Module-3.1 Flashcards
Before the periodic tablle
-arostot;e beleived he world was mad eup of four elements - earth fire wate air and fire .
384-322 BCE .
-it could be argued that these four elemtns are similar to what we now call the sttes of mttter - solid liqui gs eith fire represenign more usunal phenomen sucha splasm .
PERIODIC TABLE 1. Antoinee 0 Laurent de Lvoiser
-1789
is now cisndired to be the first modern chemical textbook .
-in this he compiled the fist extensive list of elemnts , which he dusibed as ‘substances that coud not be broekn down further .@
WHAT was lavoisers list of elemnts
-Oxygen , nitrogen ,hydrogen , phosphours m meury zinc and sulfur ,.
-also disitnguished between metlss + non -metals ,
-unofrutntley it also included some compounds and micutr ,along witht emrs suchas ‘ligt; and ‘caloric’ heat which he belied to be material sustances .
- devleopmet of the periodic table
Jons Jakob Bezerlius
1828
-he published a table of atomic weights and determiend the compsotion by masso f many compounds .
-Berzelisus was also responsibel for introuvin lettter based symbosl for elements .
- development og the periodic tble Johann Wolfang Dobereiner
Dobereine rnoticed that certain groups of three leements .
-Ordered by tomic wiehgt would have a middle elemnts with a weeight and propeties (such s density) that were roughly n averga eof the other two elements .
THE TRIADS THAT Dobereine rconsdiered
-calcium strontium and brium .
-Chlorine bromine and iodien ,
-Litiium ,s odium and potassium
- development of the periodic table
-JOhn Newlands -devise a periodic table that had elemnts arraned in order of their relaitve aomic weihts .
-In1865 he suggested that rather than bein intraids , elements show simialr propeties tot eh elemens eight palces adfre it int he t bale .
-He called it his ‘law of octaves; .
5.development of the periodic table
dmitri Medeleev
-modern periodic table is based on theo ne publsihed by medeleev in 1869 .
-his table showed eleemnt otfrtrf by atomic weightd , similr to newlands and ordered periodically .
dmitiri edeelev (1)
-eleements withs imialr properies are rranged in verticals columns .
-gap were felt where no elements fitted the repeating patterns and the propeties of the missing elements were predicted .
-these missing elements hae eens ixned found nd atched .
elements Tht have been FOUND AND MATCHED MENDEELEV;S PREDICTION
Gallium , germainium , and scandium .
dmitiri mendeleev (2)
-the order of elements was rearranged where their properied did not fi .
-For e.g tellurium had a higher atomic weiht than iodine , but mendeelv rervesed them to make the properties fit with the res of th e table .
-The arrangeent of the element or groups of element in ord of their atomic weights corresponded to their os -capped valencies as well as thier distincitve chemical pepowperies
-Apparent series suchas , LI , Be , B , C, N , O and F.
6.development of the periodic table
1913-Henry Moseley determined the tomic number for ll the known elements .
-Moseley mdoifided MENDEELV’S peridoic law to red that the porpoeties of the lements vary periodically with rhier atomic numbers rather thn atomic weights .
-moseleys modifided periodic law put the elements tellurium and iodine in the correct order , a it did ofr argon nd potssiuma nd for cobalt nd nicle .
7.Development of the periodic table
-Seaborg dISCOVERED the transuranic elements fom 4 plutonium to 102 nobrlium ,
-He lso remodelled the periodic tbale by plcting the actnicide seirs belows the lanthanide series at the bottom of the table .
In the periodic tbale , elemnts are ordered by increasing atomic number .
The period number tells us the number of highest energy electron shell for the elemnts int hat period .
(e.g all the elements in period 4 , have the highest energy electron is 4 .)
KEY - there is a repeating trend in the propeties of the element across a period .
E.g —> period 2 , LEFT side elements are metals ,but on he right they re non - metals . This repeats to period three also .
Beteen the metals and non-metals we have the metalloides .
Metalloides have properties of oth metals and non-metls .
