Module 2.2 Flashcards
(167 cards)
1
Q
A quick recap of GCSE
A
-Electrons exist in electron shells .
The first shell has two electrons .
-The second shell can hold up to eight electrons .
-As well as the third before we put electrons in the fourth shell .
-KEY ; as we move further from the nuecleus , the energy of the shells increases .
2
Q
-Each shell is given a number …
A
… scientists call this the prinipal quantum number or ~n”
-The first shell has the principal quantum number 1 .
-The second shell has a principal quantum numebr 2 .
-Third has 3
-The fourth ha four .
3
Q
How do we clculate the mac=ximum nmber of electrons that each shell can hold .
A
Maximum n.o of electrons = 2nsquared
CHECK SHEET FOR THE SHELLS .
4
Q
Key; The electrons in a shell are found in regions called atmomic orbitals .
So what are atomic orbitals ?
A
-Atmoic orbitals ; a region around the nucleus that can hold up to two electrons with opposite spins .
-You do need to know electrons can either have an up spin or down spin .
-When we have two electrons in the same orbital , THESE TWO ELECTRONS MUST HAVE OPPOSITE SPINS .
-WE CANNOT have two electrons both with an up spin or both with a down spin , must be OPPOSITE .
5
Q
GCSE RECAP AGAIN
A
-At GCSE we considered an electront obe a particle .
-HOWEVER , at A-elevel , an electron is considered to be a cloud of NEGATIVE CHARGE /
-The negaitve charge cloud has the SHAPE OF THE ORBITAL , occupied by the elctron
6
Q
-Scientists can never be certain of the exact LOACATION of an electron …
A
… however , the atomic orbital shows us a 95% probability of where an electron exists .
7
Q
There are several different types of atomic orbitals + each one has a letter
A
These are S,P,D,F
8
Q
S OrbitaL (CHECK SHEET AND TABLE ON SHEET )
A
-hAS A SPHERICAL SHAPE .
-EVERY ELECTRON SHELL CONTAINS A SINGLE s ORBITAL.
-YOU NEEED to know how to draw it so pracitseee.
9
Q
P ORBITAL check sheet and table once again.
A
There are three p orbitls . They are shaped a bit like a dumbell .
-EVery electron shell aparfrom the first one contains three p orbitals .
10
Q
D orbitals .
A
there are five d orbitals but these re complicted shapes we don;t need to know .
-every shell apart from shells 1 an 2 have five d orbitals .
11
Q
f orbitals
A
there are seven different f orbitals . wE FIND THESE IN SHELLS FOUR AND ABOUVE .
remember ; EACH ORBITAL CAN ONL HOLD A MAXIMUM OF TWO ELECTERONS . (add up all the electorns in each shell and check table )
12
Q
subshell
A
all the orbitals of the same tpe in the same shell .
13
Q
key;
A
different subshells have different energies .
-As we move AWAY from the nueclues , the energy of the subshells increases /
14
Q
in orderto show how electrons fill the orbtials , we need to follow three rules :
A
…
15
Q
rule 1
A
orbtials with the lowest energy are filled first .
16
Q
rule 2
A
we can have up to two electrons in the same orbitals ut they must have opposite spins .
17
Q
rule 3
A
if we have orbitaals wtht he same energy , then we have to put electrosn into invidual orbitals before we pir them .
-this is because electrons int he same orbital repel .
CHECK SHEET FOR EXAMPLES .
18
Q
sheck example on electron configuration part 2 - whats interesting about shell four .
A
-the energy of the 4s subshell is less than the energyof the 3d subshell .
-Meanign we fll in the 4s subshell befpre we start filling in the 3d subshell .
19
Q
check sheet - what is star one saying ?
A
check sheegt
20
Q
WHATA RE THE TWO EXCEPTIONS TO THE ELECTRON CONFIGURTIONR UELE .
A
CHROMIUM AND COPPER .
-iN BOTH cases 4s subshell contains only one electron eventhough there are electrons in the 3d subshell.
-This is because the 3d subshell is more sstable when it is either half full /completely full
21
Q
In the case of chromium ; by having only one elecron in the 4s subshell it can have a half full 3d subshell .
A
in the case of copper , by onl having one electron int he 4s subshell it cn have a compeltleyf ull 3d subshell .
22
Q
do more practise ont his !!
A
…
23
Q
CHECK SHEET
A
Scientists divide the periodic table into different blocks , each block is named fter the subshell containign the highest energy electron for the elemnts in that block .
24
Q
short hand configuration (sodium is example .
A
1.look at the position of sodium on the perioidc table .
25
stage two
looka t the noble gas in the period before sodium (neon=10 electrons).
26
stage 3
compare the elctron configurations compare of neon compred to sodium .
27
stage four
tje onner electrons hells of sodium ajve the same electron confiuguration as neon .
-Remeer the elctrons in the inner shell are NOT INvolved inc hemical reactions is ontl the ONES IN THE outerhsell (so s31 for soidum)
28
stage five
so thartheer than writing out the whole electron configruation of soidum , we can represent the elcetorns in the inner shel ike this .
