Part II The Microscopic World I/P.1
PART II: THE MICROSCOPIC WORLD I
I. Atomic Structure
A. Elements
Substances which cannot be broken down into anything simpler by chemical means are called elements.
e.g. Water (H2O) can be broken down into hydrogen and oxygen, therefore, water is not an element.
Oxygen cannot be broken down into other substances, therefore, oxygen is an element.
B. Names and symbols of some elements
It is useful to give each element a chemical symbol.
Chemical
Element
Latin Name
Element
Symbol
Aluminium
Al
Lead
Argon
Ar
Lithium
Barium
Ba
Magnesium
Beryllium
Be
Manganese
Boron
B
Mercury
Bromine
Br
Neon
Calcium
Ca
Nickel
Carbon
C
Nitrogen
Chlorine
Cl
Oxygen
Chromium
Cr
Phosphorus
Cobalt
Co
Platinum
Copper
Cuprum
Cu
Potassium
Fluorine
F
Silicon
Gold
Aurum
Au
Silver
Helium
He
Sodium
Hydrogen
H
Sulphur
Iodine
I
Tin
Iron
Ferrum
Fe
Zinc
Latin Name
Plumbum
Hydragyrum
Kalium
Argentum
Natrium
Stannum
C. States of Elements
Elements exist in different states at room temperature and pressure.
e.g.
Silver and sulphur are solids
Bromine and mercury are liquids
Chlorine is gas
D. Classification of elements
a.. In general, elements can be classified into metals and non-metals
The metals colours (e.g. sulphur is yellow; phosphorus is red or yellow; carbon is black)
Characteristics of Metals
Characteristics of Non-metals
a. Solids (except Mercury)
a. Gases or solids (except bromine)
b. Shiny
b. Dull
c. Good conductor of heat and electricity
c. Poor conductor of heat and electricity
d. Malleable and ductile
d. Brittle (in solid state)
e. Hard and strong
e. Not uniform in hardness and strength
Chemical
Symbol
Pb
Li
Mg
Mn
Hg
Ne
Ni
N
O
P
Pt
K
Si
Ag
Na
S
Sn
Zn
Part II The Microscopic World I/P.2
f. High melting and boiling points
g. High density
b. (i)
f. Usually low melting and boiling points
g. Low density
Some elements can be further classified as semi-metals or metalloids 類金屬.
(ii)
Metalloids have some properties similar to metals and non-metals e.g. boron, silicon and germanium
Semi-metals
Hardness and strength
Brittle
Melting point and boiling point
High
Appearance
Grey and shiny crystals, or brown powder
Electrical conductivity
Not conduct electricity;
When heated or impure= conduct electricity well
(iii) Most metalloids have important uses in industry. An example is silicon which is a semi-conductor. It is
wisely used in making computer chips.
Classwork
Study the following descriptions of three elements. Classify each as a metal, non-metal or metalloid. Explain you
choice in each case.
Element
Description
X
A yellow solid that melts at 119oC. Both the solid and liquid forms do not conduct electricity.
Y
A shiny solid which can be bent or hammered into shape easily.
Z
A shiny brittle solid which can conduct electricity
Answer
X is a non-metal element because it has a low melting point and it does not conduct electricity in both solid
and liquid state.
Part II The Microscopic World I/P.3
Y is a metal element because it is malleable.
Z is a metalloid because it is a brittle solid but it also can conduct electricity.
E. Basic Structure of an Atom
a. Fundamental sub-atomic particles
Atoms are made up of three fundamental subatomic particles - protons, neutrons and electrons.
(i)
The center of an atom is a very tiny and extremely dense region called the nucleus. The nucleus
contains protons and neutrons packed tightly together.
(ii)
Electrons are spinning very fast around the nucleus.
(iii)
There is empty space in-between the nucleus and electrons.
Sub-atomic
Particle
Proton
Neutron
Electron
Symbol
Relative mass
Relative charge
Position within the atom
p
n
e-
1
1
1/1837
+1
0
-1
inside nucleus
inside nucleus
move freely at great speed around
nucleus
b. Building Up Different Atoms
(i)
Different atoms have different numbers of protons, neutrons and electrons.
