ANSWER PAPER 1

ANSWER FOR PAPER 1 SPM 2013

1	D	26	C
2	A	27	D
3	D	28	D
4	C	29	C
5	B	30	B
6	C	31	A
7	B	32	B
8	C	33	A
9	A	34	D
10	** B/C	35	D
11	D	36	A
12	B	37	A
13	C	38	D
14	C	39	B
15	D	40	A
16	D	41	D
17	C	42	C
18	B	43	A
19	A	44	B
20	B	45	C
21	D	46	D
22	A	47	A
23	B	48	B
24	C	49	A
25	C	50	C

question 10

B  fats found in animal but oil found in vegetables and fishes ( animal right?）

C  Fats are mixtures of organic acid( fatty acid/ carboxylic acid is organic acid)  and glycerol.

so hopefully this question is a bonus for students.

Electrolysis Of Molten Compounds

Electrolysis Of Molten Compounds

Process of breaking down compounds in molten or aqueous state into their constituent elements by passing electricity through them.

v  Anode (Positive electrode)

─       Negative ions or anions are attracted to anode.

─       Anion undergo discharge by releasing electrons to form neutral atoms.

v  Cathode (Negative electrode)

─       Positive ions or cations are attracted to cathode.

─       Cation  undergo discharge to form atoms.

Electrolysis Of Molten Lead (II) Bromide

v  Anode (Positive electrode)

─       Anions, Br ^- are attracted to the anode.

─       Bromide ions, Br^- discharge by releasing electron to form bromine gas, Br₂

Half equation :

2Br^-(l)      →       Br₂(g) + 2e^-

v  Cathode (negative electrodes)

─       Lead(II) ions, Pb²⁺ are attracted to the cathode.

─       Lead (II) ions, Pb²⁺ undergo discharge by accepting two electrons to form a lead atom/lead metal.

Half equation :

Pb²⁺ (I) + 2e^-       →         Pb (s)

Overall equation:

Pb²⁺ (I) + 2Br ^-  → Pb(s) + Br₂(g)

Electrolytes & Non-Electrolytes

Electrolytes

v  Substance that can conduct electricity in molten state or aqueous solution.

v  Contains free moving ion that carry electrical charges.

v  Ionic compounds such as molten lead (II) chloride, PbCI₂.

v  Acid solutions such as nitric acid, HNO₃.

Non-Electrolytes

v  Substance which cannot conduct electricity in all states.

v  Made from molecules only-no ions that can carry electrical charges and conduct electricity.

v  Covalent compounds such as alcohol, sugar and naphthalene.

Elements In A Period

Elements In A Period

v  The Change In The Physical Properties of Elements Across Period 3

1)      Proton number increases by one from one element to another element.

2)      Number of valence electron increases by 1.

3)      Physical state change from solid to gas.

4)      Atomic radius (atomic size) of elements decreases because:

Ø  Force of attraction between nucleus and valence electron increases.

Ø  The stronger the force of attraction, the smaller the atomic size.

5)      Electronegativity of element increases because:

Ø  As the size of atom decreases, force of attraction between nucleus and valence electron increases.

Ø  Thus, the strength of the atom to attract electrons to the nucleus increases.

Ø  Metallic properties decrease across the period.

v  The Changes In Oxide Properties of Elements Across Period 3

The properties of the oxide of elements in Period 3 change from basic to acidic

Na2O MgO	Basic oxides	§  Metal oxides are basic. §  Metal oxides dissolve in water to form alkaline solution. v  Example: Na2O (s)  +  H2O (l) →  2NaOH (aq)
Al2O3	Amphoteric	§  Amphoteric means oxide that shows both properties of acid and base. v  Examples: Al2O3 (s)  +  6 HCl (aq) →  2AlCl3 (aq)      +      3H2O (l) Al2O3 (s)   +   2 NaoH (aq)  →  2NaAlO2 (aq)  +  H2O (l)
SiO2 P4O10 SO2 Cl2O7	Acidic	§  Non-metal oxides are acidic. §  Non-metal oxides dissolve in water to form acidic solutions. v  Example: SO2 (s)   +   H2O (l) → H2SO3 (aq)

Transition Elements

Transition Elements

Ø  Elements that is located in Group 3 to Group 12 in the Periodic Table.

Ø  Examples: scandium, iron, nickel, tungsten and copper

Physical  Properties:

§  Solid

§  Shiny surface

§  Ductile and malleable

§  Good conductor of heat and electricity.

§  High density.

§  High melting point and boiling point.

Special characteristics

v  Show different oxidation numbers in their compounds.

v  Form coloured ions or compounds.

v  Act as catalyst.

v  Form complex ions.

•          These characteristics are not possessed by other metals.

Ions	Colour of aqueous solution
Fe2+	Green
Fe3+	Yellow/ brownish yellow/ brown
Mn2+	Pink
MnO4-	Purple
Cr2O72-	Orange
CrO42-	Yellow

The Uses Of Transition Metals In Industries

Process	Manufacture	Catalyst
Contact Process	Sulphuric acid, H2SO4	Vanadium oxide, V2O5
Ostwald process	Nitric acid, HNO3	Platinum
Haber process	Ammonia, NH3	Iron
	Margarine	Nickel

GROUP 1

GROUP 1

1.       Alkali metals

2.       Exist naturally as compounds

3.       Elements: Lithium, sodium, potassium, rubidium, caesium and francium.

4.       General physical properties of lithium, sodium and potassium:

·         Exist as solid that have silvery and shiny surfaces at room temperature.

