Year 10 Advanced Science – Chemistry Notes by Mr. E Hung @2010 May – revised 2013
1. Differences between Element, compound and mixture
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Element = a pure substance consists of only one kind of atom that can’t be
separated into simpler parts. Eg. Gold, Sodium, Oxygen, Mercury
Compound = a pure substance consists of two or more kinds of atoms (elements)
chemically bind together that can’t be separated by physical means.
Eg. carbon dioxide, sodium chloride, copper sulphate
Mixture is a substance consists of 2 or more substances that can be elements or
compounds. A mixture can be separated into its components by physical methods.
Eg. Air, Mineral water, milk, cake mix
A homogeneous mixture has the same uniform appearance and composition
throughout. Eg. sugar solution, mineral water, colloid
A heterogeneous mixture consists of visibly different substances or phases.
Eg. A suspension, emulsion
2. What is a molecule?
Molecule = the simplest structural unit of an element or compound. Some elements can
exit as a single atom, eg. Na, Ne, Au but others can’t. Eg Oxygen as O2, Nitrogen as N2,
Chlorine as Cl2. Each is the smallest molecule or structure of an element. A compound
consists of more than one kind of atom chemically combined together, eg. Carbon dioxide
as CO2, Copper sulphate, CuSO4.
3. Symbols for First 20 and 20 other
common elements:
H
He
Li Be B C N O F Ne
Na Mg Al Si P S Cl Ar
K Ca
Horrible Heary Little Begger Boys
Catch Newts On Friday Near
Naples, Magnificent Albert Sits
Planning Some Clever
Arrangement Keeping Calm.
Fe Ag Au Zn Pb Cu Ti U Br I Hg Ni W Ba Sr Cr Co Sn Cs Fr
Year 10 Advanced Chemistry Notes
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4. Electronic Configuration of the First 20 Elements
Period 1
Gp 1
Gp 2
Alkali
Metals
Alkaline
Earth
Metals
1
Gp 3
Gp 4
Gp 5
Gp 6
Gp 7
Gp 8
Halogens
Noble
Gases
2
H
He
(1)
Period 2
3
Li
(2, 1)
Period 3
11
Na
(2, 8, 1)
Period 4
19
K
(2, 8, 8, 1)
(2)
4
Be
(2, 2)
12
Mg
(2, 8, 2)
5
B
(2, 3)
13
Al
(2, 8, 3)
6
C
(2, 4)
14
Si
(2, 8, 4)
7
N
8
(2, 5)
15
P
(2, 8, 5)
O
9
(2, 6)
16
S
(2, 8, 6)
20
Ca
F
(2, 7)
17
Cl
(2, 8, 7)
10
Ne
(2, 8)
18
Ar
(2, 8, 8)
Br
(2, 8, 8, 2)
Sr
I
Nina
Pascoe
Ba
OutStanding
5. Atomic Structure – showing the Protons, Neutrons and Electrons (electronic configuration)
PEA (Proton number = Electron number = Atomic number)
Atomic Mass = Protons + Neutrons (Neutrons = Mass – Protons)
Magnesium (Mg) (2,8,2)
P=E=A=12
Nucleus:12P + 12N
Mass = P+N= 24
Carbon (C) (2,4)
PEA = 6
Nucleus: 6P + 6N
Mass = 6+6 = 12
Sodium (Na) (2,8,1)
PEA = 11
Nucleus: 11P + 12 N
Mass = 11+12 = 23
Argon (Ar) (2,8,8)
PEA = 18
Nucleus: 18P + 22N
Mass = 18+22 = 40
6. What are Isotopes?
 Isotopes are different types of atoms of the same chemical element, each having a
different number of neutrons.
 They differ in mass number but not in atomic number.
 The number of protons (the atomic number) is the same because that is what
characterizes a chemical element.
 Eg. carbon-12, carbon-13 and carbon-14 are three isotopes of the element carbon with
mass numbers 12, 13 and 14, respectively.
 The atomic number of carbon is 6, so the neutron numbers in these isotopes of carbon
are therefore 12−6 = 6, 13−6 = 7, and 14–6 = 8, respectively.
 The mass number of an atom is not a whole number because of the presence of isotopes.
 Eg. Chlorine has a mass number of 35.5 (Proton = 17, Neutron = 18 or 20)
Chlorine consists of atoms of isotopic masses 35 (75%) and 37 (25%)
Average Mass =( 35 x 0.75 ) + ( 37 x 0.25 ) = 35.5
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7. Electron Dot or Lewis Diagram:
 Scientists have devised a simple way to illustrate these outer shell electrons using
simple diagram. These electrons also called valence electrons.