-Same tredn repeating across other periods . Scientiss calla pattern of repeatin trends PERIODICIT .
As we saw before ,e ch period represents a new highest energy electron she..
-Period 1 , highest energy electrons are in the first shell .
-period 2 , highest energy electrons are in shell 2
For each electron shell , electrons fill the s subshell before they fill the p subshell .
(check video if stuck )
KEY ; filling of electrons in subshells follows a periodic pattern .
-Within the highet energys hel , the s subshell fills beforore the p susbshell .
Each block in the periodic table is named after the sushell containign the highest energy electron for elements in that block .
Within a group , each elemnt has the same number of electrons in the outer subshell .
E.g , all the elemnt sin group 1 have one electron int heir outer s subshell .
As we move cross a period …
Each element has one more protoon in its nucleus than the element to its left .
We cannot measure the radius of an electron of an atom directly , as electron clouds , do not have a clear cut off point .
-Way of calculting atomic raidus it to look at the pair of idenitcal atoms , that have formed a bond .
-We take the atomic raidus as half the distant between the nuclei of the two atoms .
check sheet for atomic raidus of elemnts lithium–> calcium
ONE trend we can see is that ;
-Atomic radius decreases as we move across a period , from left to right .
look at period 2 we can see it decreasing from LITHIUM TO FLUROINE . and atomicrdius deceasing again from sodium tochlorine .
Going to explain why , the first trend happens using period 2 , but this applies to the other periods as well .
-As we move across a period (left to right ) , each element has one more proton in its nucleus than the elemnts before .
-Meaning that h POSITIVE CHARGE in the nucleus increases across a period ,
-Due to this , there is an INCREASED ATTRACTIOn , between the nucleus and the electrons .
-DRAWING the electrons CLOSER to the nucleus .
-Causing the atomic radius to decrease across the period .
Remember ; Outer shell electrons are partially shielded , fromt he attraction of the neuclues by electrons in innershells .
However , elements in period 2 , only have one inner electro shell .
-Meaning sheidling due to the inner elecron shell is the same across the period .
Second trend
-Atomic radius increses MOVING DOWN the group , Can see this , check sheet , with Lithium , Sodium and Potassium .
-As we move down group 1 , atomic radius increases (same trend with group 2 ) .
-The number of electron shells increases ,a s we move down the grup .
-So the outer electron shell is furhter from the nucleus .
-Each element has one more , full inner electron shell .
–ncreasing the amount of shielding between the nueclues and the otuer electron .
-Therefore ,t here is less atraction between the outer electron and the nueclues .
-Therefore ,t he atomic radius INCREASES , moving down a group .
Remember the defenition of first ionisation energy
the energy needed to remove 1 electron from each atom of an element in 1 mole of gaseous atoms, to form 1 mole of gaseous ions with a +1 charge.
BECAUSE OF THIS
-Atomic radius decreases across a period .
-Botht he increased nueclear charge and decreased atomic radius .decreases across a period .
Botht he increased nucler chagre nd decreased atomic radius means that the outer electrons are more attrcted to the nucleus .
-Causing 1st ionisation energy to increase across a period .
check shee for table
-How first ionisation energy varies across
-Plotted period 2 st ionisation energy against the atomic number of each element.
As you can see , the first ionisation energy ,t ends to icnrese as we move acros a period from elft to right .
-Moving across a period , the postive charge in the nucelus increase ,a s the number of protons increases .
-This increases the atraction between the nucelus and the electrons .
In all of these elemnts , we are remoivng on eelectron fromt he second electron shell .
Meaning the sheilding effect due to the inner electron shell is the sme for ehc eleemn .
-Boron oxygen do not fit the patern of increasing first ionisation energyu .
-Need to look t subshells Li+E removing electron from 2S SUBSHELL .
-However , for Boron , outer electron is int he 2P subshel and 2P subshell has a higher energy than the 2S subshell .
-Meaniing , it takes less energy to take an otuer electron from Bboron comapred o BE .
-Which s why it has a lwoer st ionsatione nergy than Beryilium .
