Edexcel​ ​IGCSE​ ​Chemistry
Topic​ ​1:​ ​Principles​ ​of​ ​chemistry
States​ ​of​ ​matter
Notes
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1.1
understand​ ​the​ ​three​ ​states​ ​of​ ​matter​ ​in​ ​terms​ ​of​ ​the​ ​arrangement,
movement​ ​and​ ​energy​ ​of​ ​the​ ​particles
● The​ ​three​ ​states​ ​of​ ​matter​ ​are​ ​solid,​ ​liquid​ ​and​ ​gas
● They​ ​can​ ​be​ ​represented​ ​by​ ​the​ ​simple​ ​model​ ​above,​ ​particles​ ​are​ ​represented
by​ ​small​ ​solid​ ​spheres
● Gas:​ ​particles​ ​have​ ​the​ ​most​ ​energy​ ​–​ ​shown​ ​by​ ​the​ ​diagram,​ ​as​ ​the​ ​particles​ ​are
the​ ​most​ ​spread​ ​apart​ ​with​ ​a​ ​random​ ​arrangement
● Liquid:​ ​particles​ ​have​ ​more​ ​energy​ ​than​ ​those​ ​in​ ​a​ ​solid,​ ​but​ ​less​ ​than​ ​those​ ​in​ ​a
gas​ ​and​ ​the​ ​particles​ ​are​ ​closer​ ​together​ ​but​ ​have​ ​a​ ​random​ ​arrangement
● solid​ ​has​ ​least​ ​energy​ ​–​ ​particles​ ​are​ ​not​ ​moving/are​ ​just​ ​vibrating​ ​and​ ​they​ ​are
arranged​ ​regularly​ ​and​ ​very​ ​closely​ ​together
1.2 understand​ ​the​ ​interconversions​ ​between​ ​the​ ​three​ ​states​ ​of​ ​matter
in​ ​terms​ ​of:​ ​the​ ​names​ ​of​ ​the​ ​interconversions,​ ​how​ ​they​ ​are
achieved,​ ​the​ ​changes​ ​in​ ​arrangement,​ ​movement​ ​and​ ​energy​ ​of​ ​the
particles
● Physical​ ​changes​ ​–​ ​therefore​ ​involves​ ​the​ ​forces​ ​between​ ​the​ ​particles​ ​of​ ​the
substances,​ ​instead​ ​of​ ​these​ ​interconversions​ ​being​ ​chemical​ ​changes
● Melting​ ​and​ ​freezing​ ​take​ ​place​ ​at​ ​the​ ​melting​ ​point:
o solid​ ​→​ ​liquid:​ ​melting
o liquid​ ​→​ ​solid:​ ​freezing
● Boiling​ ​and​ ​condensing​ ​take​ ​place​ ​at​ ​the​ ​boiling​ ​point:
o liquid​ ​→​ ​gas:​ ​boiling
o gas​ ​→​ ​liquid:​ ​condensing
● when​ ​you​ ​change​ ​from​ ​solid​ ​to​ ​liquid​ ​to​ ​gas:​ ​the​ ​particles​ ​gain​ ​more​ ​kinetic
energy,​ ​move​ ​around​ ​more​ ​and​ ​become​ ​more​ ​randomly​ ​arranged​ ​and​ ​further
apart
● when​ ​you​ ​change​ ​from​ ​gas​ ​to​ ​liquid​ ​to​ ​solid:​ ​the​ ​particles​ ​lose​ ​kinetic​ ​energy,
move​ ​less​ ​and​ ​become​ ​more​ ​regularly​ ​arranged​ ​and​ ​closer​ ​together
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understand​ ​how​ ​the​ ​results​ ​of​ ​experiments​ ​involving​ ​the​ ​dilution​ ​of
coloured​ ​solutions​ ​and​ ​diffusion​ ​of​ ​gases​ ​can​ ​be​ ​explained
1.3
● Diffusion
o Movement​ ​of​ ​particles​ ​from​ ​an​ ​area​ ​of​ ​high​ ​concentration​ ​to​ ​an​ ​area​ ​of
low​ ​concentration
o For​ ​this​ ​to​ ​work,​ ​particles​ ​must​ ​be​ ​able​ ​to​ ​move
▪ Therefore,​ ​diffusion​ ​does​ ​not​ ​occur​ ​in​ ​solids,​ ​since​ ​the​ ​particles
cannot​ ​move​ ​from​ ​place​ ​to​ ​place​ ​(only​ ​vibrate)
o Therefore,​ ​coloured​ ​solutions​ ​are​ ​diluted​ ​by​ ​adding​ ​water,​ ​because​ ​the
particles​ ​of​ ​the​ ​colour​ ​diffuse​ ​to​ ​the​ ​air​ ​of​ ​low​ ​concentration,​ ​mixing​ ​with
the​ ​water​ ​molecules,​ ​causing​ ​dilution​ ​to​ ​occur
know​ ​what​ ​is​ ​meant​ ​by​ ​the​ ​terms:​ ​solvent,​ ​solute,​ ​solution​ ​and
saturated​ ​solution
1.4
●
●
●
●
1.5
Solvent​ ​=​ ​liquid​ ​in​ ​which​ ​a​ ​solute​ ​dissolves
Solute​ ​=​ ​substance​ ​that​ ​dissolves​ ​in​ ​a​ ​liquid​ ​to​ ​form​ ​a​ ​solution
Solution​ ​=​ ​mixture​ ​formed​ ​when​ ​a​ ​solute​ ​has​ ​dissolved​ ​in​ ​a​ ​solvent
Saturated​ ​solution​ ​=​ ​solution​ ​in​ ​which​ ​no​ ​more​ ​solvent​ ​can​ ​be​ ​dissolved
(chemistry​ ​only)​ ​know​ ​what​ ​is​ ​meant​ ​by​ ​the​ ​term​ ​solubility​ ​in​ ​the
units​ ​g​ ​per​ ​100​ ​g​ ​of​ ​solvent
● Solubility​ ​is​ ​shown​ ​as​ ​the​ ​grams​ ​of​ ​a​ ​solute​ ​that​ ​will​ ​dissolve​ ​in​ ​100​ ​g​ ​of​ ​water
1.6
(chemistry​ ​only)​ ​understand​ ​how​ ​to​ ​plot​ ​and​ ​interpret​ ​solubility
curves
● generally:
○ solubility​ ​of​ ​solids​ ​increases​ ​when
temperature​ ​increases
○ solubility​ ​of​ ​gases​ ​increases​ ​when
pressure​ ​increases
○ any​ ​mass​ ​below​ ​the​ ​line​ ​for​ ​a​ ​solute​ ​at​ ​a
specific​ ​temperature​ ​would​ ​mean​ ​the
solution​ ​was​ ​unsaturated
○ any​ ​mass​ ​above​ ​the​ ​line​ ​for​ ​a​ ​solute​ ​at​ ​a
specific​ ​temperature​ ​would​ ​mean​ ​the
solution​ ​was​ ​supersaturated​ ​and​ ​unstable
1.7
(chemistry​ ​only)​ ​practical:​ ​investigate​ ​the
solubility​ ​of​ ​a​ ​solid​ ​in​ ​water​ ​at​ ​a​ ​specific​ ​temperature
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Edexcel​ ​IGCSE​ ​Chemistry
Topic​ ​1:​ ​Principles​ ​of​ ​chemistry
Elements,​ ​compounds​ ​and​ ​mixtures
Notes
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1.8​ ​understand​ ​how​ ​to​ ​classify​ ​a​ ​substance​ ​as​ ​an​ ​element,​ ​compound​ ​or
mixture
● Element​ ​=​ ​substance​ ​made​ ​from​ ​only​ ​one​ ​type​ ​of​ ​atom
● Compound​ ​=​ ​substance​ ​made​ ​from​ ​two​ ​or​ ​more​ ​elements​ ​that​ ​have​ ​reacted
chemically​ ​with​ ​each​ ​other
● A​ ​mixture:
o Consists​ ​of​ ​2​ ​or​ ​more​ ​elements​ ​or​ ​compounds​ ​not​ ​chemically​ ​combined
together
o Chemical​ ​properties​ ​of​ ​each​ ​substance​ ​in​ ​the​ ​mixture​ ​are​ ​unchanged
1.9​ ​understand​ ​that​ ​a​ ​pure​ ​substance​ ​has​ ​a​ ​fixed​ ​melting​ ​and​ ​boiling​ ​point,
but​ ​that​ ​a​ ​mixture​ ​may​ ​melt​ ​or​ ​boil​ ​over​ ​a​ ​range​ ​of​ ​temperatures
● A​ ​pure​ ​substance​ ​=​ ​a​ ​single​ ​element​ ​or​ ​compound,​ ​not​ ​mixed​ ​with​ ​any​ ​other
substance
● In​ ​everyday​ ​language,​ ​a​ ​pure​ ​substance​ ​=​ ​substance​ ​that​ ​has​ ​had​ ​nothing​ ​added
to​ ​it,​ ​so​ ​it​ ​is​ ​unadulterated​ ​and​ ​in​ ​its​ ​natural​ ​state,​ ​e.g.​ ​pure​ ​milk
● Pure​ ​substances​ ​melt​ ​and​ ​boil​ ​at​ ​specific​ ​temperatures
o This​ ​melting​ ​and​ ​boiling​ ​points​ ​data​ ​can​ ​be​ ​used​ ​to​ ​distinguish​ ​pure
substances​ ​from​ ​mixtures​ ​(which​ ​melt​ ​over​ ​a​ ​range​ ​of​ ​temperatures​ ​due
to​ ​them​ ​consisting​ ​of​ ​2​ ​or​ ​more​ ​elements​ ​or​ ​compounds)
1.10​ ​describe​ ​these​ ​experimental​ ​techniques​ ​for​ ​the​ ​separation​ ​of​ ​mixtures:
simple​ ​distillation,​ ​fractional​ ​distillation,​ ​filtration,​ ​crystallisation,​ ​paper
chromatography
● Simple​ ​distillation
o Used​ ​to​ ​separate​ ​a​ ​pure​ ​liquid​ ​from​ ​a​ ​mixture​ ​of​ ​liquids
▪ Works​ ​when​ ​the​ ​liquids​ ​have​ ​different​ ​boiling​ ​points
▪ Commonly​ ​used​ ​to​ ​separate​ ​ethanol​ ​from​ ​water
▪ (Taking​ ​the​ ​example​ ​of​ ​ethanol…)​ ​ethanol​ ​has​ ​a​ ​lower​ ​bp​ ​than
water​ ​so​ ​it​ ​evaporates​ ​first.​ ​The​ ​ethanol​ ​vapour​ ​is​ ​then​ ​cooled​ ​and
condensed​ ​inside​ ​the​ ​condenser​ ​to​ ​form​ ​a​ ​pure​ ​liquid.