-sodium (ne)3s1 --> shorthand . IN EXAM WRITE GFULL HAND UNESS ASKED TO .
29
Use manganese as an example .
-check sheet .
-dblock still showed in electron configuration eventhough not outershella s electrons in the dsusbhell are invovled in the chemical reaciton .
-ALWAYS loose fromt he 4s subshell first=
30
Why do elements react and bond together ?
-elements invovled in chemical reactions will often become more stable by combining with other elements or transferrign electons .
-The most stable and unreactive elemnts are the noble gases .
-these elements re stable because they contain a full outershell of electrons .
31
(full outer shell inerms of orbitals)
-eight electrons in the outer shell of a noble fas , are made up of two int he s orbital and two each in the three p -orbitals .
-during reactions , other elemnts iwll bond with a tendecey to acquire this noble gas conficuration .
-s, p bodning are used for bodnignsuualuf r firt 18 .
32
types of chemical bonding
-Chemical bonds are classified into three main tpes ; ionic , covalent and metallic .
-A compound is formed when atoms of different elemnts are alwas in same proptions (h20always has two hdyrogen atoms to one oxygena tom . )
33
Ionic bonding
-In general , ionic bdoning occurs in compound consiting of a metal and a non-metal .
-if w eimagie a bond forming between atoms ,e elctros are tansfered fromt h metal atom tot he non-metal atom to form opposiely charegd ions that attrct achother .
34
Covalent Bonding
-occurs in compounds consisting of two non-metals .
-If we imagine BOND FORMING BEWEEN ATOMS , ELECTRONS ARE SHARED Between the atoms and are atracted tothe nueclei of both bonded atoms .
-covalent bdonging will also usuallyr esult in the elemtn inovlved achieing noble gas configurtations .
35
metallic bonding
-occurs in metlas .
-electron are shared between ll the atoms .
36
Ionic bonding how it works
-electrons are transferred from the metal atom to the non-metal atom .
-opposielyc harged ions are formed ,w hich are bonded together by electrostatic atraction .
-the metal ion is positive ..
-non -metal ions is negative .
37
giant ionic lattices
-each ion id surroudded by oppositelyc hrged ions .
-These ions atract each other from all directions , forming a three-dimensional giant ionic latice .
38
giant ionic lattices properites
Each ion attracts oppositelc hrged ions from al directions ;
-Each ion is surrounded by oppositelyc harged ions .
-The ions attracts each othe r, forming giant ionci alttice .
-because giant ionic latices involve hueg amouts of electorstatic atreaction between ions ,t hry have unique propeties .
39
Properties of ionic compounds - High meltign and boling poitns
SOLID AT ROOM TEMPERTURE - a large moutn of energy is needed to break the strong electorstatic bonds that hold the oppositely charged ions together in the soidl lattic e.
therefore they ahve hig mp p .
40
electrical conductivit of ionic compounds
-IN SOLID ionic lttice
-ions re held in ficed postions and no ions can mvoe .
-The ionic compound does not conduct eleictiruu .
41
electrical conducitivity of ionic compounds (2)
-the solid lattice breaks down and the ions are free to move .
-The ionic compound is now conductor of electircy .
42
Ionic compounds - solubilitt
-an ioionc lttice dissolves (soluble) in polar solvents , such s wte r..
-Pola solvents contain substncs that have polar bonds .
-a polar bond occur between atoms that do not share electrons equally nd it reuslts in the atoms having very smallc ages .
43
ioncic compoounds - solubilit part twp
polar water moelcues beak downa niconic lttice by surroudnign each ion to form a solotuion .
-the slight changes within the polar substance are able to atracr the chargd ion in the ionci lticie .
44
Covalent bonding -shared pairs of electrosn
-the negativley chrged shared pair of electrons is ttrcted to the positive charge of btoh nuclei .
-This atrction overcomes the repulsion between the two postiecl chrged nuclei .
-the resulting attraction is the ocvlent bond tht hodls the two atoms together .
-twpo electrons are shared . This is in contrast to the trnsfer of ele ctrons thr result in an ionic bond .
45
average bond enthalpy
-not all covalent bonds are the same strength . Some are much stronger than others ,w hich means that more energyw ould be required to break them , e.g during a reaction .
46
How are atoms in a solid metal held together by metallic bonding ?
-in metallic bonding , the atoms are ionsied . Positive ions (cations) ocupy fixed positions in a lattice .
-the outer shell electrons are delocalsied - they are shared between all the atoms in the metallic strcuture .
-metal is held together by the attractions between all the positive ions + the negative elctrons - especially the delocalised electrons .
47
Whya re metals able to conduct electricty ?
-As they have delcoalised electrons can move freely anwayhere within the emtallic strcutrue .
-Allowing the metal to conduct electricity even in the solid state .
48
why are metals mallebale + ductile ?
-as the delocalised electrons are largely responsible for these properties .Because , they can move , the etallic strucutres has a degree of 'give' which allows atom or layers to slid past eachother .
49
why do metals (typically) hve high metlign boilign poitns .