Ionisation energy decreases at Oxgen ..
-Look at the subshell one of the 2p orbitals contians a PAIR OF ELECTRONS .
-These electros repel eachother
This means it taes less energyt o remove one of these electrons .
-Than if the electrons were in sepeatre orbitals .
Which is why Oxgen is the s ionsiatione egry is less thn N .
-Simialr paern for priod 3 (ame dips ut were are looking at the 3rd electrons hell , NOT the second shell .)
Check the ionsiation energies for group 1
-Ionisation energy in group 1 ,
-as you can see first ionisation energy decreases as we go down group.
–THIS IS DUE TO TWO FACTROS
FACTOR ONE
-As you go down the group , atomic raidus increases .
-Meaning hte outer elecron shell is furhter way from the nucleus .
Factor two
-Going down the group , thenumer of intenrl eleectron shells also increase .
-Menaing there is more shiedling between nuclejus and the outer electorns .
-Both of these factors mens ,that theattration etwene te nucles nd outer eecorns decreases .
-Causing 1st ionsation energy to fall .
-ALTHOUGH , nuclear charg goes up dwont he grup it is oFFSET , by the two actrs i mentioned above .
CHECK THE GRAPH
-sucessive iosniation energy for the elemnt oxygen .
two things to nocie
FACTOR ONE
CHECK star one on grah too
-In the case of oxygne , we cns ee a gradual INCREASE , in ionsition energy ,a s we remove the first SI electrons .
-THIS IS BECASUE ,e ach time we remove an elecrrn ,t he remianing electrons in the outershella re pulled LSIGHTLY CLOSER ot the nucleus .
-Menaing htwre is a grter attraction , between the outer electrons and hte nucleus csuing ioinsiiton energy to GRADUALL INCREASE .
factor two
-check star two
-THere is a MASSIVE increase in ionsiationenergyw ehn we remvoe the sixth electron .
-can be eexplaiend b lookign t the electronsi no the ocygen atom
-The first six elecrons are removed in the SECOND ELECTRONS HELL .
-Once we have removed these elctron the 7TH ELECTRON IS REMOOVED FROM THE FIRST ELECTRON SHELL .
STAR TWO (2)
-compared to the econd electron shell , the first electron shell is clsoer to the neuclues and electrons int he first shell exoerience mcuh LESS HIELIDNG .
-Meaning electrons int he first shell , have a greater attaction to the nueclues comapred to the second hsell.
-EXPALINING why ionisaiton energy is MUCH GREATER for the 7th and 8th electrons ocmpared to the 6th .
EXAM - Could be asked to use first ionisation energy data to identify an element .
-First ionisation energyies for an elemnt in period 3 . (sShwoing not all jsut the first 6) .
Wokrout number of elecrons int he outer shell .
-STAR ONE - as you can see , ionisation enegry gradually increases up to the FOURTH ionisation .
-Form star two , massively increases whenw e get to ionisation number five .
-Telling us that the element has electrons in the outer hell .
-The fifth electron must hae been removed from an INTENRAL SHELL ,w hich is wh y its ionsition energy is MUCH GREATER .
As this electron hs foru electrons in the outer shell .
Period 4 group 4 the element is sillicon .
We find metals on the etside and non-metlas on the rightsid e.
-we find metallidus between them which are also known as SEMI METALS.
Lithium (2)
Because the lithiuma otms hve doanted thier outer electron we now describe them as CATIONS .
-tHE NEGATIVE delocalised electrons are strongly attrcted to postive cations by electrostatic attraction .
-This electrostatic attraction is metallic bdoning .
Period 2
LHS we ahve emtals LI and BE
Lithium ATOMS have three protons int he nucleus , and two electrons in the shell and 1 electron on t he outer shell .
-in metals m electrons in the outershella re DECOLAISED (in other words they are shared ).
What is the overall structure of a metal ?
GIANT METALLIC LATTICE
One key thing about giant metallic stucture
cations are fixed in place and nnot move
second thing about giant metallic structure
delocalsied electrons are free to move .