 It is these valence electrons that make similar property of the elements within the
group.
8. Formation of Ionic Compounds
Mg
[Mg]2+[ O ]2-
O
One Magnesium atom donates 2 electrons, being accepted by one atom of oxygen.
(Two hands to two holes) MgO
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9. Formation of Covalent Compounds – Covalent Bonding
simply
simply
O=O
O=C=O
simply
10. Differences between Ionic and Covalent Compounds
Elements involved
Electrons
Exist as
Bonding
Melting and boiling points
Conducting electricity
Examples
Ionic Compounds
Between Metal and Non-metal
Donating and Receiving Electrons
Ions in water or solid lattice
Weaker forces between ions
But strong forces in solid state
Very high
Conducts electricity in molten or aqueous
form due to the presence on ions
Solid like Sodium chloride
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Covalent Compounds
Non-metals only
Sharing Electrons
Neutral molecules
Strong forces between atoms
But weak forces between molecules
Very low
Can’t conduct electricity – no ions –
neutral molecules only
Solid like glucose, liquid like water and gas
like carbon dioxide
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Each Na+ ion in NaCl is surrounded by six Cl- ions, and vice versa.
Removing an ion from this compound therefore involves breaking at
least six bonds.
Some of these bonds would have to be broken to melt NaCl, and
they would all have to be broken to boil this compound.
As a result, ionic compounds such as NaCl tend to have high
melting points and boiling points.
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Ionic compounds are therefore solids at
room temperature.
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Cl2 consists of molecules in which one
atom is tightly bound to another, as shown in the figure above.
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The covalent bonds within these molecules are at least as strong as
an ionic bond, but we don't have to break these covalent bonds to separate
one Cl2 molecule from another.
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As a result, it is much easier to melt Cl2 to form a liquid or boil it to
form a gas, and Cl2 is a gas at room temperature.
11. What are metallic bonds?
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In the metallic bond, an atom achieves a more stable
configuration by sharing the electrons in its outer shell with
many other atoms.
The valence electrons are not closely associated with
individual atoms, but instead move around amongst the atoms within the crystal.
Therefore, the individual atoms can "slip" over one another
yet remain firmly held together by the electrostatic forces
exerted by the electrons.
This is why most metals can be hammered into thin sheets
(malleable) or drawn into thin wires (ductile).
When an electrical potential difference is applied, the
electrons move freely between atoms, and a current flows.
12. Writing Chemical Formula for Ionic Compounds: M=Metal; N=Non-metal
Cation
Anion/ Radical
Chemical Formula
Examples
M 1+
N 1-
MN
NaCl, LiF, KI, NaOH, KNO3, NH4Cl
M 2+
N 2-
MN
MgO, CaS, MgSO4
3+
3-
MN
AlPO4
M
N
M 3+
N 2-
M2N3
Al2O3
M 2+
N 3-
M3N2
Ca3(PO4)2
M 2+
N-
MN2
BaBr2, MgCl2
M 1+
N 2-
M2N
Li2O, (NH4)2SO4
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D
Di
13. Giant Covalent Bonds: eg. Diamond, Graphite and Silicon Dioxide (glass/sand).
Silicon dioxide
Diamond
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very high melting points because all atoms are held firmly in place by strong covalent bonds.
In graphite each carbon atom is held in place by three strong covalent bonds which gives
graphite a high melting point.
In diamond 4 strong covalent bonds holds each atom in place. This also gives diamond a very
high melting point. The four bonds make diamond very hard.
Graphite has weak bonds between layers so the layers slip over each other making graphite soft.
Silicon dioxide has a high melting point - around 1700°C. Very strong silicon-oxygen covalent
bonds have to be broken throughout the structure before melting occurs.
14. Physical Properties of Metals:
 They are good conductors of heat and electricity.
 They are malleable and ductile in their solid state.
 They show metallic luster.
 They are opaque.
 They have high density (except lithium and sodium)
 They are solids (except mercury)
 They have crystal structure in which each atom is surrounded by eight to twelve near
neighbors
15. Chemical Properties of Metals:
 They have one to four valence electrons.
 They readily lose electrons to form Cations – forming ionic compounds.
 Metals can be easily oxidized. Oxidation is informally known as the process of corrosion.
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Iron + Oxygen + Water  Iron oxide (hydrated) = RUST
Fe(s) + O2(g) + H2O(I)  Fe2O3 . x H2O(s)
Metals can react with acid in a single displacement reaction to make hydrogen gas and
an aqueous solution of a salt.
Magnesium + Hydrochloric acid  Magnesium chloride + Hydrogen gas
Mg + HCl --------->

MgCl2
+
H2
Some reactive metals, like sodium or calcium will react with water to make a base.