▪ Sequence​ ​of​ ​events​ ​in​ ​distillation​ ​is​ ​as​ ​follows:​ ​heating​ ​->
evaporating​ ​->​ ​cooling​ ​->​ ​condensing
● Fractional​ ​distillation
o The​ ​oil​ ​is​ ​heated​ ​in​ ​the​ ​fractionating​ ​column​ ​and​ ​the​ ​oil​ ​evaporates​ ​and
condenses​ ​at​ ​a​ ​number​ ​of​ ​different​ ​temperatures.
o The​ ​many​ ​hydrocarbons​ ​in​ ​crude​ ​oil​ ​can​ ​be​ ​separated​ ​into​ ​fractions​ ​each
of​ ​which​ ​contains​ ​molecules​ ​with​ ​a​ ​similar​ ​number​ ​of​ ​carbon​ ​atoms
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o The​ ​fractionating​ ​column​ ​works​ ​continuously,​ ​heated​ ​crude​ ​oil​ ​is​ ​piped​ ​in
at​ ​the​ ​bottom.​ ​The​ ​vaporised​ ​oil​ ​rises​ ​up​ ​the​ ​column​ ​and​ ​the​ ​various
fractions​ ​are​ ​constantly​ ​tapped​ ​off​ ​at​ ​the​ ​different​ ​levels​ ​where​ ​they
condense.
o The​ ​fractions​ ​can​ ​be​ ​processed​ ​to​ ​produce​ ​fuels​ ​and​ ​feedstock​ ​for​ ​the
petrochemical​ ​industry.
● Filtration
o If​ ​you​ ​have​ ​produced​ ​e.g.​ ​a​ ​precipitate​ ​(which​ ​is​ ​an​ ​insoluble​ ​salt),​ ​you
would​ ​want​ ​to​ ​separate​ ​the​ ​salt/precipitate​ ​from​ ​the​ ​salt​ ​solution.
▪ You​ ​would​ ​do​ ​this​ ​by​ ​filtering​ ​the​ ​solution,​ ​leaving​ ​behind​ ​the
precipitate
● Crystallisation
o If​ ​you​ ​were​ ​to​ ​have​ ​produced​ ​a​ ​soluble​ ​salt​ ​and​ ​you​ ​wanted​ ​to​ ​separate
this​ ​salt​ ​from​ ​the​ ​solution​ ​that​ ​it​ ​was​ ​dissolved​ ​in
▪ You​ ​would​ ​first​ ​warm​ ​the​ ​solution​ ​in​ ​an​ ​open​ ​container,​ ​allowing
the​ ​solvent​ ​to​ ​evaporate,​ ​leaving​ ​a​ ​saturated​ ​solution
▪ Allow​ ​this​ ​solution​ ​to​ ​cool
▪ The​ ​solid​ ​will​ ​come​ ​out​ ​of​ ​the​ ​solution​ ​and​ ​crystals​ ​will​ ​start​ ​to
grow,​ ​these​ ​can​ ​then​ ​be​ ​collected​ ​and​ ​allowed​ ​to​ ​dry
● Paper​ ​chromatography
o Chromatography…
▪ Used​ ​to​ ​separate​ ​mixtures​ ​and​ ​give​ ​information​ ​to​ ​help​ ​identify
substances
▪ Involves​ ​a​ ​stationary​ ​phase​ ​and​ ​a​ ​mobile​ ​phase
▪ Separation​ ​depends​ ​on​ ​the​ ​distribution​ ​of​ ​substances​ ​between
the​ ​phases
Paper​ ​Chromatography
Analytical technique separating compounds by their
relative​ ​speeds​ ​in​ ​a​ ​solvent​ ​as​ ​it​ ​spreads​ ​through​ ​paper.
The more soluble a substance is, the further up the paper
it​ ​travels.
Pigment
Separates​ ​different​ ​pigments​ ​in​ ​a​ ​coloured​ ​substance.
Solid,​ ​coloured​ ​substance
1.11​ ​understand​ ​how​ ​a​ ​chromatogram​ ​provides​ ​information​ ​about​ ​the
composition​ ​of​ ​a​ ​mixture
● see​ ​1.10-​ ​separates​ ​mixture​ ​into​ ​individual​ ​components,​ ​so​ ​reveals​ ​number​ ​of
components​ ​in​ ​mixture​ ​and​ ​these​ ​components​ ​can​ ​be​ ​identified​ ​using​ ​Rf​ ​values
● Compounds​ ​in​ ​a​ ​mixture​ ​may​ ​separate​ ​into​ ​different​ ​spots​ ​depending​ ​on​ ​the
solvent​ ​but​ ​a​ ​pure​ ​compound​ ​will​ ​produce​ ​a​ ​single​ ​spot​ ​in​ ​all​ ​solvents
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1.12​ ​understand​ ​how​ ​to​ ​use​ ​the​ ​calculation​ ​of​ ​Rf​ ​values​ ​to​ ​identify​ ​the
components​ ​of​ ​a​ ​mixture
● Rf​ ​value​ ​=​ ​distance​ ​moved​ ​by​ ​substance​ ​/​ ​distance​ ​moved​ ​by​ ​solvent​ ​(​ ​/
represents​ ​a​ ​dividing​ ​sign)
▪ Different​ ​compounds​ ​have​ ​different​ ​Rf​ ​values​ ​in​ ​different
solvents,​ ​which​ ​can​ ​be​ ​used​ ​to​ ​help​ ​identify​ ​the​ ​compounds
1.13​ ​practical:​ ​investigate​ ​paper​ ​chromatography​ ​using​ ​inks/food
colourings
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Edexcel​ ​IGCSE​ ​Chemistry
Topic​ ​1:​ ​Principles​ ​of​ ​chemistry
Atomic​ ​structure
Notes
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1.14​ ​know​ ​what​ ​is​ ​meant​ ​by​ ​the​ ​terms​ ​atom​ ​and​ ​molecule
● All​ ​substances​ ​are​ ​made​ ​of​ ​atoms
● A​ ​substance​ ​with​ ​only​ ​one​ ​sort​ ​of​ ​atom​ ​=​ ​element
o An​ ​atom​ ​is​ ​the​ ​smallest​ ​piece​ ​of​ ​an​ ​element​ ​that​ ​can​ ​exist
● A​ ​molecule​ ​=​ ​formed​ ​when​ ​atoms​ ​join​ ​together​ ​by​ ​chemical​ ​bonds​ ​(can​ ​be​ ​made
of​ ​atoms​ ​of​ ​the​ ​same​ ​element)
1.15​ ​know​ ​that​ ​the​ ​structure​ ​of​ ​an​ ​atom​ ​in​ ​terms​ ​of​ ​the​ ​positions,​ ​relative
masses​ ​and​ ​relative​ ​charges​ ​of​ ​sub-atomic​ ​particles
subatomic​ ​particle
relative​ ​mass
relative​ ​charge
position
proton
1
+1
in​ ​the​ ​nucleus
neutron
1
0
in​ ​the​ ​nucleus
electron
1/1836
-1
in​ ​shells​ ​around
nucleus
1.16​ ​know​ ​what​ ​is​ ​meant​ ​by​ ​the​ ​terms​ ​atomic​ ​number,​ ​mass​ ​number,
isotopes​ ​and​ ​relative​ ​atomic​ ​mass​ ​(Ar)
● Atomic​ ​(proton)​ ​Number​ ​=​ ​number​ ​of​ ​protons​ ​(=​ ​number​ ​of​ ​electrons​ ​if​ ​it’s​ ​an
atom,​ ​because​ ​atoms​ ​are​ ​neutral)
● Mass​ ​(nucleon)​ ​Number​ ​=​ ​number​ ​of​ ​protons​ ​+​ ​neutrons
● Isotopes​ ​=​ ​different​ ​atoms​ ​of​ ​the​ ​same​ ​element​ ​containing​ ​the​ ​same​ ​number​ ​of
protons​ ​but​ ​different​ ​numbers​ ​of​ ​neutrons​ ​in​ ​their​ ​nuclei
● Relative​ ​atomic​ ​mass​ ​(of​ ​an​ ​element)​ ​=​ ​an​ ​average​ ​value​ ​that​ ​takes​ ​account​ ​of
the​ ​abundance​ ​of​ ​the​ ​isotopes​ ​of​ ​the​ ​element
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1.17​ ​be​ ​able​ ​to​ ​calculate​ ​the​ ​relative​ ​atomic​ ​mass​ ​of​ ​an​ ​element​ ​(Ar)​ ​from
isotopic​ ​abundances
e.g.