-As the deleoclaised electrons can move freely , this explains an important property of metals .
-MRTLS ARE GOOD CONDUCTORS OF ELECTRICITY ,W HEN THEYA RE BOTH SOLIDS AND LIQUIDS .
Imagine a voltage is applied to a metal .
-As youc an see , the deloclaised electrons are atracted to the psotive pole and move towards it .
-Int his case , the delocalised elecros are mobile charge carries , enabling metals to conduct electricity .
thirdmetallic strucutre
Aonther porepty i most metals hve relatively high melting and boilign points .
-THIS IS DUE TO THE STRENGTH OF THE METALLIC BOND .
rememebr - DELOCALISED ELECTRONS HAVE A STRONG ELECTROSTAIC ATTRACTION TO THE CATIONS .
–iT TAKES A LOT OF ENERGY TO OVERCOME THIS ATTRACTIONS SO METALS GENERaLY HAVE A HIGH Mleting point and boiling point .
Berryilium in period 2
Beryllium atoms have two outer electrons which are delocalised ,
-So in Be , the strength of the metallic bond is greater than LITHIUm .
-Therefore , beryllium has a higher melting and boiling point than lithium .
fourth thing about giant metallic strcutres
-They do not dissolve .
-When we dd a metal to water –. it reacts with the water rather than dissolving .
check grpah , showing the apporcimae melting points of the elemnts in period 2 .
Lithium and Berillyium have a high melting point due to their metallic bodning in their giant metalluc latice .
-Next Boron mettaaloid and non metal carbob , ahve a very high melting point becaue theyf ormg ian covalent strcutres .
Giant covalent strucutres
-Billions of atoms are joined by strong covalent bonds + together these atoms form giant covalent lattice .
-Takes a great deal o energy to break all the covalent bonds ina giant covalent lattice .
-Explaining why covalent strucuture have a high melting point nad oig point .
-boron forms a number of gian covalent strucutres .
Giant covalent strucutres of carbon
Carbon , is in group4 ,t ehre are 4 electrons in the outershell , carbon can covalently bond to four other carbon atoms .
-By doing this , the carbon atoms have formed a giant covalnet lattice .
-Whent his takes palce ,c ron has formed dimond .
What is the shape and bond angle of diamond ?
In diamond , toms are arranged in a tetrahedral strcutre with a bond angle of 109.5 degrees .
one key property of diamond
high melting point and boiling point , due to large amounts of energy required to break the covalent bonds .
second property of diamond
it does nto conduct lectricty , as every electron is in a covalent bodn .
-There are no delcolcised electrons to ct as charge carriers.
final property of diamond
-giant covalent lattices are insoluble .
-This is because , solvents cannot disrupt the largenumber o strong covalnet bonds .
-instead of forming 4 covalnet bonds i can form 3 and form graphiete .
bond angles and shape of graphite
graphite forms layers of planar hexagonal strctures and has a bond angle of 120 degrees .
-One electron from each carbon atom is delocalised can act as a mobile charge carrier.
-Because of these delcoalised electrons , graphite is a good conductor of electricity.
single layer of graphite is GRAPHENE
-due to its delocalsied electrons , graphene is also a good conducotr of electicity.
check star oen of graph
-after carbon , the metling points of the elemnts in period 2 , drop very sharply .
n2 , o2 , f2 , all consits of diaomtic moelcules and are gases a rtp .
Ne -> is a noble gas consisiting of indivual neon atom .
if we cool these elements down to solid form ,t hey form a simple molecular lattice , with wek inermolecular forces between the moelcules .
-These weak inermolecular forces don’t requrie a lto of energy to break ,w hich is why these elemnts have low mp nd bp .
-so the changes in melting points cross period 2 is due to strcuutre and bonding .
A similar pattern occurs in period 3
Giant metallic strucutre , giant covalent and simple moelcular
phosphour is p4 and sulfur is s8 which are solids at rtp .
First five elements in group 2 check sheet , what are they referred to as ?
Elements in group 2 are referred to
As the alkaline earth metals
Key about alkaline earth metal s
-Alkaline earth metals are reactive .
