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Wednesday, June 26, 2013

STPM CHEMISTRY NOTES AND QUESTIONS

STPM CHEMISTRY  NOTES AND QUESTIONS

TERM 1

Q1

TERM 2


Group 14
1.      Element: C  Si  Ge  Sn  Pb
2.      p-block elements
3.      with valence shell electronic configuration of ns2 np2
4.      unlike group 2 elements, different in properties (group 2 = all metal alkaline earth).
5.      Atomic radius ↑ due to ↑ proton number ↑ shell drastically
6.      IE (min energy to remove most loosely e- from atom gaseous to form uni ion) ↓ because of ↑ proton number.
7.      Down group, Zeff remain almost constant, but each element has one extra than previous.
8.      Zeff constant, ↑ extra shell, causes attraction between nucleus & electron cloud progressively weaker ↓ . Atomic radius ↑ (bigger) . 1st IE generally ↓ .
9.      Pb > Sn (IE, slightly higher) because the screening effect does not ↑ as much as expected as ↑ proton number.
10.  Down group Zeff remain almost constant but each successive element has one shell extra than prceding one.
11.  This causes the attraction between the nucleus and the electron cloud to get progressively weaker.
12.  ↑ atomic radius, ↓ attraction (nucleus – e- )
13.  ↑ atomic radius, 1st IE ↓ .
14.  IE, lead, Pb slightly higher than tin, Sn although Pb > Sn atom. Because screening effect does not ↑ as much as expected from tin, Sn to lead, Pb. ↑ of 32 protons & 32 elements.
15.  Screening effect 4f electron is less than expected.
16.  ↑ in nuclear charge is not totally being balanced by increase additional 32 electron.
17.  Zeff ↑ slightly from Sn to Pb to an ↑ in 1st IE.
18.  Electrical conductivity. Down group, properties change from non-metallic to metallic => ↑ electrical conductivity
19.  C (diamond) non-conductor because no delocalised e- in the giant covalent structure.
20.  Si & Ge = metalloids = semiconductors.
21.  Sn & Pb = metallic, good conductor of electricity, because have delocalised e- in their giant metallic structure.
22.  Melting point. C (diamond) has a giant covalent structure with strong covalent bonds hold individual atom together in 3D array. A lot of E need to break the covalent bond => very high melting point.
23.  C  >  Si  > Ge (↑ larger in size -> covalent weaker ↓) (Si & Ge = giant structure)
24.  Si & Ge have structure like diamond.
25.  Very large size ↑, covalent bond weaker ↓, melting point weaker ↓ .
26.  Ge larger than Si = melting point Ge lower than Si.
27.  Sn & Pb have giant  metallic structures.
28.  Large size, metallic bonds are relatively weak -> lower melting poin, compared C, Si, Ge.
29.  Solid Pb has a close packed structure.
30.  Sn is more open structure.
31.  Melting point : Pb > Sn
32.  Atomic radius : Pb > Sn
33.  Tetrachloride & Oxides of group 14 elements.
34.  All group 14 form tetrachlorides, MCl4.
35.  Colourless liquids at room temperature.
36.  MCl4 undergoes sp3 hybridisation to give bond angle 109.5o .
37.  Can be prepared by direct combination, by heating the element in a stream of dry chlorine gas.
38.  Boiling point of tetrachlorides.
39.  CCl4  SiCl4  GeCl4  SnCl4  PbCl4
40.  Down group: Size ↑, no. e- ↑, van der Waals ↑ stronger, boiling point ↑.
41.  CCl4 : exceptional behavior of CCl4, no explanation. M.p. -23 / b.p. 77oC .
42.  PbCl4 : no normal b.p. for PbCl4 because it decomposes on heating.
43.  Down group : MCl4 (other than CCl4) gradually ↑.
44.  CCl4 has smallest size, m.p & b.p. higher than expected. Abnormal behavior. No suitable explanation.
45.  van der Waals CCl4 > stronger than SiCl4 molecules.
46.  Thermal stability of tetrachlorides.
47.  Covalent bond : depends on the bond lenght, longer bond, weaker, easier to break bond.
48.  Down group (C, Si, Ge, Sn, Pb) M-Cl bond become longer & weaker.
49.  As atomic radius ↑.
50.  (CCl4, SiCl4, GeCl4 => stable at heat, very hight Temperature) SnCl4, (PbCl4, very unstable)
51.  Hydrolysis of the MCl4.
52.  Break of H2O.
53.  All group 14 (except CCl4) hydrolysis with water produce hydrochloric acid or hydrogen chloride gas (depend amount of H2O).
54.  C, (Si, Ge, Sn, Pb, respectively chlorides make use of their empty d-orbitals in their valence shells to form coordinate bonds with water molecules).
55.  Si atom use two empty 3d orbitals to accept lone pair electrons from 2 water molecules to form a hexa-valence intermediate.
56.  CCl4 has no d-orbitals of comparable energy. Not hydrolysed by water.
57.  Thermal stability of the oxides.
58.  Group 14 form 2 type oxides  [1. Monoxide, MO (+2), 2. Dioxide, MO2 (+4)].
59.  Table :