·         Low melting points and boiling point compared to other metals

·         Good conductors of heat and electricity

·         Low densities

5.       The Changes In the Physical Properties From Lithium, Sodium to Potassium

a)      Atomic size increases

Reason : Number of shells increases.

b)      Melting point and boiling point decreases

Reason : The increase in atomic size down the group causing the metal bond between atoms become weaker.

c)       Density increases

Reason : Increase in atomic mass more than the increase in atomic radius (volume of atom).

v  Show similar chemical properties (same number of valence electron in its outermost shell) but differ in reactivity.

v  Alkali metals donate one electron from its outermost occupied shell to form an ion with a charge of +1, achieving stable electron arrangement.

Example:

Na → Na⁺  + ē

      2.8.1      2.8

Chemical reaction for alkali metal

1.       Alkali metal + H₂O → Metal hydroxide solution + hydrogen gas

Example: 2Li(s) + 2H₂O(l) → 2 LiOH(aq) + H₂(g)

2.       Alkali metal + oxygen gas → Metal oxide

Example: 4Li(s) + O₂(g) → 2Li₂O(s)

3.       Metal oxide + water → Metal hydroxide solution

Example: Li₂O(s) + H₂O(l) → 2LiOH(aq)

4.       Alkali metal + chlorine gas

Example: 2Na(s) + Cl₂(g) → 2NaCl(s)

5.       Alkali metal + Bromine gas → Metal bromide

Example: 2K(s) + Br₂(g) → 2KBr(s)

** Halogens ( chlorine, bromine are always exist as diatomic molecules, thus the chemical formula are Cl2, Br2 )

GROUP 17

GROUP 17  

Fluorine, Chlorine, Bromine, Iodine, Astatine

1. Group 17 elements are known as halogens.

2. They are exist as  DIATOMIC  MOLECULES with molecular formula of F₂, Cl₂, Br₂, I₂

3. Electronegativity is a measurement of the strength of an atom in a molecule to pull electrons towards its nucleus.

When going down to the Group 17, the electronegativity of halogens decreases.

Explanation:

-          The number of shells occupied with electrons increases.

-          This cause the outermost occupied shells further away from the nucleus.

-          Thus, the strength of nucleus to attract electrons becomes weaker.

5. General physical of halogens:

-       Very low melting points and boiling point. (Weak forces of attraction between MOLECULES )

-       Low density

-       Cannot conduct electricity

When going down the Group 17, The changes in physical properties.
Atomic radius/ atomic size	Melting point and boiling point	Densities  (  )
Increase	Increase	Increase
-       Reason : Number of shells increases	-       Molecular size increases down the group. -        Thus, forces of attraction between molecules become stronger. -       Therefore, more energy is needed to overcome the strong force of attraction.	-       Increase in atomic mass is more than the increase in atomic radius (volume of atom)

Halogens	Physical state	Colour
Fluorine	Gas	Pale yellow
Chlorine	Gas	Greenish-yellow
Bromine	Liquid	Reddish-brown
Iodine	Liquid	Purplish- black

6. All the atoms of the halogens exhibit similar chemical properties because they have same valence electrons.

7. The reactivity of halogens decreases down the Group 17.

Explanation :

-       All halogens have seven valence electrons.

-       Each atom of the halogens will gain one more electron to achieve stable octet electron arrangement and form ion with a charge of negative.

Examples:

F + ē  → F^-

-       When going down to the group 17,  atomic size of halogens increases down the group

-       Thus, outermost occupied shell becomes further away from nucleus.

-       The strength to attract one electron into outermost shell by nucleus becomes weaker.

-       It causes the reactivity to decrease down the group

GROUP 18

Group 18

·         Group 18 elements are called noble gases

·         Elements= Helium, Neon, Argon, Krypton, Xenon, Radon

Physical properties:

·         Colourless gases at room temperature and pressure.

·         Low melting points and boiling points.

·         Low densities

·         Insoluble in water.

·         Do not conduct electricity

·         Poor conductors of heat

·         Exist as monoatomic gases at room condition

The physical properties of group 18 elements change down the group

a)      Atomic radius or atomic size increases.

Reason:

Number of occupied shells in the atom increases

b)      Melting point and boiling point increases.

Reason:

Forces of attraction between atoms become stronger as atomic size of elements increases.

c)       Density of elements increases

Reason:

Atomic mass increases more than the increase of atomic radius (volume of atom).

v  All group 18 elements are inert.

v  Inert means chemically unreactive.

v  Reason : All noble gases have the maximum number of valence electron on

outermost shell, which are two or eight valence electron.

v  Elements that has two valence electrons on the first shell is called duplet electron arrangement.

v  Elements that has eight valence electrons on the outermost shell is called octet electron arrangement.

The Uses of noble gases:

Helium	§  To fill airships and hot air balloons.
Neon	§  To fill neon lamps in signboards.
Argon	§  To fill an electric bulb.
Krypton	§  It is used in lasers to repair the retina of the eye. §  To fill the flash lights of cameras.