The electrons are free to mvoe throughout the strucutre , but the positive ions remains where thrya re .
-The attrction between the positive ions +the negative delocalised elecgtrons are very strong .
- high tmeperture is needed to overome the emtallic bond and disoldged the iosn foromt heir rigiod positos in the littic e.
50
How does the volume + tomic radius of the metlas in groups 1,2,3 affect the bond strength + thus the melting points of the metls ?
-For each sucesive grop , element ahs a giant metalic lattice , the nuclear charge increase , as does the numebr of electrons in the outershell . As gthere are more electrons in as you from goup 1-3 , so most electrons to form covalent bond .
-There is a harp decrese icncresing meting points s the elmets hae simple moleculr strcutures .
-Each indivuals molecule is atrrced to other by relativelu wek intrrmoleculsr forces .
-Atomic radius decrese s you go down group 1,2,3
51
How is metallic bodning different to covalent + ionic bodning .
check squilaci metallic sheet .
52
Dative covalent bond ?
When atom uses a lone pair of electrons to form a covalent bonds , .
53
Rules for a dative covalen tbond
-the acceptor must be electron deficent .
-Avaialbe orbitals for the electrons to occupy .
54
first ionisation energy
ennergy needed to remove one mole of electrons from one mole of atoms in their gaseoust state to form one 1+ ions also in their gaseouss stae ..
55
Check magnesium
Mg ---> Mg+e-
-1 mole of magensium ions + turn into a gas .
-Includign state smbol has .
-Then take one electron from every atom to form one mole of 1+ magnesium ions in thier gasesous state . The neegry needed ot do this is first ionsiation energy .
56
Once we remove an electron , we can continue to remove lectrons to measure the ionsiaton energy each time .
Mg ---> Mg1++E-
57
Looking above we remoeved sanotehr electron this is second ionsiaton energy
the energy needed to remove on emole of electrons from one mole of 1+ ions in their gaseous state to form one mole of 2+ ions in their gaseous stae .
58
-can contine reomving electron and measuring their ionsiaorin energye ach itme
these are sucessive ionsiatoine enrgies .
59
Workout the iosniation energy
ITS THE SAME AS THE CHARGE PRODUCED
60
Key about iosniation energy ;t he elctrons in an tom are attracted to the positive portons int he nucleus .
-The GREATER the atraction between the outer electeons + the nucleus , the GREATER the ionsiatoin enegry ./
61
FIRST FACTOR THAT AFFECTS IONISAITONE ENRY ;
-The distance between the nucleous + outermost electorn (Atomic radius)
-AS ATOMIC RAIDUS ICNREASE S.
-THE FORCE OF ATTRaction between the positive nuleus + otuer electron decreses,
62
SECOND FACTOR THAT AFFEC IONSIATION NERGY .
Charge on the nueclues
-Eleectors are attracted to the positvley chagred portos in the nculeus
THE greter nummber of portons = greater force of atraction between outer electrons + nucles .
63
THIRD FACTOR THT AFFECTS IONSIATION ENERGY
SHEIDLIGN
-Electrons int he outer shell re repelled b electrons in the inner shell .
-Sheiidlings effect reduced the attaction ebtween the outer electrons and the nucleus .
64
How does sucessive iosniatoin tell us about how eelctrons are ARRANGED IN ATOMS ?
cHECK GRAPH star one
This is ocrgen rph
-graudal increase in ionsation energya s we have just removed th first isx electreons .
-AS we remove each otuer electrons , the outer shells are PULLED LSIGHTLHY CLSOER TO THE NUCLEUS .
-MEANING ,t here is a grater atteaction between the outer electron + nucleus causing IONISAITON ENERGY TO INCREASES GRADUALLy .
65
star two
-Massive increase in ionsiation energy ,w ehn we remove the SEVENTH ELECTROn .
-As compred to the second electron shell . THE
FIRST ELECTRON SHELL IS CLOSEER TO THE NCUELUS ND THE ELCTROS INT HE FIRST SHELL EPXERIECE MUCH LESS shielding .
-so electrons int he first hsell hav agreater ATTRACTIOn to the ncuelus comapred to secon shell .
-eexpalingign why ionsiation energy is GREATER FO R7TH 8 TH ELECTRONS COMAPED TO FIRST SIX
66
EXAM ; identiy eement using iosiatione energy data
-Given a grpah .
-First si ionsiatione nergyies , workout element .
67
1. workout the nmber of eleoctrons in outer shell .
-Ionsiation graudlal increases up to fourth ionistion then HAS A BIG ICNREAE AT FIVE .
-Mening elemnt has FOUR ELECTRONS ON ITS OUTER SHELL .
5th electron must have been removed from intenrll shell . WHICH IS WHY FIFTH IONSIAIOTN ENERGY IS GRETER THAN PERVIOUS ONES .
-sowe know its gorup four and silicon .
68
1st eleemnts in gorup 1
-first ionsiation energy decreases as you go down the group .
69
FIRST RESON FOR CRD 68
1.moving down grop 1 .