Monoxide
CO
SiO
GeO
SnO
PbO
Physical state
Gas
Solid
Structure
Simple molecule
Predominantly ionic
Thermal stability
Form dioxide on heating in air
Stable
Acid / base nature
Neutral
Amphoteric


60.  Table
Dioxide
CO2
SiO2
GeO2
SnO2
PbO2
Physical state
Gas
Solid
Structure
Simple molecule
Giant covalent
Intermediate between giant covalent & giant ionic
Thermal stability
Stable
Unstable
Acid / base nature
Acidic
Amphoteric
Stable to heat
/
/
/
/



61.  Acid-base nature of the oxides.
62.  Neutral oxides (CO, SiO)
63.  C, Si, Ge, Sn, Pb. Down group, metallic character ↑, with the ↑ proton number.
64.  Relative stability of +2 and +4 oxidation states of group 14.
65.  G.14 can exhibit with valence e- (+2 à s2 electron , +4 à s2 & p2 electron).
66.  +2 oxidation state: M (g) à Mg2+ (g) + 2e-  /kJ mol-1
C, Si : covalent, Ge : covalent with slight ionic characteristic, Sn, Pb : ionic
67.  +4 ions: M (g) à M4+ (g) + 4e-  / kJ mol-1 (all covalent, energy very high).
68.  Down group, atomic radius ↑ , strength covalent bonds ↓ (weaker), less energy released. Result :  stability: +2 ↑, +4 ↓
69.  Inert pair effect = tendency of the latter number of a group to exhibit a valency 2 units less than the group valency due to the reluctance of the s electrons to participate in the reaction.
70.  ↑ stablility of +2 oxidation state can be seen from the standard electrode potential :
More + Eo value: more stable.
71.  More +ve Eo value; less stable +4 oxidation state.
72.  More –ve Eo value, more stable +2 oxidation state.
73.  Left –ve, Eo value periodic table right +ve.
74.  Pb4+ strongly oxidising (powerful oxidising agent). Very high tendency to accept e- to be converted to Pb2+  .
75.  Silicon, silicone & silicates.
76.  Silicone use to make silicone (an organosilicon polymer).
77.  Si, semiconductor (electronic transistor, microchips, IC (integrated circuits), computer component).
78.  Silicones can be obtained in the form of : (oils, grease, vublaer –like solids).
79.  Silicones are chemically inert & good water repellents. Use as lubricants, hydroulic fluids, electric insulators, elastomers, resins, electrical condensers, car polishes, implant, in plastic surgery, water-proof fabric & use in moulds to prevent casting from sticking mould.
80.  Silicate, SiO4 4- (tetrahedral anion, orthosilicate ion) use in making glass. Example: asbestos, sand, quartz, granite, mica & talcum, clay, feldspar. Silicate compound containing Si & O2 .
81.  75% earth’s crust consist of silicates.
82.  Chain silicate:
a.       single (pyroxenes) : sharing 2 oxygen, long linear chain, example; Na2SiO3 (sodium silicate), Mg(SiO3)2 (magnesium silicate), CaMg(SiO3)2 (calcium magnesium silicate).
b.      Double (amphiboles) : 2 single chain silicate join, share 2/3 oxygen atom, emperical formula [Si4O11]6- ions, strong covalent bond (inter), (intra) weak forces hold individual chain, result, brittle, breaks easily, single & double chain. Example: asbestos (heat insulator), break-pads, dust-hazardous (tuberculosis & lung cancer).
83.  Sheet silicates.
84.  Framework & giant structure silicates.
85.  Zeolites : use as cation exchanges, drying agents, heteronegous catalysts, molecular sieves.
86.  Glass (as transparent amorphous solids called quartz glass). Hard, non-crystalline transparent substance, not true solid, use as container for liquids, transparent, chemically inert (except to concentrated HF and caustic alkalis), non-toxic, easily recyclable.
87.  Ceramics. Contain clay (silicates & aluminosilicates), metal oxide; hard, brittle, stable (even at high temperature). Problem = brittles, deform / suddenly shatter without warning. Uses: electrical insulator, glass-wares, pottery, water containers, bathroom tiles, as support for precious metal catalyst in catalytic converters, light weight, retain properties at > 1000oC, chemically inert (strong bond in structure). Problem: brittleness, weak point within the ceramic matrix.