A​ ​sample​ ​of​ ​chlorine​ ​gas​ ​is​ ​a​ ​mixture​ ​of​ ​2​ ​isotopes,​ ​chlorine-35​ ​and
chlorine-37.​ ​These​ ​isotopes​ ​occur​ ​in​ ​specific​ ​proportions​ ​in​ ​the​ ​sample​ ​i.e.
75%​ ​chlorine-35​ ​and​ ​25%​ ​chlorine-37.​ ​Calculate​ ​the​ ​R.A.M.​ ​of​ ​chlorine​ ​in
the​ ​sample.
The​ ​average​ ​mass,​ ​or​ ​R.A.M.​ ​of​ ​chlorine​ ​can​ ​be​ ​calculated​ ​using​ ​the
following​ ​equation:
R.A.M.
=
​ ​ ​=
​ ​ ​(mass​ ​of​ ​isotope-A ​ ​x​ ​ ​%​ ​of​ ​isotope-A)​ ​+​ ​(mass​ ​of​ ​isotope-B​ ​ ​x​ ​ ​%​ ​of​ ​isotope-B)
​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​100
(35​ ​x​ ​75)​ ​+​ ​(37​ ​x​ ​25)
​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​100
=
R.A.M.
=
3550
100
35.5
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Edexcel​ ​IGCSE​ ​Chemistry
Topic​ ​1:​ ​Principles​ ​of​ ​chemistry
The​ ​Periodic​ ​Table
Notes
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1.18​ ​understand​ ​how​ ​elements​ ​are​ ​arranged​ ​in​ ​the​ ​Periodic​ ​Table:​ ​in​ ​order
of​ ​atomic​ ​number,​ ​in​ ​groups​ ​and​ ​periods
● Elements​ ​are​ ​arranged​ ​in​ ​order​ ​of​ ​atomic​ ​(proton)​ ​number​ ​(bottom​ ​number)​ ​and
so​ ​that​ ​elements​ ​with​ ​similar​ ​properties​ ​are​ ​in​ ​columns,​ ​known​ ​as​ ​groups.
● Elements​ ​in​ ​the​ ​same​ ​periodic​ ​group​ ​have​ ​the​ ​same​ ​amount​ ​of​ ​electrons​ ​in​ ​their
outer​ ​shell,​ ​which​ ​gives​ ​them​ ​similar​ ​chemical​ ​properties.
● elements​ ​with​ ​the​ ​same​ ​number​ ​of​ ​shells​ ​of​ ​electrons​ ​are​ ​arranged​ ​in​ ​rows
called​ ​periods
1.19​ ​understand​ ​how​ ​to​ ​deduce​ ​the​ ​electronic​ ​configurations​ ​of​ ​the​ ​first​ ​20
elements​ ​from​ ​their​ ​positions​ ​in​ ​the​ ​Periodic​ ​Table
● the​ ​electronic​ ​configuration​ ​of​ ​an​ ​element​ ​tells​ ​you​ ​how​ ​many​ ​electrons​ ​are​ ​in
each​ ​shell​ ​around​ ​an​ ​electron’s​ ​nucleus
● for​ ​example,​ ​sodium​ ​has​ ​11​ ​electrons:​ ​2​ ​in​ ​its​ ​most​ ​inner​ ​shell,​ ​then​ ​8,​ ​then​ ​1​ ​in
its​ ​outermost​ ​shell.
o you​ ​can​ ​represent​ ​sodium’s​ ​electronic​ ​configuration​ ​as:​ ​2.8.1
● remember-​ ​electrons​ ​fill​ ​the​ ​shells​ ​closer​ ​to​ ​the​ ​nucleus​ ​before​ ​filling​ ​any​ ​further
out.​ ​1st​ ​shell​ ​holds​ ​2​ ​electrons,​ ​2nd​ ​and​ ​3rd​ ​hold​ ​8
1.20​ ​understand​ ​how​ ​to​ ​use​ ​electrical​ ​conductivity​ ​and​ ​the​ ​acid-base
character​ ​of​ ​oxides​ ​to​ ​classify​ ​elements​ ​as​ ​metals​ ​or​ ​non-metals
●
●
●
●
Metals​ ​are​ ​generally​ ​conductive​ ​(of​ ​electricity)
Non​ ​metals​ ​(excluding​ ​graphite)​ ​are​ ​not​ ​conductive
If​ ​an​ ​element​ ​is​ ​conductive​ ​and​ ​its​ ​oxide​ ​is​ ​basic​ ​then​ ​the​ ​element​ ​is​ ​a​ ​metal
If​ ​an​ ​element​ ​is​ ​not​ ​conductive​ ​and​ ​its​ ​oxide​ ​is​ ​acidic​ ​then​ ​it’s​ ​a​ ​non​ ​metal
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1.21​ ​identify​ ​an​ ​element​ ​as​ ​a​ ​metal​ ​or​ ​a​ ​non-metal​ ​according​ ​to​ ​its​ ​position
in​ ​the​ ​Periodic​ ​Table
● Metals​ ​=​ ​elements​ ​that​ ​react​ ​to​ ​form​ ​positive​ ​ions.
o Majority​ ​of​ ​elements​ ​are​ ​metals.
o Found​ ​to​ ​the​ ​left​ ​and​ ​towards​ ​the​ ​bottom​ ​of​ ​the​ ​periodic​ ​table.
● Non-metals​ ​=​ ​elements​ ​that​ ​do​ ​not​ ​form​ ​positive​ ​ions.
o Found​ ​towards​ ​the​ ​right​ ​and​ ​top​ ​of​ ​the​ ​periodic​ ​table
● divide​ ​can​ ​be​ ​seen​ ​by​ ​the​ ​red​ ​line​ ​in​ ​the​ ​periodic​ ​table​ ​at​ ​the​ ​top
1.22​ ​understand​ ​how​ ​the​ ​electronic​ ​configuration​ ​of​ ​a​ ​main​ ​group
element​ ​is​ ​related​ ​to​ ​its​ ​position​ ​in​ ​the​ ​Periodic​ ​Table
● group​ ​number:​ ​gives​ ​number​ ​of​ ​electrons​ ​in​ ​outer​ ​shell​ ​e.g.​ ​group​ ​3​ ​has​ ​3
electrons​ ​in​ ​outer​ ​shell
● period​ ​number:​ ​gives​ ​number​ ​of​ ​electron​ ​shells​ ​e.g.​ ​period​ ​1​ ​has​ ​1​ ​shell​ ​of
electrons
1.23​ ​understand​ ​why​ ​elements​ ​in​ ​the​ ​same​ ​group​ ​of​ ​the​ ​Periodic​ ​Table​ ​have
similar​ ​chemical​ ​properties
● number​ ​of​ ​electrons​ ​in​ ​outer​ ​shell​ ​is​ ​responsible​ ​for​ ​the​ ​way​ ​different​ ​elements
react
● this​ ​means​ ​elements​ ​with​ ​the​ ​same​ ​number​ ​of​ ​electrons​ ​in​ ​the​ ​outer​ ​shell​ ​will
undergo​ ​similar​ ​reactions
● therefore​ ​elements​ ​in​ ​the​ ​same​ ​group​ ​have​ ​similar​ ​chemical​ ​properties
1.24​ ​understand​ ​why​ ​the​ ​noble​ ​gases​ ​(Group​ ​0)​ ​do​ ​not​ ​readily​ ​react
● They​ ​have​ ​8​ ​electrons​ ​in​ ​their​ ​outer​ ​shell​ ​(except
helium,​ ​which​ ​has​ ​2).