-ATOMIC RADIUS INCREADES , Meaning outer electron shell is FURTHER AWAY FROMT HE NCULEUS /
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SECOND REASON FOR CARD 68
-going forwn the group , the nmber of internal energy levels ALSO ICNREASES , MEANING THEREIS MOR SHEIDLING BETWEEN NUCLEUS AND OUTER ELECTROSN .
-meaning goigng downa group , the attraction between nucleus + outer elecrron decreases .
-CAUSING FIRST IONSIATION ENERGY TO FALL .
71
HOWEVER IN GROUP 1 , nuclear chagre icnrease as you go down a group
but this is offset by the two fcts met ntioned befor .
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IONSIATION energy cross a period chekc graph
-first ionsition energya gasint atomic number .
73
first observation across a period ;
-first ionsiatione nergyt ends to icnrese as we move acorss a period .
AS We move across a period , nueclear chagrge increases as THE NUMBER OF PROTOS INCCRESES .
-This icnreases the attraction between then ucleus and the elctron /
74
observation across a period (2)
-noth the incresing nuclear chrge and decreased atomic radius mens that the otuer electors are MORE ATTRCTED TO THE NCULEUS .
-Causing the ifrst ionsiatoinenergy to ICNREASE .
75
KEYA BOUT PERIOD 2
-in all of these elements we are remaining in an electron from the SAME ELECTRONS HELL (2nd electrons hell)
-Meniang the shielding efect due to the inner electron shell is SIMIALR TO EACH ELEEMTN ..
76
EXCEPRIOSN TO IOSNIATION ENERGYR ULE 1
-BORORN .
-Take a look at lithium an dberrilium there are in the s subshell .
-WHOEER BROWONn has its outer electron in the 2- dubshell , which is in HIGHER ENERGY than 2S , which emans it es ESIER to removle an outer electron from boron than brrilu, .
-WHCIHC WIS WHY BORON HA SA LOWER FIRST IONSIATION ENERGY THAN Berillyoi, .
77
SECOND EXCEPTION TO THE RULE IN PERIOD 2
take a looka t the ocnfigurtion of the elctrons in nitorena and ocyen .
-NITORGEN has each electron in sepre orbitals in 2- .
Whereas in oxygen there is one porbital which hs a PAIR OF ELECTRONS . chekc graph .
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SECOND EXCEPTION TO THE RULE IN PERIOD 2 (2)
-These electrons repel eavhother , mening it tkes LESS energy to remove one of these lectrons than if electrons were in SEPRATE ORITALS .
-sof irst ionsiion energus lower in oxygen thn in ntirogen DUE TO MUTAL REPULSION /
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period 2 ...
second energy level
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period 3
3rd subshellw2g
81
how scientists represent 3d molecules ;
methane
-cenral carbon atom surroudnedby four hydrogen atoms .
82
check the diagram
solidlines...
tells us ht ebond lies ont he plane of the screen or page
-so the hydrogena toms also lie ont he plane of the page.
83
the solid wedge
tells us hte bond is coming out of the plane of the page .
-Menaing tht the hydorgen atom si coming out of the plane of the opage .
84
the dotted wedge...
this bond is projecting bck behind the palne of the page .
-so the hdyrogen is behidnt he plane of the page .
85
going ot be using these bonds whenever we need to show three dimensional shape of a moelcule ...
...
86
in order to workout the hsapes of the moelcules ,w e need to looka t electron pair replusiont heory ...
-it sates that the shape of the moelcules is detemrined by t he electron pairs surrounding the CENRAL ATOM .
-Only referring tot he outer shell in thsic ase .
-Based on the fact pairs of electrols REPEL all other elecrron pairs .
-The elcrons there move AS FAR PPRT as possible to MINImise this replusion .
87
Referring to Pairs of electons ...
covlent bonds is a pair of electrons .
88
-showing on sheet a molecule of beryillium choloride .
-in this moelcul th cetrnal aoms is beryillium .
-this is covanelty bonded to two atoms of chlorine .
-meaning we have two electron pars around the cetnral aotm .
-these two electrons repel echother and move as fr apet as possible .
-the furhtest they can move apart is ina sraight line .
89
part 2 of beryillium hclroide .
-scienitsts ay this molecule has a LINEAR strcuure aka striaght line .
-the angle betweent hese bonds is 180 DEGREES
90
CO2 has two double bonds however iwe rea it as a single bond .
KEY ; of a cenntral atom hhas TWO BONDS or bonidng ares .
-then t will hve a LINER shape with a 180 bod ngle .
91
electron pairs with three covalent bonds
-repel mvoe apart as far as possble
check sheet
-int his case the bond rrnage htemsleves towards the points of traingle like this .
-The bond anlges between them is 120 defrees .
Scienits cll this hspae a "trigonl"as it is based on a triangle .
-molecule viewed fromthe sid e, youw ould see it FLAT.
-Sciensits call this"Planar" meaning flat .
-so the shae is claled TRIGONAL PLANANR .
92
we willALWAYS see this shapes im tlaking bout when you ahve a central atom
within a cerain about of bondng electrsouns around it (1-6) as LONG AS THERE IS NO LONE PAIRS OF ELECTRONS .