-All of the elements on groip 2 have 2 electrons in the s sub shell
Of their outer sheel
When group 2 elements react , these 2 outeexteons are lost .
By transferring these two electrons
To
Other species . Group 2 elements act as reducing agents .
Check sheet for magnesium reaction
When magnesium reacts —> it looses two electrons to form Mg2+
-As it has lost two electrons and therefore is oxidised .
Showing these two electrons being transferred to a different chemical species
Check sheet
-Soecies (x) accepted electrons (x) so reduced .
Reactivity if group 2 elements INCREASES as you go down group. 2
Electron shells of magnesium .
-When group 2 elements react , they loose 2 outer electrons like Mg .
To remove an electron from an atom
Requires energy scientists call this
ionisation energy
Check sheet for magnesium first ionisation energy equation
As we are removing the first two electrons this involves ionisation energy .
Calcium is below magnesium
So therefore both the 1st and second ionisation energy for calcium is LESS than magnesium.
One reason for the card above
-Calcium has a greater atomic radius than magnesium . So the two outer electrons are further from the nucl es compared to magnesium.
Second reason
Calcium has one more inner electron shell than magnesoim .
As we have seen full inner electron shells PARTIALLY SCREEN the outer electrons .
- From the positive charge of the hue lies .So therefore outer electrons of calcium are LESS STROGMLH attracted to the hue lies Guam the outer electrons of Magnwsium.
-so
Calcium has a lower ionisation energy than mg
As if takes LESS ENERGY to remove the outer electron calcium is
More reactive .
This trend in ionisation energy explains why elements get more reactive down group 2.
REDO. REACTIONS OF GROUP 2-with oxygen and water
Oxygen water +dilute acids for
Pcr spec
First 5 elements of group 2 as we saw here group 2 elements have two electrons in their outer electrons in their outer s sub shell .
-They loose these electrons when they react and from 2+mefal
Ions so is oxidised .
These two electrons can fhn be used to reduce another. Idk pal .
-as reactivity increases as you go down the group (already said why check previous cards )
For ice sos. Only heed to magnesium to barium group 2
…
Chekc sheet for reaction of magnesium with oxygen
-If we head magnesium in air , then the magnesium reacts with oxygen producing a BRIGHT WHIFE LIHJT .
-Exakple of a redox reaction look at oxidation number oxisqfok number of one atom ONLY regardless of how many there is .
Before reacfion , both magnesium and oxygen have oxidation number of zero . As they are elements .
During the reaction each magnesium
Atom
Is oxidised loosing two electrons.
-so oxidation number is from 0 to +2 .
-each oxygen atoms is reduced gaining 2 electrons .
So oxidation number goes from
0 to -2 .
Key : changes in oxidation number must balance
Total change in oxidation number is +4 for the two amg atoms and _4 for oxygen .
We can also react group 1 elements with water
- check sheet for equation
Magnesium reacts very SLOWLY with cold water .
-As we move down group 2 the elements react more rapidly .
Reacting group 2 elements with water produces the
Alkaline metal hydroxide and hydrogen gas .
At the start calcium had oxidation £7’bet as 0 or is an element .
-Hydfirgen ions has an oxidation. +1
During the reaction the calcium is oxidised loosing two electrons .
-oxidation number of calcium from 0-+2
Z as-Tne two electrons are transferred in two hydrogen atoms forming the hydrogen gas .
-Tbese two hydrogen atoms were reduced at oxidation number +1-0
Oxidation numbers are wqual
Why did we ignore calcium hydroxide in the reaction
The hfwobhdidgen atoms have an oxidation of +1
So the oxidation. Number of the hydrogen atoms Ajax not
Hanged during the reaction
Ocr spec only need dilute acids
Check wlsheet for reacfion
When group 2 eelements react with dilute acids we make a metal salt and hydrogen gas .
