tubing is used to connect the inlet of the condenser to the tsp snd the outlet of the sink
  • water always enters the condenser at the bottom and leaves at the top to ensure the outer jacket is full
  • the vapour form the mixture rises up the inner tube of the condenser until it meets the outer jack with cold water causing it to condense
63
Q

distillate apparatus

A
  • round bottom of pear shaped flask
  • condenser
  • rubber tubing
  • heat source
  • screw cap adapter
  • receiver adapter
  • still head
  • themometer
64
Q

first three stages of distillation set up

A
  • the flask is clamped by its neck and the still head is connected to the flask (t shaped with two ground glass joints) one to fit the screw tap adapter and one the condenser
  • grease the joints so the apparatus comes apart easily at the end
  • second clamp is placed round the receiver adaptor at the point at which is is attached the condenser (remove need to clamp condenser as will be sufficiently supported)
  • rubber tubing is used to connect the inlet of the condenser to the tap and sink outlet. water always enters at the lowest point
65
Q

last three stages of distillation set up

A
  • a flask is used to collect the distillate so that the distillation apparatus is not completely airtight
  • flask is heated and moisture starts to boil, different liquids will have different boiling points
  • vapour moves out of the flask into other parts of the apparatus leaving behind the less volatile components of the mixture
  • when the vapour reaches the cold condenser then become a liquid which drips into the collecting flask
66
Q

purifying organic products

A
  • sometimes you get left with an aqueous or water layer and a separate organic layer
  • to identify which is the organic layer add some water to the mixture and the layer that gets bigger is the aqueous layer
67
Q

how to use a separating funnel

A
  • ensure tap is closed
  • pour the mixture of liquids into a separating funnel, place a stopper on top of the funnel and invert the content
  • allow the layers to settle
  • add some water to see which layer increases in volume (this is the aqueous layer)
  • place a conical flask until, remove the stopper and open tap to allow lower layer out
  • place a second conical flask under the separating funnel to collect the other layer
  • label the flasks
68
Q

how to remove named impurities

A
  • in preparation using acids impurities may come about. by adding aqueous sodium carbonate and shaking the mixture in the separating funnel
  • an acid present will react with carbonate to produce carbon dioxide. tap needs to be slowly opened, holding the stoppered separating funnel upside down to release any gas pressure that may be built up
  • finally the aqueous sodium carbonate layer is removed and the organic layer washes with water before running both layette off into two operate funnels
69
Q

drying the organic product

A

-add a drying agent (an anhydrous inorganic salt that readily takes up water to become hydrated)

70
Q

steps to dry an organic product

A
  • add the organic liquid to a conical flask
  • using a spatula add some drying agent to the liquid and gently swirl the contents of the mix together
  • place a stopper on the flask to prevent your product from evaporating away
  • if the solid has all stuck together in a lump there is still some water present so add some more drying agent until it disperse into fine powder
  • decent the liquid from the solid into another flask (should be clear)
71
Q

redistillation

A
  • some organic products have very similar boiling points so there may be some impurities
  • therefore clean and dry the apparatus and repeat the distillation process
  • only collect the product with the boiling point you are trying to make
  • the narrower the boiling range and the purer the product
72
Q

organic synthesis

A
  • the preparation of complex molecules from simple starting materials
  • this can be used by large pharmaceutical companies when testing new medicines
73
Q

target molecules

A

used to describe the compound the chemist is attempting to prepare
to do this you need to:
-identify the functional group in your starting and target molecules
-identify the intermediate the links these
-state the reagents and conditions needed for this