● They​ ​are​ ​unreactive​ ​and​ ​do​ ​not​ ​easily​ ​form​ ​molecules,
because​ ​they​ ​have​ ​a​ ​stable​ ​arrangement​ ​of​ ​electrons.
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Edexcel​ ​IGCSE​ ​Chemistry
Topic​ ​1:​ ​Principles​ ​of​ ​chemistry
Chemical​ ​formulae,​ ​equations​ ​and​ ​calculations
Notes
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1.25​ ​write​ ​word​ ​equations​ ​and​ ​balanced​ ​chemical​ ​equations​ ​(including​ ​state
symbols):​ ​for​ ​reactions​ ​studied​ ​in​ ​this​ ​specification,​ ​for​ ​unfamiliar​ ​reactions
where​ ​suitable​ ​information​ ​is​ ​provided
● (g)​ ​means​ ​gas,​ ​(s)​ ​means​ ​solid,​ ​(l)​ ​means​ ​liquid,​ ​(aq)​ ​means​ ​aqueous
● Example​ ​of​ ​word​ ​equation:​ ​hydrochloric​ ​acid​ ​+​ ​sodium​ ​hydroxide​ ​->​ ​sodium
chloride​ ​+​ ​water
● Example​ ​of​ ​balanced​ ​chemical​ ​equation:​ ​HCl​ ​+​ ​NaOH​ ​->​ ​NaCl​ ​+​ ​H​2​O
● to​ ​balance​ ​an​ ​equation:​ ​you​ ​need​ ​to​ ​make​ ​sure​ ​there​ ​are​ ​the​ ​same​ ​number​ ​of
each​ ​element​ ​on​ ​each​ ​side​ ​of​ ​the​ ​equation​ ​and​ ​if​ ​there​ ​isn’t​ ​use​ ​big​ ​numbers​ ​at
the​ ​front​ ​of​ ​a​ ​compound​ ​to​ ​balance​ ​it​ ​e.g.​ ​3H​2​O
1.26​ ​calculate​ ​relative​ ​formula​ ​masses​ ​(including​ ​relative​ ​molecular​ ​masses)
(Mr)​ ​from​ ​relative​ ​atomic​ ​masses​ ​(Ar)
● Relative​ ​formula​ ​mass​ ​(Mr)​ ​of​ ​a​ ​compound:​ ​sum​ ​of​ ​the​ ​relative​ ​atomic​ ​masses​ ​of
the​ ​atoms​ ​in​ ​the​ ​numbers​ ​shown​ ​in​ ​the​ ​formula
● In​ ​a​ ​balanced​ ​chemical​ ​equation:
sum​ ​of​ ​Mr​ ​of​ ​reactants​ ​in​ ​quantities​ ​shown​ ​=​ ​sum​ ​of​ ​Mr​ ​of​ ​products​ ​in​ ​quantities
shown
1.27​ ​know​ ​that​ ​the​ ​mole​ ​(mol)​ ​is​ ​the​ ​unit​ ​for​ ​the​ ​amount​ ​of​ ​a​ ​substance
● Chemical​ ​amounts​ ​are​ ​measured​ ​in​ ​moles​ ​(therefore​ ​it​ ​is​ ​the​ ​amount​ ​of
substance).​ ​The​ ​symbol​ ​for​ ​the​ ​unit​ ​mole​ ​is​ ​mol.
● The​ ​mass​ ​of​ ​one​ ​mole​ ​of​ ​a​ ​substance​ ​in​ ​grams​ ​is​ ​numerically​ ​equal​ ​to​ ​its​ ​relative
formula​ ​mass.
● For​ ​example,​ ​the​ ​Ar​ ​of​ ​Iron​ ​is​ ​56,​ ​so​ ​one​ ​mole​ ​of​ ​iron​ ​weighs​ ​56g.
● The​ ​Mr​ ​of​ ​nitrogen​ ​gas​ ​(N​2​)​ ​is​ ​28​ ​(2x14),​ ​so​ ​one​ ​mole​ ​is​ ​28g.
● One​ ​mole​ ​of​ ​a​ ​substance​ ​contains​ ​the​ ​same​ ​number​ ​of​ ​the​ ​stated​ ​particles,
atoms,​ ​molecules​ ​or​ ​ions​ ​as​ ​one​ ​mole​ ​of​ ​any​ ​other​ ​substance
1.28​ ​understand​ ​how​ ​to​ ​carry​ ​out​ ​calculations​ ​involving​ ​amount​ ​of
substance,​ ​relative​ ​atomic​ ​mass​ ​(Ar)​ ​and​ ​relative​ ​formula​ ​mass​ ​(Mr)
● You​ ​can​ ​convert​ ​between​ ​moles​ ​and​ ​grams​ ​by​ ​using​ ​this​ ​triangle
or​ ​the​ ​equation:
moles​ ​=​ ​mass​ ​÷​ ​relative​ ​atomic​ ​mass
mass​ ​=​ ​moles​ ​x​ ​relative​ ​atomic​ ​mass
o E.g​ ​how​ ​many​ ​moles​ ​are​ ​there​ ​in​ ​42g​ ​of​ ​carbon?
▪ Moles​ ​=​ ​Mass​ ​/​ ​Mr​ ​=​ ​42/12​ ​=​ ​3.5​ ​moles
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1.29​ ​calculate​ ​reacting​ ​masses​ ​using​ ​experimental​ ​data​ ​and​ ​chemical
equations
● Chemical​ ​equations​ ​can​ ​be​ ​interpreted​ ​in​ ​terms​ ​of​ ​moles
o E.g.​ ​Mg​ ​+​ ​2HCl​ ​→MgCl​2​​ ​+​ ​H​2​​ ​shows​ ​that​ ​1​ ​mol.​ ​Mg​ ​reacts​ ​with​ ​2​ ​mol.​ ​HCl
to​ ​produce​ ​1​ ​mol.​ ​MgCl​2​​ ​and​ ​1​ ​mol.​ ​H​2
● Masses​ ​of​ ​reactants​ ​&​ ​products​ ​can​ ​be​ ​calculated​ ​from​ ​balanced​ ​symbol
equations.​ ​If​ ​you​ ​are​ ​given​ ​the​ ​reacting​ ​mass​ ​of​ ​one​ ​reactant​ ​and​ ​asked​ ​to​ ​find
the​ ​mass​ ​of​ ​one​ ​product​ ​formed:
o Find​ ​moles​ ​of​ ​that​ ​one​ ​substance:​ ​moles​ ​=​ ​mass​ ​/molar​ ​mass
o Use​ ​balancing​ ​numbers​ ​to​ ​find​ ​the​ ​moles​ ​of​ ​desired​ ​reactant​ ​or​ ​product
(e.g.​ ​if​ ​you​ ​had​ ​the​ ​equation:2NaOH​ ​+​ ​Mg​ ​→​ ​Mg(OH)​2​​ ​+​ ​2Na,​ ​if​ ​you​ ​had​ ​2
moles​ ​of​ ​Mg,​ ​you​ ​would​ ​form​ ​2x2=4​ ​moles​ ​of​ ​Na)
o Mass​ ​=​ ​moles​ ​x​ ​molar​ ​mass(of​ ​the​ ​product)​ ​to​ ​find​ ​mass
1.30​ ​calculate​ ​percentage​ ​yield
Percentage​ ​yield​ ​=​ ​ ​ ​ ​ ​ ​ ​ ​Amount​ ​of​ ​product​ ​produced​​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​x​ ​100
​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​Maximum​ ​amount​ ​of​ ​product​ ​possible
● It​ ​is​ ​not​ ​always​ ​possible​ ​to​ ​obtain​ ​the​ ​calculated​ ​amount​ ​of​ ​a​ ​product​ ​for​ ​3
reasons…
o Reaction​ ​may​ ​not​ ​go​ ​to​ ​completion​ ​because​ ​it​ ​is​ ​reversible
o Some​ ​of​ ​the​ ​product​ ​may​ ​be​ ​lost​ ​when​ ​it​ ​is​ ​separated​ ​from​ ​the​ ​reaction
mixture