Check the sheet for how I wrote that out
Being lazy 😀😀
We can see the same pattern with the same reactions
1.magneosum
And nitric acid
2. Magnesium and hydrochloride scods
Don’t need to know
Chlorine
… but it’s on the sheet ahyways so
Reactivity of group 2
I group. 2 elements as you go down the group. .
Magneosum reacts extremely slowly with water
Bradoum
Decays eapildy
Another what of making a group 2 hydroxide
Is to react a group 2 oxide with water
Check sheet
KEY : as the calcium hydroxide form, it dissolves in the water to form calcium hydroxide solution .
However , calcium gyrixude is only slightly soluble in water . So solution becomes saturated VETS WUIXKLY
Check sheet
-as we contribute to form calcium hydroxide . There is no length dissolves but I stwa sit confines to form solid cal iymbysteo.2
Key : one imprison trend we need to learn
Magnesium hydroxide extend leg low solubility in water .
Barium hydroxide relatively soluble in water .
-The solubility of group 2 hydroxide increased as you go down group 2 .
This trend is important as it explains alkalinity of group 2 hydroxides
Equation on sheet z
-When group 2 hydroxides dissolve in water , they release a metal ion snd two hydroxide ions .
Star one
The aqueous hydroxide ions make Yh e solution ALKALINE
-Concentrafiom of the aqueous hydroxide ion depends on the alkalinity or the solution.
High concentration of hydroxide ions - high ph
Magneosum hydroxide - extremely soluble Ph10 low concentrafiom of hydroxide ions .
- ph of this solution therefore is ph10
Barium hydroxide
Relatively soluble in water.
Relatively hydrogen ion concentration so ph is high
Can use group 2 compounds in neutralisation reactions in agriculture .
E.g1 solid calcium hydroxide also known as lime
-offen spread in fields .
Calcium hydroxides neutralises acids on the so .Making soil more fertile idk chekc cod .
Equation for caoh and general acid.
Check equation
-in reaction we produce water +salg depends on the acid
Another use of group 2 compounds
-treats indigestion . Indigestion often caused by excess hcl in the stomach this can treated I n two ways
Treatment one check equation
A suspension of magnesium
Hydroxide is called milk of magnesium .
-This neutralises hydrochloride acid to produce magnesium chloride and water .
Second treatment check sheet for equation
Alternatively tablets of calcium carbonate can be -see here
-calcium
Carbonate neutralise hcl acid to produce the products (check the sheet ).
Why don’t we use Calcoum hydroxide
As the alkalinity of calcium hudrpcide woild be harmful to body tissues , example lining the throat .
Group 7
Halogens —> all of the members lf
Group 7 form diatomic molecules example fluorine
At RTM , fluorine is a pale yellow gas and chlorine is a pale green gas .
Bromine is a red-brown liquid and Isidore is a grey -black solis
Why don’t we look at astatine ?
Aaa stains is highly radioactive and never has been observed
What is the trend in melting point and boiling point down group 7
They both increase -why ?
If we look down the halogens they form
A simple molecular lattice .
-the covalent bond between the halogen atom is strong .
-However , between the halogen molecules there are induced doodle dipole interactions Aka London forces .
-These intermolecular forces are relatively weak and do not take a lot of energy to break .
What is the strength of London forces determined by ?
The number of electrons .
-Iodine has 108 electrons whereas fluorine has 18
-So the London forces between iodine molecules require more energy to break that those between fluorine molecules .
Expwin why my wnd booking point increase down g 7
Trend in electronegativity
Two compounds contain halogen
Halogens are electronegative elements (bonding )
Electronegativity is the ability of an atom to attract the bonding electrons in a co a leng bond
Fluorine atom
Is
More electronegative than hydrogen ar
.
-Meaning the electron paid in the covalent bond is more attracted to the fluorine atom that the hydrogen atom …
…
Due to this
The electron pair in the covalent bond SHIFTS SLIGHTLY towards the fluorine atom .
-Making hydrogen fluoride a polar
Molecule . With the fluorine atom having a partially negative charge .
-As we mode down the group 7 , halogens are LESS ELECTRONEGATIVE
Bromine has a greater nuclear charge than fluorine .