88.  Tin Alloys. Use as plating iron/ steel container to form ‘tin can’, as long as tin layer not scratched, iron beneath will not rust. Making alloys: bronze (70:30 Cu:Sn), Pewter (95:3:1 Sn:Cu: Sb), solder (30:70 Sn: Pb).



Group 17
1.      Physical properties, known as halogen, F Cl  Br  I  At, all 7 e- in valence shell, ns2 np5 , simple diatomic molecules, X2, Cl2  , Br2  , I2  , electron sharing.
2.      All reactive non-metals. Most reactive group of element in periodic table:
a.       Low bond dissociation energy,
b.      High electronegativity.
3.      Down group ↑ colour intensity.
Atomic radius & ionic radius. Down group, Zeff constant, no. of electronic shell ↑ .
4.      Size: Ionic > atomic
5.      Volatility G.17. m. p. / b. p. Measure of the strength of the intermolecular forces hold the molecules together. Stronger force higher m.p & b.p. ↑
6.      X2 --- X2 van der Waals forces (intermolecular forces).
7.      Down group 17, size ↑ total number e- ↑ , cause ↑ intermolecular van der Waals force. Result m.p & b.p ↑
8.      Down group: less volatile; colour gets darker (↑ proton number).
9.      Solubility in water. Halogens sparingly soluble in water. Cannot form hydrogen bonds with water molecules. Cl2 (g) + H2O ó HCl (aq) + HOCl (aq)
10.  Iodine completely soluble in aq potassium iodide, KI (aq). Formation of water – soluble complete ion, I3 - , I2 (aq) + I- (aq) à I3 (aq)
11.  F2 reacts violently with water. Liberate oxygen gas, 2F2 + 2H2O à 2HF + O2
12.  2HOCl, chloric (I) acid (aq) à(strong sunlight, exposes) 2HCl (aq) + O2 (g)
13.  Reactions group 17 (selected only) Bond Energy = energy required to break covalent bond per mole. X-X (g) à 2X (g) DHo = Bond Energy
14.  B.E. ↓ down group (Cl2 , Br2, I2, : size atom ↑, covalent bonds longer & weaker.
15.  F2 ; lower B.E eventhough it has smaller atomic size. Because closeness of the atoms in the molecules, repulsion between non-bonding electrons, bond weaker.
16.  As oxidising agents. Halogen act as electron acceptor to achieve octet configurations.
X (g) +e- ó X- (aq)
17.  All powerful oxidising agents => very large +ve Eo value, down group lower +ve Eo value (F, Cl, Br, I). (oxidising strength ↓)
18.  Down group, size increases causing decrease in tendency to accept electrons (oxidising strength ↓).
19.  Cl2 (g) + 2Br- (aq) à 2Cl- (aq) + Br2 (aq)
Cl2 (g) + 2I- (aq) à 2Cl- (aq) + I2 (aq)
20.  Br2 (aq) + 2I- (aq) à 2Br- (aq) + I2 (aq)
21.  Who can oxidise : F, Cl can Br & I, Br only I
22.  Reaction with hydrogen. X2 (g) + H2 (g) à 2HX (g)
23.  Down group ; oxidising strenght halogens ↓ also reactivity towards H ↓ ( Cl, Br, I)
24.  Cl2 (g) + H2 (g) à 2HCl (g)
25.  Br2 (g) + H2 (g) heat à 2HBr (g)
26.  I2 (g) + H2 (g) ó 2HI (g)
27.  Hydrogen Halides. Boiling point. Simple molecule of HX, all colourless gases (fume in moist air)
HCl , HBr , HI (size ↑ à) total number electron in molecule : ↑ (molar mass / gram)
Intermolecular van der Waals forces : ↑ stronger ,  b.p. ↑ à
28.  Thermal Stability. Depends on strength of H-X bond.
29.  2HX (g) all hydrogen halide decompose on heating à H2 (g) + X2 (g)
HX down group:  get longer, weaker (thermal stability ↓) .
30.  Reaction of chlorine with NaOH, sodium hydroxide.
31.  Cl2 (g) + dilute 2 NaOH (aq) à NaCl (aq) + NaOCl, sodium chlorate (I) use as domestic bleach (aq) + H2O (l) ,  concentrated sodium hydroxide at 70oC à sodium chlorate (v) formed.
Oxidation state : 0 to +1, 0 to -1 ; disproportionation [ chlorine reduced to chloride (-1) & oxidised to chlorate (I) (+1) at the same time].
32.  3Cl2 (g) + 6 NaOH heat à 5NaCl (aq) + NaClO3 (aq) use in explosive & weed killer + 3H2O (l)
Oxidation state : disproportionation, chlorine reduced to chloride (-1) & oxidised to chlorate (v) (+5) at the same time.  