o Some​ ​of​ ​the​ ​reactants​ ​may​ ​react​ ​in​ ​ways​ ​different​ ​to​ ​the​ ​expected
reaction
● Amount​ ​of​ ​product​ ​obtained​ ​is​ ​known​ ​as​ ​yield
1.31​ ​understand​ ​how​ ​the​ ​formulae​ ​of​ ​simple​ ​compounds​ ​can​ ​be​ ​obtained
experimentally,​ ​including​ ​metal​ ​oxides,​ ​water​ ​and​ ​salts​ ​containing​ ​water​ ​of
crystallisation
example​ ​experiment​ ​to​ ​find​ ​formula​ ​of​ ​magnesium​ ​oxide:
● weigh​ ​some​ ​pure​ ​magnesium
● Heat​ ​magnesium​ ​to​ ​burning​ ​in​ ​a​ ​crucible​ ​to​ ​form​ ​magnesium​ ​oxide,​ ​as​ ​the
magnesium​ ​will​ ​react​ ​with​ ​the​ ​oxygen​ ​in​ ​the​ ​air
● weigh​ ​the​ ​mass​ ​of​ ​the​ ​magnesium​ ​oxide
● Known​ ​quantities:​ ​ mass​ ​of​ ​magnesium​ ​used​ ​&​ ​mass​ ​of​ ​magnesium​ ​oxide
produced​ ​
● Required​ ​calculations:​ ​
○ mass​ ​oxygen​ ​=​ ​mass​ ​magnesium​ ​oxide​ ​-​ ​mass​ ​magnesium
○ moles​ ​magnesium​ ​=​ ​mass​ ​magnesium​ ​÷​ ​molar​ ​mass​ ​magnesium​ ​
○ moles​ ​oxygen​ ​=​ ​mass​ ​oxygen​ ​÷​ ​molar​ ​mass​ ​oxygen​ ​
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○ calculate​ ​ratio​ ​of​ ​moles​ ​of​ ​magnesium​ ​to​ ​moles​ ​of​ ​oxygen
○ use​ ​ratio​ t​ o​ ​form​ ​empirical​ ​formula
1.32​ ​know​ ​what​ ​is​ ​meant​ ​by​ ​the​ ​terms​ ​empirical​ ​formula​ ​and​ ​molecular
formula
● molecular​ ​formula-​ ​the​ ​number​ ​of​ ​atoms​ ​of​ ​each​ ​element​ ​in​ ​a​ ​compound
● empirical​ ​formula-​ ​the​ ​simplest​ ​whole​ ​number​ ​ratio​ ​of​ ​atoms​ ​of​ ​each​ ​element​ ​in
a​ ​compound
1.33​ ​calculate​ ​empirical​ ​and​ ​molecular​ ​formulae​ ​from​ ​experimental​ ​data
● Empirical​ ​formula​ ​from​ ​the​ ​formula​ ​of​ ​molecule:
● if​ ​you​ ​have​ ​a​ ​common​ ​multiple​ ​e.g.​ ​Fe​2​O​4​,​ ​the​ ​empirical​ ​formula​ ​is​ ​the
simplest​ ​whole​ ​number​ ​ratio,​ ​which​ ​would​ ​be​ ​FeO​2
● if​ ​there​ ​is​ ​no​ ​common​ ​multiple,​ ​you​ ​already​ ​have​ ​the​ ​empirical​ ​formula
● Molecular​ ​formula​ ​from​ ​empirical​ ​formula​ ​and​ ​relative​ ​molecular​ ​mass
● Find​ ​relative​ ​molecular​ ​mass​ ​of​ ​the​ ​empirical​ ​formula
● Divide​ ​relative​ ​molecular​ ​mass​ ​of​ ​compound​ ​by​ ​that​ ​of​ ​the​ ​empirical
formula
● Multiply​ ​the​ ​number​ ​of​ ​each​ ​type​ ​of​ ​atom​ ​in​ ​the​ ​empirical​ ​formula​ ​by​ ​this
number
● e.g.​ ​if​ ​answer​ ​was​ ​2​ ​and​ ​the​ ​empirical​ ​formula​ ​was​ ​Fe​2​O​3​​ ​then​ ​the
molecular​ ​formula​ ​would​ ​be​ ​empirical​ ​formula​ ​x​ ​2​ ​=​ ​Fe​4​O​6
1.34​ ​(chemistry​ ​only)​ ​understand​ ​how​ ​to​ ​carry​ ​out​ ​calculations​ ​involving
amount​ ​of​ ​substance,​ ​volume​ ​and​ ​concentration​ ​(in​ ​mol/dm3​ ​)​ ​of​ ​solution
● Concentration​ ​of​ ​a​ ​solution​ ​can​ ​be​ ​measured​ ​in​ ​mass​ ​per​ ​given​ ​volume​ ​of
solution​ ​e.g.​ ​grams​ ​per​ ​dm​3​​ ​(g/dm​3​)
● to​ ​calculate​ ​concentration​ ​of​ ​a​ ​solution​ ​use​ ​the​ ​equation
concentration​ ​(g​ ​dm​-3​)​ ​=​ ​mass​ ​of​ ​solute​ ​(g)​ ​¨​ ​volume​ ​(dm​3​)
● To​ ​calculate​ ​mass​ ​of​ ​solute​ ​in​ ​a​ ​given​ ​volume​ ​of​ ​a​ ​known
concentration​ ​use​ ​the​ ​equation:​ ​ ​mass​ ​=​ ​conc​ ​x​ ​vol​​ ​i.e.​ ​g​ ​=
g/dm​3​​ ​x​ ​dm​3​​ ​ ​(think​ ​about​ ​the​ ​units!)
1.35​ ​(chemistry​ ​only)​ ​understand​ ​how​ ​to​ ​carry​ ​out​ ​calculations​ ​involving​ ​gas
volumes​ ​and​ ​the​ ​molar​ ​volume​ ​of​ ​a​ ​gas​ ​(24dm​3​​ ​and​ ​24000​ ​cm3​ ​​ ​at​ ​room
temperature​ ​and​ ​pressure​ ​(rtp))
● Equal​ ​amounts​ ​in​ ​mol.​ ​of​ ​gases​ ​occupy​ ​the​ ​same​ ​volume​ ​under​ ​the​ ​same
conditions​ ​of​ ​temperature​ ​and​ ​pressure​ ​(e.g.​ ​RTP)
● Volume​ ​of​ ​1​ ​mol.​ ​of​ ​any​ ​gas​ ​at​ ​RTP​ ​(room​ ​temperature​ ​and​ ​pressure:​ ​20​ ​degrees
C​ ​and​ ​1​ ​atmosphere​ ​pressure)​ ​is​ ​24​ ​dm​3
● This​ ​sets​ ​up​ ​the​ ​equation:
Volume​ ​(dm​3​)​ ​of​ ​gas​ ​at​ ​RTP​ ​=​ ​Mol.​ ​x​ ​24
● Use​ ​this​ ​equation​ ​to​ ​calculate​ ​the​ ​volumes​ ​of​ ​gaseous​ ​reactants​ ​and​ ​products​ ​at
RTP
o e.g.​ ​5​ ​moles​ ​of​ ​H​2​​ ​would​ ​occupy​ ​a​ ​volume​ ​of​ ​24​ ​x​ ​5​ ​=​ ​120​ ​dm​3​​ ​at​ ​RTP
1.36​ ​practical:​ ​know​ ​how​ ​to​ ​determine​ ​the​ ​formula​ ​of​ ​a​ ​metal​ ​oxide​ ​by
combustion​ ​(e.g.​ ​magnesium​ ​oxide)​ ​or​ ​by​ ​reduction​ ​(e.g.​ ​copper(II)​ ​oxide)
● see​ ​1.31
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Edexcel​ ​IGCSE​ ​Chemistry
Topic​ ​1:​ ​Principles​ ​of​ ​chemistry
Ionic​ ​bonding
Notes
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1.37​ ​understand​ ​how​ ​ions​ ​are​ ​formed​ ​by​ ​electron​ ​loss​ ​or​ ​gain
● Ions​​ ​–​ ​Atoms​ ​that​ ​have​ ​lost​ ​or​ ​gained​ ​electron/electrons.