93
Central atom with four pairs of bodning electrons?
Tetrahedral moelcule
-good exmpale ; MEthane
-All thebond angles in a tetrahedral moelcule is 109.5 degree /
94
central atom with five pairs o bodning electrons arondit ?
example - phosphorous pentrachloride .
-Ino rder to minimise REPUSion - two of the bodning pairs move to OPPOSITE SIDES o the mocleus .
-The other three bonding pairs now take up a centrla position ling ont he same plane .
-Theys pread themlselevs out as far as hey can /
95
central atom with five has two bond angles to consider
-bonds pointing up nd downa re at 90 degrees to the central plane .
-wheres he angl between thebonds lying on the centrl plane is 120 degrees .
-This shape is called trigonal bipryamidal .
(trigonal ) becuase the three atoms ont hec entral palne are formigna triangle
-biprimidal as these formt wo pyrmaidal shapes with the other oms .
96
six bodning pars around the cetnral atom
exmpale - sulfurhexafluroide
-haing a bdonign pair baove + below the cetnral palnea r.
-four bdongin pirs lying ont he central plae .
-bonds poitnignt up and downa t 90 degrees to the cetnral plane .
-howeve rint hsic ase , the angle between the bonds lyoin on he cetnral plane is also 90 degrees .
97
SHAPESo moelcule tha re ions Crbonste CO32-
CHECK SHEET
-IN THIS ION THE CETNRAL CARBON ATOM SIS BONDEDTO THREE ATOMS OF OXYGEN .
-hAVE GOT TWO SINGLE BONDS DN ONE DOUBLE BOND (TREAT DOUBLE BOND LIKE A ISNGLEBOND)
-due to the fact theat there is threee bdonging areas around htis central atom .
-Ion forms a triogonal planar strucutr with BOND ANGLES OF 120.
98
why does a double bond ont chnge the shape of the moleucle ?
-as dative bonds behave int he same wayas regulr covalent bods
99
-lone pairs rpeel more strongly than bodning pairs .
-this extra repulsion decreses bond ngles by 2.5
100
looka t ammonia on sheet
-nitrgne atoms have a lone pair of elecrons
-so we have THREE bdongin apir nd one lone pair .
-We alrady know four pairs of electrons form a tetrahedral strcutre .
-so the shape of the ammonia moelucle is based on a tetrahedron .
*however a slone pairs repel more storngly and extra repulsion decresases bond ngle by 2.5 degrees
-THEBOND ANGLE IN THE AMMONIA MOELCUIE IS 107 degrees .
-Scienists call this hape PYRAMIDAL s it looks like a pyramid .
101
-Ammonia can react with hydorgen to make ammoniam NNH4
-int ihs cas e, the lone pair forms a dative covalent bond
-a DATIVE COVALENTBOND has the same LEVEL OF REPUSIONA S REGULAR COVALEN BOND .
-mENAING THE BOND ANGLE IN THE AMMONIM returns back tot eh tetrahedral 109.5 .
102
-Another example is ater
-moelcuel of water .
-oxygen has tw isngle covalen t bodn to hydrognaosm .
-But oxygen also has TWO LONE PAIRS OF ELECTONS .
-THE TWO LONE PAIRS MEANS WE HAVE four bodnign pairs in total
-TETRAHEDRAL BOND ANDLGE 109.5
ONE LONE PAIR REDUCES THE ANGLE BY 2.5
BUT WE HAVE TWO SO
-2 LONE PAIRS ANLGE REDUCES TO 10.5
sciensits call this a non-linear or vshaped moelcule .
103
check sheet electron pair geoemtieis as there is a lot more to know off by heart
dr squailaci ...
104
also check sheet as there is a tehcnique on how to owkrout the bodning
see squillaci 2.2
105
check electronengsvity sheet
-pair of electrons in the covalent bond in blck
-hydorgen meolcuele and hclorine molecule
-looking at these tw moleucles , we can see that the pair of electron sint he covalen tbond , lie exactl MIDWAY between the two neuclei .
106
Hydorgen Chloride MOLECULE
-In this moleuclue , the electrons in the covalent bond are MUCH CLOSER , to the chlorine nucleus than they hydogen nucleus .
-As chlorine atoms are MORE ELECTRONEGAtive than hydrogen atoms .
107
What is electronegavity
The abiolity of an atom to attract the pair of electrons in a covalent bond.
108
goign back tothe hydrogen oemlcule both atoms on either side of the covalent bond are HYDORGEN
-AS both of them are hydorgena toms , they have the same electronegavity .
-So the pair of electrons in the covalent bond are equally attracted the two hydrgoen nuclei (same with the chlorine moelcule )
109
HOWEVER
In the case of hydrogen chloride , the chlorine atom is MORE ELECTRONEGAVITYE , than the hydrogena toms .
-Menaing that the electron pair int he covlent bond are CLOSER to the chlorine nucleus .
110
-Electron negativities of the elements in period 1 , 2 3
-Noble gases are not incldued as theyare NOT REACTIVE .