33.  Reaction of selected Halide Ions.
34.  Reaction of halide ions with silver nitrate, Ag+ (aq) silver ion + X- (aq) halide ions à AgX (s) coloured precipitate.
35.  AgCl (s) white, AgBr (s) cream, AgI (s) yellow.
36.  AgCl (s) silver chloride + 2NH3 (aq) ammonia, dilute/ concentrated à [Ag(NH3)2]+ (aq) diammine silver complex ion, water soluble colouless solution + Cl- (aq)
37.  AgBr (s) silver bromide + 2NH3 (aq) ammonia dilute insoluble/ concentrated soluble à [Ag(NH3)2]+ (aq) + Br- (aq)
38.  AgI (s) insoluble for both dilute & concentrated NH3 ammonia.
39.  Reaction halides with concentrated H2SO4 sulphuric acid.
40.  X- (s) solid halide/ halide ion + H2SO4 (aq) heat à HX (g) hydrogen halide + HSO4 (aq)
41.  Example : NaCl (s) sodium chloride + H2SO4 (aq) concentrated sulphuric acid heat à NaHSO4 (aq) + HCl (g) hydrogen chloride white fume
42.  KBr (s) potassium bromide + H2SO4 (aq) concentrated sulphuric acid heat à KHSO4 (aq) + HBr (g) hydrogen bromide white fume
When continue heating/ prolong heating: 2HBr (g) hydrogen bromide + H2SO4 (aq) concentrated sulphuric acid, act as oxidising agent, heat à Br2 (g) bromine, reddish brown fume + SO2 (g) + 2H2O(g)
43.  KI (s) potassium iodide + H2SO4 (aq) concentrated sulphuric acid, heat à KHSO4 (aq) + HI (g) hydrogen iodide, white fume.
Continue / prolong heating : 2HI (g) + H2SO4 (aq), heat à I2 (g) iodine violet fumes + SO2 (g) + 2H2O (g)
44.  Concentrated H2SO4 not suitable agent to prepare HBr & HI. H3PO4 suitable because = non-oxidising agent.
KBr (s) + H3PO4 (aq) concentrated phosphoric acid heat à HBr (g) + KH2PO4 (aq)
KI (s) + H3PO4 (aq) heat à HI (g) + KH2PO4 (aq)
45.  Industrial Aplications of  X & Their Compounds.
46.  Uses of chlorine. Chlorine; purify water (swimming pool/ piped water), bleaching agent (paper industry).
47.  Sodium chlorate (I); domestic bleach, antiseptic.
48.  Chlorinated organic compounds; solvent, CCl4 , CHCl3, chloroform.
49.  Trichloroethane; dry-cleaning agent.
50.  Freons [CCl2F2 & CFCl3] ; cleaning agents, refrigerants, propellant in aerosol cans.
51.  Dichlorodiphenyltrichloroethane, DDT; powerful insecticide.
52.  Postassium chlorate (v) ; weed-killer, explosives.
53.  Chloroethene, CH2 = CHCl ; monomer & PVC, plastic like (pipes, leather).
54.  Uses of Br. Bromine (dyes, drugs).
55.  1,2-dibromoethane, BrCH2CH2Br (petrol additive remove lead compounds (Pb) from engines.
56.  Silver bromide, manufacture of photographic films (AgBr is more sensitive to light than AgCl).
57.  Uses of iodine. Iodine (make dyes, colour photography).
58.  Iodoform (triiodomethane) ; antiseptic.
59.  AgI, silver iodide ; cloud seeding (provides nucleus condensation of water vapor in cloud).
60.  Iodine solution in ethanol (aq) ; use antiseptic.
61.  Black & white photography. AgCl (s) à Ag (s) metallic silver + 1/2Cl2 (g)
AgBr (s) à Ag (s) + 1/2Br2 (l)
AgBr more prefer, more sensitive to light.
AgI à Ag (s) + 1/2I2 (s)   
Silver halides sensitive to light. AgX when expose to light (turn dark).
AgX decompose to form metallic silver & halogen; photosensitive.
Black & white photographic film = clear cellulose strip coated with grains of AgBr, silver bromide.
AgBr (s) silver bromide, sunlight strike the film, activated à AgBr * (s)
Then exposed film treated with aqueous hydroquinone, C6H2O2 (reducing agent). Reduce activated AgBr à Ag
2AgBr* (s) + C6H6O2 (aq) hydroquinone à 2Ag (s) + 2Hbr (aq) + C6H4O2 (aq) quinone.
AgBr (s) unactivated silver bromide (removed from the film) + 2S2O3 2-  (aq) sodium thiosulphate à Ag(S2O3)2 3- (aq) + Br- (aq)
Film negative : struck by light appear black (due to deposition of metallic silver).