● Metal​ r​ eacting​ ​with​ ​a​ ​nonmetal:​ ​electrons​ ​in​ ​the​ ​outer​ ​shell​ ​of​ ​the​ ​metal​ ​atom
are​ ​transferred
o Metal​ ​atoms​ ​lose​ ​electrons​ ​to​ ​become​ ​positively​ ​charged​ ​ions
o Nonmetal​ ​atoms​ ​gain​ ​electrons​ ​to​ ​become​ ​negatively​ ​charged​ ​ions
● Cation​ ​=​ ​positive​ ​ion​ ​(+​ ​→​ ​ca+ion)
● Anion​ ​=​ ​negative​ ​ion​ ​(Negative​ ​→​ ​aNion)
1.38​ ​know​ ​the​ ​charges​ ​of​ ​these​ ​ions:​ ​metals​ ​in​ ​Groups​ ​1,​ ​2​ ​and​ ​3,​ ​nonmetals
in​ ​Groups​ ​5,​ ​6​ ​and​ ​7,​ ​Ag+​ ​,​ ​Cu2+
​ ​,​ ​Fe2+
​ ​,​ ​Fe3+
​ ​,​ ​Pb2+
​ ​,​ ​Zn2+
​ ​,​ ​hydrogen​ ​(H+​ ​),​ ​hydroxide
​
-​
2-​
(OH-​ ​),​ ​ammonium​ ​(NH​4+​ ​),​ ​carbonate​ ​(CO​32-)
​ ,​ ​nitrate​ ​(NO​3​ ),​ ​sulfate​ ​(SO​4​ )
●
●
●
●
●
●
●
group​ ​1​ ​→​ ​+1
group​ ​2​ ​→​ ​+2
group​ ​3​ ​→​ ​+3
group​ ​5​ ​→​ ​-3
group​ ​6​ ​→​ ​-2
group​ ​7​ ​→​ ​-1
the​ ​rest​ ​above​ ​just​ ​need​ ​to​ ​be​ ​learnt
1.39​ ​write​ ​formulae​ ​for​ ​compounds​ ​formed​ ​between​ ​the​ ​ions​ ​listed​ ​above
● compounds​ ​have​ ​no​ ​overall​ ​charge,​ ​therefore​ ​charges​ ​of​ ​ions​ ​must​ ​cancel​ ​out
1.40​ ​draw​ ​dot-and-cross​ ​diagrams​ ​to​ ​show​ ​the​ ​formation​ ​of​ ​ionic
compounds​ ​by​ ​electron​ ​transfer,​ ​limited​ ​to​ ​combinations​ ​of​ ​elements​ ​from
Groups​ ​1,​ ​2,​ ​3​ ​and​ ​5,​ ​6,​ ​7​ ​only​ ​outer​ ​electrons​ ​need​ ​to​ ​be​ ​shown
● ionic​ ​compounds​ ​are​ ​formed​ ​when​ ​a​ ​metal​ ​and​ ​nonmetal
react.
● Ionic​ ​bonds​ ​are​ ​formed​ ​by​ ​the​ ​transfer​ ​of​ ​electrons​ ​from
the​ ​outer​ ​shell​ ​of​ ​the​ ​metal​ ​to​ ​the​ ​outer​ ​shell​ ​of​ ​the
nonmetal.
● The​ ​metal​ ​therefore​ ​forms​ ​a​ ​positive​ ​ion​ ​and​ ​the
nonmetal​ ​forms​ ​a​ ​negative​ ​ion
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1.41​ ​understand​ ​ionic​ ​bonding​ ​in​ ​terms​ ​of​ ​electrostatic​ ​attractions
● A​ ​giant​ ​structure​ ​of​ ​ions​ ​=​ ​ionic​ ​compound
● Held​ ​together​ ​by​ ​strong​ ​electrostatic​ ​forces​ ​of​ ​attraction​ ​between​ ​oppositely
charged​ ​ions
● The​ ​forces​ ​act​ ​in​ ​all​ ​directions​ ​in​ ​the​ ​lattice,​ ​and​ ​this​ ​is​ ​called​ ​ionic​ ​bonding.
An​ ​example​ ​is​ ​sodium​ ​chloride​ ​(salt):
Na​+​​ ​(small​ ​blue​ ​particles)​ ​and​ ​Cl​-​​ ​(larger​ ​green​ ​ones)
1.42​ ​understand​ ​why​ ​compounds​ ​with​ ​giant​ ​ionic​ ​lattices​ ​have​ ​high​ ​melting
and​ ​boiling​ ​points
● Strong​ ​electrostatic​ ​forces​ ​of​ ​attraction​ ​between​ ​oppositely​ ​charged​ ​ions
● Requires​ ​a​ ​lot​ ​of​ ​energy​ ​to​ ​overcome​ ​these​ ​forces​ ​of​ ​attraction
● Therefore,​ ​the​ ​compounds​ ​have​ ​high​ ​melting​ ​and​ ​boiling​ ​points
1.43​ ​know​ ​that​ ​ionic​ ​compounds​ ​do​ ​not​ ​conduct​ ​electricity​ ​when​ ​solid,​ ​but
do​ ​conduct​ ​electricity​ ​when​ ​molten​ ​and​ ​in​ ​aqueous​ ​solution
● As​ ​a​ ​solid,​ ​the​ ​ions​ ​are​ ​in​ ​fixed​ ​positions​ ​so​ ​can’t​ ​conduct​ ​electricity
● when​ ​molten​ ​or​ ​in​ ​aqueous​ ​solution​ ​the​ ​ions​ ​are​ ​free​ ​to​ ​move​ ​carrying​ ​charge
and​ ​conducting​ ​electricity
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Edexcel​ ​IGCSE​ ​Chemistry
Topic​ ​1:​ ​Principles​ ​of​ ​chemistry
Covalent​ ​bonding
Notes
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1.44​ ​know​ ​that​ ​a​ ​covalent​ ​bond​ ​is​ ​formed​ ​between​ ​atoms​ ​by​ ​the​ ​sharing​ ​of
a​ ​pair​ ​of​ ​electrons
● Covalent​ ​bonding​ ​occurs​ ​in​ ​most​ ​non-metallic​ ​elements​ ​and​ ​in​ ​compounds​ ​of
nonmetals
● When​ ​atoms​ ​share​ ​pairs​ ​of​ ​electrons,​ ​they​ ​form​ ​covalent​ ​bonds.​ ​These​ ​bonds
between​ ​atoms​ ​are​ ​strong.
1.45​ ​understand​ ​covalent​ ​bonds​ ​in​ ​terms​ ​of​ ​electrostatic​ ​attractions
● Strong​ ​bonds​ ​between​ ​atoms​ ​that​ ​are​ ​covalently​ ​bonded​ ​are​ ​the​ ​result​ ​of
electrostatic​ ​attraction​ ​between​ ​the​ ​positive​ ​nuclei​ ​of​ ​the​ ​atoms​ ​and​ ​the​ ​pairs​ ​of
negative​ ​electrons​ ​that​ ​are​ ​shared​ ​between​ ​them
1.46​ ​understand​ ​how​ ​to​ ​use​ ​dot-and-cross​ ​diagrams​ ​to​ ​represent​ ​covalent
bonds​ ​in:​ ​diatomic​ ​molecules,​ ​including​ ​hydrogen,​ ​oxygen,​ ​nitrogen,
halogens​ ​and​ ​hydrogen​ ​halides,​ ​inorganic​ ​molecules​ ​including​ ​water,
ammonia​ ​and​ ​carbon​ ​dioxide,​ ​organic​ ​molecules​ ​containing​ ​up​ ​to​ ​two
carbon​ ​atoms,​ ​including​ ​methane,​ ​ethane,​ ​ethene​ ​and​ ​those​ ​containing
halogen​ ​atoms
some​ ​from​ ​the​ ​above​ ​list:
hydrogen​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​hydrogen​ ​chloride
​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​water​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​methane
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​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​oxygen​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​carbon​ ​dioxide
1.47​ ​explain​ ​why​ ​substances​ ​with​ ​a​ ​simple​ ​molecular​ ​structures​ ​are​ ​gases​ ​or
liquids,​ ​or​ ​solids​ ​with​ ​low​ ​melting​ ​and​ ​boiling​ ​points;​ ​the​ ​term
intermolecular​ ​forces​ ​of​ ​attraction​ ​can​ ​be​ ​used​ ​to​ ​represent​ ​all​ ​forces
between​ ​molecules
● Substances​ ​that​ ​consist​ ​of​ ​small​ ​molecules​ ​are​ ​usually​ ​gases​ ​or​ ​liquids​ ​that​ ​have
low​ ​boiling​ ​and​ ​melting​ ​points.
● Substances​ ​that​ ​consist​ ​of​ ​small​ ​molecules​ ​have​ ​weak​ ​intermolecular​ ​forces
between​ ​the​ ​molecules.​ ​These​ ​are​ ​broken​ ​in​ ​boiling​ ​or​ ​melting,​ ​not​ ​the​ ​covalent
bonds.