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16. What are Transition Methods?
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A group of hard, tough metal elements, which conduct electricity.
Includes well known metals such as gold, silver, iron, copper and platinum.
There are 38 of them with 1 to 4 valence electrons,
Many have variable valence electrons:
Copper (I) & Copper (II), Iron (II) & Iron (III)
Very hard and strong
Higher densities
High melting points
High boiling points
High electrical conductivity
Malleable and Ductile
Not as reactive as Group 1 and Group 2 elements (metals)
Compounds with colours
From left to right, aqueous solutions of: Co(NO3)2 (red); K2Cr2O7 (orange);
K2CrO4 (yellow); NiCl2 (turquoise); CuSO4 (blue); KMnO4 (purple).
17. There are 7 evidences of chemical reactions:
(1) A change in colour
(2) Heat may be given out as you can feel the warmth of the test tube – exothermic
(3) Heat may be absorbed from the surrounding, giving you the feeling of cold - endothermic
(4) A Precipitate is formed from mixing 2 solutions
(5) Fizzing or bubbling implies a gas is formed.
(6) Sound may be given out.
(7) Light may be given out.
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18. Different types of Chemical Reactions:
Reaction type
What happens?
Combination
Or Synthesis
2 simple chemicals react together
to produce a more complicated
chemical. A+ B ----- > C
Combustion
(Burning)
The combination of a substance
with oxygen. The products are
carbon dioxide, water
CH4(g) + O2(g)  CO2(g) + H2O (l)
A larger molecule is broken down
into smaller or simpler chemicals.
D ------ > E + F
Calcium Carbonate  Calcium Oxide
+ Carbon dioxide
2CaCO3  2CaO + 2CO2
Decomposition
One Example
C2H5OH(g) + 2O2(g)  CO2(g) + 3 H2O (l)
Hydrogen peroxide
2H2O2  2H2O + O2
Single
Displacement
Double
Displacement
(Replacement)
An element replaces the metal part
of a compounds and the metal in
this compound is deposited as free
metal. (related to Metal Reactivity
Series)
M1 + M2N -------- > M1N + M2
Zinc + Copper Sulphate
 Zinc Sulphate + Copper
(The blue colour of the Copper
sulphate solution disappeared)
No reaction if it is:
Copper + Zinc sulphate (as copper is a
weaker metal than Zinc)
Two solutions with ions are mixing
up with the formation of new
products. Both cations and anions
are swapping around.
M1N1+ M2N2 -------- > M1N2 + M2N1
NB:
Neutralisation Reaction &
Precipitation Reaction are Double
Displacement reaction
Precipitation
Reaction
A precipitate (a solid) is formed
from mixing 2 solutions together.
NB: Precipitation Reactions are
Double Displacement reaction
Lead Nitrate + Potassium Iodide
 Lead Iodide + Potassium Nitrate
(Lead Iodide is a yellow precipitate)
Neutralisation
(Acid and
Base)
An acid reacts with an alkali to
form a salt and water
The hydrogen ion and hydroxyl ion
combine to form water
NB: Neutralisation Reactions are
Double Displacement reaction
Hydrochloric acid + Sodium hydroxide
 Sodium chloride + water
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19. Reactivity Series:
 By studying replacement reactions we can arrange the metals in decreasing order of
reactivity.
K
Na
Ca
Mg
Al
Zn
Fe
Pb
Cu
Ag
Au
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Most Reactive ---------------------------------------------------------------------- Least Reactive
aluminum will replace zinc in an aqueous solution of a zinc compound. Silver will not
replace tin. (Single Replacement Reaction)
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Summary of the Reactivity of metals with water and acid:
Stored in Oil
Metals
K, Na
Remarks
Extremely reactive; reacts vigorously with oxygen (air))
Reacting with water
K, Na, Ca
Highly reactive metals
Reacting with acid
K, Na, Ca
Mg, Al, Zn, Fe, Pb
Reactive metals
No Reaction
Gold Au
Unreactive – not reacting with oxygen
20. What are Alloys?
 A mixture of two or more metals, or a metal and a non-metal. For example, brass is an alloy
of copper and zinc, and steel is an alloy of iron and carbon.
 An alloy is created for its desirable properties that can’t be obtained from the pure metal.
 Some common Alloys used in our daily life:
21. Methods of preventing Corrosion of Metals:
1. Painting - to prevent contact between the metal and oxygen
2. Coating with Plastic, Oil, Grease or Tar - to prevent oxygen contact with the metal
3. Coating with Metal - Galvanising iron with zinc
4. Sacrificial Protection - enclosing a metal to be protected with a more reactive metal
which will corrode first
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