However , bromine also has. Greater atomic. Radius so the outer electrons in bromine are further from
The nucleus .
Bromine has. Greater juexluds charge than fluorine however , bromine also has a greater afomicmfadisu. So the outer electron in bromine are further from
The nucleus
Bromine also has more
Electronnsuells fuan fluorine .
-Thess internal electron shielding the bonding electrons from the nuclear charge .
So together, the atomic radius and increased number of electron shells reduce the attraction between the bonding electrons in the nucleus .
-This means moving down group 7 elements become less electronegative .
Reactivi
UP 7 keu
Reactivity of group 7
When halogens react , they can remove an electron from
Another species .
-in this case xx I am showing the other species as a negative electron :
-chlorine atom removing an electron from x meaning x is oxidised
The chlorine atom is acting as an oxidising agent . As chlorine atom has GAINED an electron
It is being reduced to form the chlorine ion .
-chlorine ion now has the electron configuration of the noble gas argon .
Key halogens exist as diatomic molecules
-so when we write half equations we have to write as , check sheet
Fluorine to iodine arrow
Going UP THE GROUP the elements are more powerful oxidising agents )iodine is the least powerful oxidising agent and fluorine is the most powerful oxidising agent.
Look at electronic structure of halogens
-fluorine
When fluorine acts as an oxidising agent , the fluorine atoms removed another electron from its speed .
Tjis electrojbsdds into the outer shell of the fluorine atom to form the fluoride ion
Bromine gains an electron
Compared to fluorine br is a less powerful oxidising agent WHY ?
Reason 1
Bromine has a greater atomic radius than fluorine .
-meaning the outer electrons in fluorine are further from the nucleus
Reason 2
Bromine has more inner electron shells so greater shielding between the nueclued and outer electrons
For therese reasons there is LESS ATTRACTION BETWEEN ATHE NUCLEUS AND OUTER ELECTRONS making
It harder for the bromine atom to gain an electron from another species compared to flirpine .
So bromine is less reactive than fluorine and a less powerful oxidising agent .
In a displacement reaction halogen I’m surplus solution and react with an aqueous solution of a metal halide
Chlorine is MORE reactive then bromine so displaces it so we form bromine and sodium chloride .
Chlorine is a more powerful oxidising agent than chlorine so has a greater ability to remove an electron than bromine .
Chlorine 0 as it is an element
Be equals minus one in the reaction, each chlorine atom, removes small electron from a bromide ion
Chlorine Nat oxidising the bromide ions are the same chlorine, atoms being reduced to chloride ions.
At the end bromine has an oxidation of 0 citizen element.
Chloride ions, each other oxidation, number of minus one.
Check sheet for changes in sprinting
Reaction takes place as chlorine is a more powerful oxidising agent than bromine .
Similar reaction between cl +Naz
We get a similar reaction with bromine and iodine as bromine is a more powerful oxidising agent than iodine
Check sheet for equafokn
-why haven’t I include the metal ool for the redox reacfion
As it is a spectator ion sk Dosent take part in the reaped reaction
(Although o used sodium halides in example ANY OTHER GROUP 1 metal Halides would apply
When we carry out a displacement reaction , we produce a halogen
Halogens are coloured elements
- meaning displacement reactions produce a colour change .
PROBLEM : in aqueous solutions , bromine and iodine can appear orange and brown
How do we solve this
Add NON-09”-; organic solutions example cyclohexane
-cyclohexane forms an upper layer and does not take part in the reacfion
However , the halogens formed dissolve in the cyclohexane allowing us to see the colours of the halogens more clearly
In organic solvents , bromine appears orange and iodine appears violet .
Looking a fraction between chlorine and sodium bromide
At the start of the reaction, chlorine solution appears Pilgreen reaction forms bromine, which is orange.
When adding cyclohexane the Roman dissolves in the upper layer, which turns orange.
If we react more and so do I, then we form iodine
In Aqueous solution, Irene appears brown.
However, in the cyclohexane layer, iodine appears violet.