Light passes through negative onto a photosensitive paper, dark area would appear white. While clear area would appear black.




TERM 3


Notes/ Power Point Presentation slides:

CHAP 01 CHEMISTRY FORM 6 TERM 3 INTRO TO ORGANIC CHEMISTRY
CHAP 02 CHEMISTRY FORM 6 SEM 3 HIDROCARBONS
CHAP 03 CHEMISTRY FORM 6 SEM 3 BENZENE AND ITS COMPOUND
CHAP 04 CHEMISTRY FORM 6 SEM 3 HALOALKANE
CHAP 05 CHEMISTRY FORM 6 SEM 3 ALCOHOL
CHAP 06 CHEMISTRY FORM 6 SEM 3 CARBONYL COMPOUNDS

CHAPTER 14 : INTRODUCTION TO ORGANIC CHEMISTRY

* bonds = single, double, triple, straight chain, branched chain, ring.
* 3 types of chain = straight, branched & ring/ cyclic.
* 2 main types  of isomerism = structural & stereoisomerism.
* Hybridisation = is the mixing of two or more non-equivalent atomic orbitals to give a set of equivalent hybrid orbital.
* sp3 hybridisation = mixing an s-orbital and 3 p-orbital to form 4 equivalent sp3 hybrid orbitals.
* sp2 hybridisation = mixing of an s-orbital and 2 p-orbitals to form 3 sp2 hybrid orbitals.
* s p hybridisation = 2 unhibrid 2p orbitals.
*spis pronounced as s-p-three
*A π, pi bond is weaker than a σ sigma bond.
* Empirical formula = simplest whole number ratio between the different type of atom (or element) in a molecule of the compound.
* Molecular formula = of an organic compound shows the actual number of the different type of atoms (or element) in one molecule of the compound.
* Structural formula = shows the actual number of the different type of atom (or elements) in a molecule of the compound and how the atom are connected to one another.
* Straight Chain, example: _________
* Branched Chain, example: _________
* Ring / Cyclic structure, example: _________
* Double bond, example: _________
* Triple bond, example: _________
* methane, ______ (molecular formula)
* ethane, ______ (molecular formula)
* ethene, ______ (molecular formula)
* ethyne, ______ (molecular formula)
* benzene, structural formula = _________

* Draw 









s-orbital
p-orbital
d-orbital








* Draw Structural Formula & It's Skeletal :
1. ethanoic acid
2. ethane
3. 2-propanol
4. propanone
5. 2-methylpropene
6. popanal

No. of C
Alkane
Alkene
Alkyne

1
Meth
Draw:



Name:
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2
Eth
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3
Prop
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Name:
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Name:
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4
But
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5
Pent
Draw:



Name:
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Name:
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Name:

6
Hex
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Name:
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7
Hept
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8
Oct
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Name:
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9
Nona
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10
Dec
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*
No. of Carbon
shrt
Alkane
Alkene
Alkyne
1
Meth
Methane
Methene
Methyne
2
Eth
Ethane
Ethene
Ethyne
3
Prop
Propane
Propene
Propyne
4
But
Butane
Butene
Butyne
5
Pent
Pentane
Pentene
Pentyne
6
Hex
Hexane
Hexene
Hexyne
7
Hept
Heptane
Heptene
Heptyne
8
Oct
Octane
Octene
Octyne
9
Nano
Nonane
Nonene
Nonyne
10
Dec
Decane
Decene
Decyne














*

======================================================================== TERM 2

TERM 1

Chapter 1 : ATOMS, MOLECULES AND STOICHIOMETRY

* proton number (symbol Z) =____ * nucleon number (symbol A) =
* A (_______) = Z (______) + N (______) *A->Z->X = _____
* isotopes = _______
* In atomic physics and quantum chemistry, the electron configuration is......
* Half-life = ____ * Nuclear reaction. fission = _____, fusion = ______.
* Radioactive isotope = _____  * Relative Isotopic Mass = _____
* Relative Atomic Mass = ______  * Relative Molecular Mass = _____
* Relative Formula Mass = _______ * one mole = ______  * 1 mole = 6.02....?
* n =MV/???? 
* Calculate the mass of lead that contains:
a. 3.45 mol of lead = ______ g ,
b. 8.4 X 10^22 atoms = _____ g ,
c. 9.8 X 10^24 atoms = ______ g
* Avogadro's number / Avogadro's constant =____  * Solution = ______
* Solute = _____ * Solvent = _____ * Concentration = _____

* Molar Concentration / Molarity = ____ * Molarity (mol dm^-3) = _____

Chapter 1 : ATOMS, MOLECULES AND STOICHIOMETRY

* proton number (symbol Z) = is the total number of protons in the nucleus of an atom. This is also equal to the number of electrons in the atom.
* nucleon number (symbol A) = is the total number of protons and neutrons in the nucleus of an atom.
* A (nucleon) = Z (proton no.) + N (neutron)
*A->Z->X
* isotopes = are atoms having the same number of protons but different number of neutrons.
* In atomic physics and quantum chemistry, the electron configuration is the distribution of electrons of an atom or molecule (or other physical structure) in atomic or molecular orbitals.
* Half-life = of a radioisotope is the time taken for the amount of the isotope or its activity (disintegration per second) to decrease to half its original value.
* Nuclear reaction. fission = breaking, fusion = joining (needs very high temperature for it to occur)
* Radioactive isotope = Isotope with unstable nucleus and undergoes spontaneous disintegration to form nucleus of smaller isotopes.
* Relative Mass
* Relative Isotopic Mass = mass of one atom of the isotope relative to 1/12 times the mass of one atom of the C-12 isotope.
* Relative Atomic Mass = of an element is the average mass of one atom of the element relative to 1/12 times the mass of a C-12 atom.
* Relative Molecular Mass = of a molecular species (elements of compounds) is the mass of one molecule of the substance relative to 1/12 times the mass of one carbon-12 atom.
* Relative Formula Mass = of an ionic compound is the sum of the relative atomic mass of all the atoms in one formula unit of the compound. The mass of an ion is taken to be the same as the mass of its neutral atom.
* one mole = amount of substance which contains the same number of particles (atom, molecules, ions or electrons) as the number of atoms in 12 grams of carbon-12.
* 1 mole = 6.02 x 10^23
* n =MV/1000
* Calculate the mass of lead that contains: a. 3.45 mol of lead = 714.15 g , b. 8.4 X 10^22 atoms = 28.88 g , c. 9.8 X 10^24 atoms = 3369.8 g
* Avogadro's number / Avogadro's constant = number of atoms in exactly 12 g of carbon-12 is 6.02 X 10^23. The symbols are NA or L.
* Solution = a homogeneous mixture of two or more substances.
* Solute = substance present in smaller quantity.
* Solvent = substance present in large quantity. most common solvent is water.
* Concentration = amount of solute present in a fixed quantity of the solution.
* Molar Concentration / Molarity = concentration in mol dm^-3
* Molarity (mol dm^-3) = [concentration (g dm^-3)] / [molar mass of solute (g mol^-1)]
*











Homework
15 multiple-choice questions with answer
                                 +
4 structured/essay questions with answer
19 (format for final exam)






CHEMISTRY'S TITLE OF EXPERIMENTS AND PROJECT FOR STPM

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