-Based on the Pauling electronegativity scale
-with four bening the most electronegstive .
111
electronegavity changes in a periodic tbe;e
-increasing electroongegarivyt to the top right of the period tbale .
(increases across the period towards fluroien)
-increases up a group towards fluroine )
112
-What are the most electronegaivie elements
Fluorine , Oxygen , Nitorgen and Chlroine .
-As electronegativity is BASED ont he attrction between electron pairs and the NUCLEUS of the atom in quesiton .
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FIRST FACTOR THAT AFFECTS ELECTRONEGATITVIVTY .
-SIze of the positive charge on the nucleus .
-Moving left to right across a period ,t he number of protons int he nucleus increases .
-This INCREASED POSITIVE CHARGE , increases the ATTRCTIOn between then ucleus and the pir of electrons int he covalnet bonds .
-SO elements on the RIGHT of the periodic table are MORE ELECTRONEGATIVE than those ont he left .
114
SECOND FACTOR THAT AFFECTS ELECTRONEGATIVITY
-The SMALLER the atomic raidus , the closer the bonding electrons will be to the nucleus of ana tom .
-as we move left t oright across a period , atomic radius decreses .
-Mening elements on the right of the period table are MORE ELECTRONGEATIVE ,t hn thos eont he left .
115
THIRD FACTOR THAT AFFECTS ELECTORNEGTIVITY
-Shielding of then uelcues by electrons in the inner shells .
(example on page is fluroien and hclirne )
--Chloirne has ONE MORE inner electron shell than fluroien .
-ELctrons in the inner shell SCREEN electrons in the outer shell from the positive charge of the nucleus .
-SO the GREATER numbe rof inner shell s, the LOWER the electronegativity .
-So chlorine has a LOWER electrongetivyt than fluroine .
116
Looking at cl molecule as both atoms are chlorine , so both have the same electronegativity
-because of this the electron pair in the covalent bond , lies midway between the two nuclei .
-scientists call this a pure covalent bond
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Two ways to show a bond is polar
1.write delta positive and delta negative to show chargers
-delta means charge is small , as the electrons has only shifted towards the more electronegative atom .
-delta negative sign on note electronegative elements
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Another way to show bond polarity is to draw an arrow like in the sheets .
-arrow points towards more electronegative element .
-As, hydrogen chloride only has one bond and that bond is polar . this means hydrogen chloride molecule molecule hwe an overall polarity . Scientists call this overall polarity the DIPOLE MOMENT
- molecules with more than one bond ?
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Molecules woth more than one bond ?
-carbon dioxide has two carbon to oxygen doible bonds .
- both bonds are polar .
-shoe get these bonds poking in opposite directions in a straight line . Meaning the dipoles cancel out .
Therefore a carbon dioxide molecule has no overall polarity .
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Tetrchloromethane
Carbon to hydrogen is effectively non polar .
As carbon + hydrogen have very similar electronegativites .
-in this case , the dipoles on the carbon to chlorine bonds cannot cancel so tie is a polar molecule
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Trichloroethane
-the carbon to hydrogen bond is effectively non polar .
As carbon +hydrogen have very similar electronegatives .
-In this case , the dipoles on the carbon to chlorine bonds cannot cancel so this is a polar molecule
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simple molecular subtances
menaing they conists of relatively small moleucles .
-Each moleulces has a fixed number of atoms .
-Giant covalent strucutres - graphie - would not fall into this category .
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Why do simple molecular structures have a relatively low boiling point .
-example is hydorgen chloride .
-In hydorgem hchloride , the chlroine atom + hydorgen atoma re chemically bonded nby a covalne tbond .
-you need to remember a covalnet bond ---> extremelys STRONG AND ONLY NROKED in a chemical reacotn .
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covalen tbonf are not affectred t all when a substance boils .
...
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other types of bonding ; intermoelcular forces
Hydorgen chldprde mpeic;es --> moelcules ae raltivley clsoe together . So this oucld be in liqui form .
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-Intermolecular forces of attraction attrac inbetween moelcules - the world "inter2 means inbetwene ..
-intermoelcla rforces are mcuh weakee than covalen tbond s.
-are easily broken (by high temperatue
-when we heat a simplemoelcuar usbtacnes , this causes the moelcules to MOVE FASTER .
-At a certian temeprature , the intemrolcur force sbrek and llows the moelcules to move away frome achother .
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-f0or exmpale the liquid hydrgoen chlooirde is turing into a gas .
LIKE WE SW BEFORE
-simple moleuclar usbtacnes sucha shyforgen chlrodie , have low boiling poitns .
-As it takes vry little energyt o break the intermolecualr forces .
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THERE RE SEVERAL TYPES O INTERMOLEUCLAR FORCES
-induced dipole-dipole interactions (London forces dipseiersion forces)
-Permanent dipole -dipole interactoins .
-Hydorgen Bonds.
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intermolecular forces
are between molecules .
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key ; intermoleuclar forces are WEAKER than covalne tbonds
and are eaisly broken e.g by high temperature .