● Substances​ ​that​ ​consist​ ​of​ ​small​ ​molecules​ ​don’t​ ​conduct​ ​electricity,​ ​because
small​ ​molecules​ ​do​ ​not​ ​have​ ​an​ ​overall​ ​electric​ ​charge.
1.48​ ​explain​ ​why​ ​the​ ​melting​ ​and​ ​boiling​ ​points​ ​of​ ​substances​ ​with​ ​simple
molecular​ ​structures​ ​increase,​ ​in​ ​general,​ ​with​ ​increasing​ ​relative​ ​molecular
mass
● The​ ​intermolecular​ ​forces​ ​increase​ ​with​ ​the​ ​size​ ​of​ ​the​ ​molecules,​ ​so​ ​larger
molecules​ ​(i.e.​ ​molecules​ ​with​ ​greater​ ​relative​ ​molecular​ ​masses)​ ​have​ ​higher
melting​ ​and​ ​boiling​ ​points.
1.49​ ​explain​ ​why​ ​substances​ ​with​ ​giant​ ​covalent​ ​structures​ ​are​ ​solids​ ​with
high​ ​melting​ ​and​ ​boiling​ ​points
● Substances​ ​that​ ​consist​ ​of​ ​giant​ ​covalent​ ​structures​ ​are​ ​solids​ ​with​ ​very​ ​high
melting​ ​points.
● All​ ​of​ ​the​ ​atoms​ ​in​ ​these​ ​structures​ ​are​ ​linked​ ​to​ ​other​ ​atoms​ ​by​ ​strong​ ​covalent
bonds.
● These​ ​bonds​ ​must​ ​be​ ​overcome​ ​to​ ​melt​ ​or​ ​boil​ ​these​ ​substances.
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1.50​ ​explain​ ​how​ ​the​ ​structures​ ​of​ ​diamond,​ ​graphite​ ​and​ ​C60​ ​fullerene
influence​ ​their​ ​physical​ ​properties,​ ​including​ ​electrical​ ​conductivity​ ​and
hardness
Diamond
● In​ ​diamond​ ​(right),​ ​each​ ​carbon​ ​is​ ​joined​ ​to​ ​4​ ​other​ ​carbons
covalently.
o It’s​ ​very​ ​hard,​ ​has​ ​a​ ​very​ ​high​ ​melting​ ​point​ ​and​ ​does​ ​not
conduct​ ​electricity.
Graphite
● In​ ​graphite,​ ​each​ ​carbon​ ​is​ ​covalently​ ​bonded​ ​to​ ​3​ ​other​ ​carbons,​ ​forming​ ​layers
of​ ​hexagonal​ ​rings,​ ​which​ ​have​ ​no​ ​covalent​ ​bonds​ ​between​ ​the​ ​layers.
o The​ ​layers​ ​can​ ​slide​ ​over​ ​each​ ​other​ ​due​ ​to​ ​no​ ​covalent​ ​bonds​ ​between
the​ ​layers,​ ​but​ ​weak​ ​intermolecular​ ​forces.​ ​Meaning​ ​that​ ​graphite​ ​is​ ​soft
and​ ​slippery.
● One​ ​electron​ ​from​ ​each​ ​carbon​ ​atom​ ​is​ ​delocalised.
o This​ ​makes​ ​graphite​ ​similar​ ​to​ ​metals,​ ​because​ ​of​ ​its​ ​delocalised
electrons.
o It​ ​can​ ​conduct​ ​electricity​ ​–​ ​unlike​ ​Diamond.
Graphene
● Single​ ​layer​ ​of​ ​graphite
● Has​ ​properties​ ​that​ ​make​ ​it​ ​useful​ ​in​ ​electronics​ ​and​ ​composites
Carbon​ ​can​ ​also​ ​form​ ​fullerenes​ ​with​ ​different​ ​numbers​ ​of​ ​carbon​ ​atoms​.
● Molecules​ ​of​ ​carbon​ ​atoms​ ​with​ ​hollow​ ​shapes
● They​ ​are​ ​based​ ​on​ ​hexagonal​ ​rings​ ​of​ ​carbon​ ​atoms,​ ​but​ ​they​ ​may​ ​also​ ​contain
rings​ ​with​ ​five​ ​or​ ​seven​ ​carbon​ ​atoms
● The​ ​first​ ​fullerene​ ​to​ ​be​ ​discovered​ ​was​ ​Buckminsterfullerene​ ​(C60),​ ​which​ ​has​ ​a
spherical​ ​shape
Carbon​ ​nanotubes
● Cylindrical​ ​fullerenes​ ​with​ ​very​ ​high​ ​length​ ​to​ ​diameter​ ​ratios
● Their​ ​properties​ ​make​ ​them​ ​useful​ ​for​ ​nanotechnology,​ ​electronics​ ​and​ ​materials
1.51​ ​know​ ​that​ ​covalent​ ​compounds​ ​do​ ​not​ ​usually​ ​conduct​ ​electricity
● exceptions​ ​include:​ ​graphite​ ​and​ ​graphene
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Edexcel​ ​IGCSE​ ​Chemistry
Topic​ ​1:​ ​Principles​ ​of​ ​chemistry
Metallic​ ​bonding
Notes
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1.52​ ​(chemistry​ ​only)​ ​know​ ​how​ ​to​ ​represent​ ​a​ ​metallic​ ​lattice​ ​by​ ​a​ ​2-D
diagram
● Metals​ ​consist​ ​of​ ​giant​ ​structures​ ​of​ ​atoms​ ​arranged​ ​in​ ​a​ ​regular​ ​pattern.
● The​ ​electrons​ ​in​ ​the​ ​outer​ ​shell​ ​of​ ​metal​ ​atoms​ ​are​ ​delocalised​ ​and​ ​so​ ​are​ ​free​ ​to
move​ ​through​ ​the​ ​whole​ ​structure.
● The​ ​sharing​ ​of​ ​delocalised​ ​electrons​ ​gives​ ​rise​ ​to​ ​strong​ ​metallic​ ​bonds.
1.53​ ​(chemistry​ ​only)​ ​understand​ ​metallic​ ​bonding​ ​in​ ​terms​ ​of​ ​electrostatic
attractions
● Strong​ ​electrostatic​ ​attraction​ ​between​ ​negatively​ ​charged​ ​electrons​ ​and
positive​ ​metal​ ​ions
1.54​ ​(chemistry​ ​only)​ ​explain​ ​typical​ ​physical​ ​properties​ ​of​ ​metals,​ ​including
electrical​ ​conductivity​ ​and​ ​malleability
● Metals​ ​have​ ​giant​ ​structures​ ​of​ ​atoms​ ​with​ ​strong​ ​metallic​ ​bonding.
o Therefore,​ ​most​ ​metals​ ​have​ ​high​ ​melting​ ​and​ ​boiling​ ​points.
o They​ ​can​ ​conduct​ ​heat​ ​and​ ​electricity​ ​because​ ​of​ ​the​ ​delocalised
electrons​ ​in​ ​their​ ​structures.
o The​ ​layers​ ​of​ ​atoms​ ​in​ ​metals​ ​are​ ​able​ ​to​ ​slide​ ​over​ ​each​ ​other,​ ​so​ ​metals
can​ ​be​ ​bent​ ​and​ ​shaped.