Reacting to the sodium iodide forms iodine
Disappears, brown and accuracy solutions, violet in the cyclohexane.
Chlorine is a widely used halogen
widewidly used halogenOne of the uses is in drinking water .
-Small amouts of chloirne is aded to drinking water to kill harmful bacter ia .
Check sheet for the euation
REACTION BETWEEN CHLORINE AND WATER .
Chlorine is a widely used halogen
-one of the uses is in drinking water .
Small amounts of chlorine is added to drinking water to kill bacteria
Check the sheet for equation
Two productions are formed in the reaction between chlorine and wate f
Two products are formed : choroid acid and hydrochloride acid .
Caloric acid is a powerful oxidising agent and is responsible for killing bacteria
If we look at the oxidation number , we can see this is a redo. Reaction
Chlorine the oxidation number is 0 as it’s an element
In hcl the oxidation number for chlorine is -1 showing that it had been reduced
But in choice one acid fhe oxidation number is +1 showing it has been oxidised
What is a disproportionation reaction !
-a redox reaction , in which aroma of the same element are oxidised and reduced and this is called a disporprtionsfion reacfion .
By adding small unties to water dilpreet one acids kills bacteria .
Preventing disease such as cholera
-in conditions where chlorinated drinking water is not available these diseases can be FATAL
Risks to using chlorine
Chlorine is a toxic gas and has to be handles carefully at water treatment plants
-there is a possibility that chlorine and water can react with natrually occurring hydrocarbons ( decaying materials ) .The chlorinated hydrocarbons produced could increase the risk of cancer in humans . (Ting risks £ and health benefits of chlorinated water is is heavily Ty our weighted .
Chlorine has a relatively low solubility in water
So only a relatively low concentration of products choral
Reacfion between chlorine and cold dilute Aqupus sodium hydroxide
In thsi reacfion , a much greater elb of chlorine can reaction and again an example of a disproportion reacfion .
-chlorite (1) ion is a powerful oxidising game t.
The solution made in this react is used as household bleach .
Cold conditions violate (1) is sdbale
But hot conditions it will further decompose …?
Check images for tests of gases .
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CONSOLIDATION ON GRAPHENE
Grapevine froms interlocking hexagonal rings that make up a lattice only one atom thick
Making gelaheen strong and extremely light so versatile
Check page 125 consolidation for odidwrioj reactions
Consolidation jow to identify anions what do I use
Check page 126 for
Carbonate ions
Sulfate ions
Halide ions
What is the equation for carbonate ions and the method
Check page 126 for equation
Method
Add a foul Ite strong acid to suspected carbonate
Collect any gas formed and pass it through line water
Carbonate test positive test observations
Fizzing , colourless gas is produced
Gas turns limewater cloudy
Sulfate ions equation and method add dilute hydrochloride acid and barium chloride to the suspected sulfate
Positive redt observation
White pot of the barium league is produced
What order should the tests be carried out and whyy
Carbonate test
Sulfate test
Halide test
Barium ions for an episode of if you hadn’t already ordered our carbonate, and you wouldn’t know if the precipitate with this carbonate or barium sulphate.
Silver ions for inside of precipitate. If you hadn’t ordered sulphate, you wouldn’t know if it was silver or silver highlight.
Halide ions method and equation
Dissolve suspected mort.
Add an acre solution of silver nitrate.
The colour of any precipitate formed.
If the colour is half distinguish and acres ammonia (first dilute then find )
The solubility of the precipitate and in aqueous ammonia
Halide ions positive test observations
Silver chloride, white precipitate soluble in dilute, NH, three
– silver wide – cream, precipitate, solid concentrated, and H3 only.
– Silver – yellow precipitate, insoluble, and concentrated and H3.
How to
Test for ammonium ions equation and method
Add sodium hydroxide solution to the suspected ammonium compound and warm vent.
– test any glass of water with red litmus paper
Positive test observations forammonoum ions
Ammonia gas will turn madness play blue.
– Ammonia gas has a distinctive smell – ammonia gas is hazardous, so you should always do this with care.