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What are induced dipole dipole interactions AKA
-London forces / dispersion forces .
-Also referred to as vander Waals forces .
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check sheet - there is three atoms of neeon . (this could apply to any atom or moelcule )
-Electrons are represented as negative signs .
-Electorns are moving randomly fro place to plce so , there is NO REASON AT ALL as to why electrons should be evenly spread out .
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Look at atom 1
-Imagine for a fraction of a second there were more electrons on the right hand side of the atoms then the left .
-NOw , for a fraction of a second , the aotm HAS A DIPOLE .
-On the left side of the atom 1 has a slight POSITIVE CHARGE .
-On the right side of atom 1 has a sligh NEGATIVE CHARGE
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Key; this is not a permanent dipole , it is just formed instantaneuolsy due to RANDOM ELECTRON MVOMENT .
Scieniststs call this an insanenous dipole because it happen at a PARTICULAR ISNANT .
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As the right side of the atom 1 has a slight negaice charge , this repels the elctrons in atom 2 .
-This causes the electrons ina tom 2 , to move towards the rightside of aotm 2 .
-Meaning atom 2 , onw has a DIPOLE .
-This is called an induced dipole The world "INDUECE" MEANS CAUSED BY SOEMTHING ELESE .
-tEHREFROE THE dipoel INA TOM 2 , DID NOT HPPEN RANODMLY , IT WAS CUASEDBY THE dipole in aotm 1 .
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NOw the negative charge on the irgh side of atom 2 repels the elecrons ina tom 3 . Causing the electrons ina om 3 to move ot the rightisde of atom 3 .
ATOM3 now has a dipole ,.
-This is an induced dipole as it ws caused by the dipole of otm 2 .
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-now all of these dipoles now have a force of attraction .
This is clled a LONDON force .
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key points baout london forces .
1.London forces are WEAK and ESIL BROKEN .
-ALL intermoleuclar forces are MUCH WEAKER than covalnet bonds .
2.-London forces are causedby random electron mvoment .
-Meaning every single atom or moelcule , will experience London forces . Even f they eperience OTHER INTERMOLEUCLAR FORCES AS WELL .
3.tHE STENGTH OF THE lONDON FORCED DPEEND ON THE NUMBER OF ELECTONR S.
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CHECK SHEET FOR BOILING POINT OF NOBLE GASES
-As ou can see , the toms with GREATER NUMBER of electrons hae a HIGHER boiling point .
-As they experience stronger london forces .
-This is the same forALKANES .
-
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Many molecules have a permamnet dipole - HCL .
-The chlorine atom has a greater electronegativty than the hydorgen atom .
-This means that the elecgtron pair in the covlent bond is MORE ttracted to the chlorine aotm , making the hclorine aotm ,s lgihtly NEGTIVE compred to the hydorgen
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Check sheet on permanent dipole - dipole interctions .
-When two moleucles of hydorgen chlroide get near enoguh , then their permanen tidpoles can lead to n attraction like this .
-this is an example o an intermoleuclar force .
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KEY only molecules with a permamne tdipole can experience , this type of intermolulcr ofrce .
check sheet for tetrchloromoethane and tetrchlordie .
-oth molecules contain polar bonds .
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-Tricholoromehane ha a permanent overal dipole aka dipole meomnt .
-However , int he case of tetrachlroomemthane , this molecule IS COMPLETEL SYMMETRICL .
-therefore , the bond polarities CANCEL OUT .
-so TETRACHLOROETHANE HAS NO overall PERMANENT DIPOLE .
-mEANING TRICHLOROMETHANE HAS PERMANENT DIPOLE-DIPOLE INTERACOTNS , whres TETRACHLROMETHN does not .
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key ; as trichloromethne has a permament dipole dipole interactions ,w e might epect it to he a higher boiling point than terchlromethane .
-However , this is not the case .
-In fact tetrachlroomethane has a hihger booilign point .
-TO expalint his ,w e need to remebermer permanet dipole dipole inerctions ar NOT THE ONLY ONLY FORCE DSCTING .
-As all the moelcules lso experience LODNON FORCES .
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The SIEZZ of the London FOrces dpened ont he NUMBER of electrons present .
tricholormehtnae = 58 electorns
tetrahclroide =74 electrons .
-London forces are STRONGER in tetrachlormethane , which is whyit has a HIGHER boling point .
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We can see the same effect in hydorgen halides
-All of these moelcues have a permnent dipole .
-So permamnet dipole dipole intrections are acting in all three cases.
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what is the effec we see in the table ?
-The strenght of the permanent dipole decreases . going DOWN the table .
-From hydrgoen chloride too hydorgen iodide .
-Due to the DECREASING electronegativyt of the halogens ,a s we move down group 7 /
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HOwever , boiling point increases as we go down the group ,w hy ?
-AS the number of electrons increases ,s o Lodnon forces increase as we go down the tble .
-Cusing the boiing points of the hydorgen halides ot INCREASE .
-
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Key; when looking at the pemrmnet dipole-dipole interactions
Always need to remember London forcs will also be acitng .