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Edexcel​ ​IGCSE​ ​Chemistry
Topic​ ​1:​ ​Principles​ ​of​ ​chemistry
Electrolysis
Notes
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1.55​ ​(chemistry​ ​only)​ ​understand​ ​why​ ​covalent​ ​compounds​ ​do​ ​not​ ​conduct
electricity
● They​ ​do​ ​not​ ​have​ ​free​ ​electrons​ ​–​ ​the​ ​electrons​ ​are​ ​shared​ ​in​ ​a​ ​covalent​ ​bond
1.56​ ​(chemistry​ ​only)​ ​understand​ ​why​ ​ionic​ ​compounds​ ​conduct​ ​electricity
only​ ​when​ ​molten​ ​or​ ​in​ ​aqueous​ ​solution
● Ions​ ​are​ ​fixed​ ​when​ ​ionic​ ​compounds​ ​are​ ​solid,​ ​meaning​ ​they​ ​can’t​ ​move​ ​so​ ​can’t
conduct​ ​electricity
● when​ ​the​ ​compounds​ ​are​ ​molten​ ​or​ ​in​ ​aqueous​ ​solution,​ ​the​ ​ions​ ​(that​ ​are
electrically​ ​charged)​ ​are​ ​able​ ​to​ ​move​ ​and​ ​carry​ ​charge
1.57​ ​(chemistry​ ​only)​ ​know​ ​that​ ​anion​ ​and​ ​cation​ ​are​ ​terms​ ​used​ ​to​ ​refer​ ​to
negative​ ​and​ ​positive​ ​ions​ ​respectively
● aNion​ ​=​ ​Negatively​ ​charged​ ​ion​ ​(-)
● ca+ion​ =​ ​ ​posi+ively​ ​charged​ ​ion​ ​(+)
1.58​ ​(chemistry​ ​only)​ ​describe​ ​experiments​ ​to​ ​investigate​ ​electrolysis,​ ​using
inert​ ​electrodes,​ ​of​ ​molten​ ​compounds​ ​(including​ ​lead(II)​ ​bromide)​ ​and
aqueous​ ​solutions​ ​(including​ ​sodium​ ​chloride,​ ​dilute​ ​sulfuric​ ​acid​ ​and​ ​copper
(II)​ ​sulfate)​ ​and​ ​to​ ​predict​ ​the​ ​products
● During​ ​electrolysis,​ ​positively​ ​charged​​ ​ions​ ​move​ ​to​ ​the​ ​negative​ ​electrode
(cathode)​,​ ​and​ ​negatively​ ​charged​ ​ions​ ​move​ ​to​ ​the​ ​positive​ ​electrode​ ​(anode).
● Ions​ ​are​ ​discharged​ ​at​ ​the​ ​electrodes​ ​producing​ ​elements,​ ​this​ ​process​ ​is​ ​called
electrolysis
● When​ ​you​ ​have​ ​a​ ​ionic​ ​solution​ ​(NOT​ ​a​ ​molten​ ​ionic​ ​compound),​ ​your​ ​solution
will​ ​contain:​ ​the​ ​ions​ ​that​ ​make​ ​up​ ​the​ ​ionic​ ​compound,​ ​and​ ​the​ ​ions​ ​in​ ​water
(OH​-​​ ​and​ ​H​+​)
● at​ ​the​ ​cathode​ ​(-):
○ hydrogen​ ​(from​ ​H+​​ ​ ​in​ ​water)​ ​is​ ​produced​ ​UNLESS​ ​the​ ​+​ ​ions​ ​in​ ​the​ ​ionic
compound​ ​are​ ​from​ ​a​ ​metal​ ​less​ ​reactive​ ​than​ ​hydrogen
○ if​ ​the​ ​metal​ ​is​ ​less​ ​reactive,​ ​it​ ​will​ ​be​ ​produced​ ​instead
● at​ ​the​ ​anode​ ​(+):
○ oxygen​ ​(from​ ​OH​-​​ ​in​ ​water)​ ​will​ ​be​ ​produced​ ​UNLESS​ ​the​ ​ionic​ ​compound
contains​ ​halide​ ​ions​ ​(Cl​-​,​ ​Br​-​,​ ​I​-​)
○ if​ ​there​ ​are​ ​halide​ ​ions,​ ​the​ ​halogen​ ​will​ ​be​ ​produced​ ​instead​ ​(e.g.​ ​Cl​2​)
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● Using​ ​the​ ​logic​ ​above…
● Electrolysis​ ​of:
○ Sodium​ ​chloride​ ​solution
■ H​+​​ ​ions​ ​go​ ​to​ ​cathode​,​ ​H2​​ ​ ​(g)​ ​is​ ​produced​ ​(Na​ ​is​ ​more​ ​reactive​ ​than
hydrogen)
■ Cl​-​​ ​ions​ ​go​ ​to​ ​anode​,​ ​Cl​2​​ ​(g)​ ​is​ ​produced​ ​(Cl​-​​ ​are​ ​halide​ ​ions)
○ Copper​ ​(II)​ ​sulfate​ ​solution
■ Cu​+​​ ​ions​ ​go​ ​to​ ​cathode​,​ ​Cu​ ​(s)​ ​is​ ​produced​ ​(Cu​ ​is​ ​less​ ​reactive​ ​than
hydrogen)
■ OH​-​​ ​ions​ ​go​ ​to​ ​anode​,​ ​O2​​ ​ ​(g)​ ​is​ ​produced​ ​(SO​4​2-​​ ​ions​ ​are​ ​not​ ​halide
ions)
○ Water​ ​acidified​ ​with​ ​sulfuric​ ​acid
■ H​+​​ ​to​ ​cathode​,​ ​H2​​ ​ ​(g)​ ​is​ ​produced​ ​(these​ ​are​ ​the​ ​other​ ​ions​ ​present
in​ ​sulfuric​ ​acid​ ​H​2​SO​4​)
■ OH​-​​ ​to​ ​anode​,​ ​O2​​ ​ ​(g)​ ​is​ ​produced​ ​(SO​4​2-​​ ​ions​ ​are​ ​not​ ​halide​ ​ions)
○ Molten​ ​lead​ ​(II)​ ​bromide​ ​(demonstration)
■ Pb​2+​​ ​to​ ​cathode​,​ ​Pb​ ​(s)​ ​is​ ​produced​ ​(not​ ​in​ ​solution​ ​so​ ​these​ ​are
the​ ​only​ ​+​ ​ions​ ​present)
■ Br​-​​ ​to​ ​anode​,​ ​Br​2​​ ​(l)​ ​is​ ​produced​ ​ ​(not​ ​in​ ​solution​ ​so​ ​these​ ​are​ ​the
only​ ​-​ ​ions​ ​present)
1.59​ ​(chemistry​ ​only)​ ​write​ ​ionic​ ​half-equations​ ​representing​ ​the​ ​reactions
at​ ​the​ ​electrodes​ ​during​ ​electrolysis​ ​and​ ​understand​ ​why​ ​these​ ​reactions​ ​are
classified​ ​as​ ​oxidation​ ​or​ ​reduction
● This​ ​is​ ​an​ ​example​ ​of​ ​a​ ​half​ ​equation;​ ​the​ ​small​ ​number​ ​is
always​ ​the​ ​same​ ​as​ ​the​ ​2​ ​larger​ ​numbers​ ​within​ ​the
equation.​ ​&​ ​electrons​ ​are​ ​represented​ ​by​ ​the​ ​symbol​ ​‘e-‘
● Oxidation​ ​Is​ ​Loss​ ​(of​ ​electrons)
● Reduction​ I​ s​ ​Gain​ ​(of​ ​electrons)
● writing​ ​half​ ​equations​ ​for​ ​the​ ​reactions​ ​at​ ​each​ ​electrode:
○ negative​ ​electrode:​ ​X​+​​ ​->​ ​X,​ ​so​ ​ionic​ ​equation​ ​must​ ​be:
X​+​​ ​+​ ​e-​​ ​ ​->​ ​X,​ ​electrons​ ​gained,​ ​so​ ​positive​ ​ions​ ​are​ ​reduced
○ positive​ ​electrode:​ ​X​-​​ ​->​ ​X,​ ​so​ ​ionic​ ​equation​ ​must​ ​be:
X​-​​ ​->​ ​e-​​ ​ ​+​ ​X,​ ​electrons​ ​are​ ​lost,​ ​so​ ​negative​ ​ions​ ​are​ ​oxidised
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1.60​ ​(chemistry​ ​only)​ ​practical:​ ​investigate​ ​the​ ​electrolysis​ ​of​ ​aqueous
solutions
example-​ ​copper​ ​sulfate​ ​solution​ ​using​ ​copper​ ​electrodes
● set​ ​up:
○ anode​ ​is​ ​made​ ​of​ ​impure​ ​copper​ ​(that​ ​you​ ​are​ ​purifying)
○ cathode​ ​is​ ​made​ ​of​ ​pure​ ​copper
○ the​ ​solution​ ​is​ ​copper​ ​sulfate
● what​ ​happens:
○ Cu​2+​​ ​ions​ ​from​ ​the​ ​anode​ ​move​ ​to​ ​the​ ​cathode,​ ​where​ ​they​ ​gain​ ​electrons
and​ ​are​ ​discharged​ ​as​ ​pure​ ​copper
○ impurities​ ​form​ ​as​ ​sludge​ ​below​ ​the​ ​anode
● the​ ​cathode​ ​will​ ​increase​ ​in​ ​mass​ ​as​ ​it​ ​gains​ ​pure​ ​copper,​ ​whilst​ ​the​ ​anode​ ​will
lose​ ​mass​ ​as​ ​copper​ ​ions​ ​are​ ​lost​ ​(they​ ​replace​ ​the​ ​ones​ ​from​ ​the​ ​CuSO​4​​ ​solution
that​ ​go​ ​to​ ​the​ ​cathode)​ ​and​ ​so​ ​are​ ​impurities
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