HYDORGEN FLUROIDE is not int he tbale as it is involded in HDYROGEN BODNING .
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Two key points about hydorgen bonding
(example in sheet is hdyorgen fluroide)
-In hydorgen fluroide ,t he fluroide ATOM HAS THREE LONE PAIRS of eectrons .
-Fluroine is the most electroegaive aelemnt ,s o this strongl attracts the pair of electrons in the covlnetbond .
-Meaning the moelcule is POLAR , with the hydrogen atom having a positivechagre ,
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two molecules of hydrogen fluroide
-posiive hydrogen on the rihgt side is attracted to the LONE PAIR OF ELECTRONS o the fluroide atom on the left hnd moleucle .
-*THis attraction is also known s a HYDROGEN BOND .
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Key; when drawing a hydorgen bond , it must be from the hydorgen atom DIRECTLY TO THE LONE PAIR OF ELECTRONS .
-hydorgen bodns are the STRONGEST type of intermolecular force .
-For hydogen bonding to take place we need to conditions .
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Two bonds to take hydrogenbonding
-hydorgena toms bonded to a strong electronegaitivy element .
-2. the electronegaitve atom must have at least oNE lone pair of electrons .
MEANING we find hydorgen bodning with lfuroine oxygen an nitorgen ./
-ALL THREE OF THESE ELEMENTS re strongl elelectronegtive and all three hae at least one lone pir electorn s.
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check sheet for
hydrogen bodning in water , ammonia
IN EXAMS YOU WOULD BE ASKED TO DRW THESE DIAGRAMS .
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-Hydrogen bonding in here a signifcan effect on the properies of molecules .
both compoinds have 2 hydorgen atoms covalbent bonds to em .
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WATER
-the hydoren atoms ae bodned to oxygen , WHICH IS HE SECONDmost electrongetive leemnts .
-Menaing wter undergoes hydorgen nodning .
-Therefore , whenw e boil water it needs a grea DEAL F EERGY to break the hydorgen bonds .
-Meaning it has a high boiling point of 100 degrees .
-Hydorgen bonds also mean water has a high ELTING POINT .
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hYDROGEN SUFLIFE
-hYDORGEN SFLIDE , THE HYDOGRWNA TOMS RE BONDED TO SULFUR .
-EECTRONGETIVITY OF SLUFUR IS VERY similar to hydorgen . Meaning hydrogen slifde cannot form hydrogen bonds .
-Hydorgen sulifde has a much lower boiling poitn of -60 degrees .
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-AS well AS HAVING A REatively high boiling point , water also has some unusal propertied due to hydorgen BODNING .
...
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ICEBERG ---> solid form of water (ice)
---> is less dense than the liquid form .
-Therefore ice floats on the surfaceof he water .
DUE TO THE HYDROGNE BONDS .
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Why is this
, as in liquid water molecules we close together and sometimes they are further apart .
-hydrogen bonds are constantly being broken and formed
-as we cool water down , water molecules move more slowly .As w reach the reach the freezing point of water (zeroooo degrees) , the water molecules arrange themselves into an ordeee structure ice .
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How does hydrogen bonds effect ice formation ?
-in the ice the water molecules are further apart than in liquid , making the ice less dense than liquid water .
-therefore I’ve , floats on the surface of the water .
-really important for organisms that live in water , As ice insulated the water below preventing the water from freezing completely
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ice is less dense than water
-i n almost all materials the solid is denser than the liquid , water is the excpetion
.
-wghen ice forms , water meocluels arrang themsleves into an ordely patttern and h ydrogen bonds form between the moclules (these iwll occur int he iquid pahse but not s often as moelcules cn move pas eachother and overcome hese bodns . )
-ice has an open lattice with hdyrgoen bonds holding the water moelcule apart .
-when ice melts ,t he rigid hydrogen bodns comllapse allowing h20 moelcules to move close together .
-so ice is lessdense thanw aer
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what is the denisy of ice lie using hydorgen bodns ?
-the ice lttic ei s an oepn netrok of water moelcules .
-ntoice how h hydrogen moelcues hold water moelculea part . it is ebvudr the hydrogen bpnds make ice tke up more space that itti s lss dense hnw ater . notice lso the hssape of the larric .
-oxyena tom bons two covalent two hydrogen hydrobem nonds re slighl stronger lmaooo .
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Why is an ionic lattice soluble in polar solvents
-polar solvents have substances that have polar bonds .
- Polar bond occurs when atoms do not share electrons equally between mean and it results in an atom having very small charges on them. Space space
– chemicals broke down ionic that is ceramic each end from solution. The slight charges within the protocol and are able to track the charter on the dark ironic Gladys. This means the letters is dropped out of it.
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Why is an ionic lattice soluble in polar solvents
-polar solvents have substances that have polar bonds .
- Polar bond occurs when atoms do not share electrons equally between mean and it results in an atom having very small charges on them. Space space
– chemicals broke down ionic that is ceramic each end from solution. The slight charges within the protocol and are able to track the charter on the dark ironic Gladys. This means the letters is dropped out of it.