Atom
Symbol
Number of p Number of n
Number of eHydrogen
H
1
0
1
Electronic arrangement
1
Helium
He
2
2
2
2
Lithium
Li
3
4
3
2,1
Beryllium
Be
4
5
4
2,2
Boron
B
5
6
5
2,3
Carbon
C
6
6
6
2,4
Nitrogen
N
7
7
7
2,5
Oxygen
O
8
8
8
2,6
Fluorine
F
9
10
9
2,7
Neon
Ne
10
10
10
2,8
Sodium
Na
11
12
11
2,8,1
Magnesium
Mg
12
12
12
2,8,2
Aluminium
Al
13
14
13
2,8,3
Silicon
Si
14
14
14
2,8,4
Phosphorus
P
15
16
15
2,8,5
Sulphur
S
16
16
16
2,8,6
Chlorine
Cl
17
18
17
2,8,7
Argon
Ar
18
22
18
2,8,8
Potassium
K
19
20
19
2,8,8,1
Part II The Microscopic World I/P.4
Calcium
Ca
20
20
20
2,8,8,2
(ii) An atom is electrically neutral. This is because any atom always has equal numbers of protons and
electrons.
(iii) On the other hand, the number of neutrons may not be equal to that of protons.
F. Atomic Number and Mass Number
a.
Atomic Number (Z)
The atomic number = number of protons
For example, sodium has eleven protons in its nucleus and so its atomic number is equal to 11.
b.
Mass Number (A)
The mass number = protons+ neutrons
For example, a helium atom has two protons and two neutrons in its nucleus. Therefore, the mass number of
the helium atom is equal to 4.
Atomic number
Mass number
=
Number of protons
=
Number of electrons (because the atom is electrically neutral)
=
Number of protons + Number of neutrons
=
Atomic number + Number of neutrons
Therefore, if we want to calculate the number of neutrons in an atom, we can do the following subtraction:
Number of neutrons
=
Mass number - Number of protons
To be more convenient, atomic number and mass number of an atom are usually expressed as a simplied
notation.
For example, 208
82 Pb
Which represents a lead atom having 82 protons, 82 electrons and (208-82) = 126 neutrons.
G. Isotopes
a. Definition: Isotopes are atoms of the same element which have different numbers of neutrons.
e.g.
35
17 Cl
and
37
17 Cl
are the two isotopes of chlorine.
b. Properties of Isotopes
(i) Isotopes have the same chemical properties because they have the same number of protons and
outermost shell electrons.
(ii) Isotopes have different physical properties because they have different number of neutrons.
c. Relative isotopic mass
(i) The relative isotopic mass of a particular isotope of an element is the mass of one atom of that isotope
on the 12C = 12.00 scale.
Part II The Microscopic World I/P.5
The
12
C isotope has been chosen as the reference standard of mass. One
12
C atom is given a
relative mass of exactly 12.00. Masses of all other atoms are compared with the reference
standard to give their relative masses.
(ii) 1. The mass of a hydrogen atom ( 11 H ) is equal to
1
of the mass of a carbon-12 atom, so is relative
12
isotopic mass is 1.
24
2. A magnesium atom ( 12
Mg ) is twice as heavy as a carbon-12 atom, so its relative isotopic mass is
24.
Relative isotopic mass = mass number
H. Relative Atomic Mass
(i) The relative atomic mass of an element is the weighted average of the isotopic masses of its natural
isotopes on the 12C = 12.00 scale.
Example
a. Chlorine consists of two natural isotopes,
35
17 Cl
and
37
17 Cl ,
with percentage abundance of 75.4% and 24.6%
respectively. Calculate the relative atomic mass of chlorine.
Ans: 35.5
b. Neon in the air contains two isotopes:
the relative abundance of the isotopes.
20
Ans: 10
Ne (90%); 1022 Ne (10%)
20
10
Ne and 1022 Ne . The relative atomic mass of neon is 20.2. Calculate
(ii) Relative atomic mass has no unit.
(iii) Relative atomic masses of some common elements
(iv) Relative atomic mass = a% x M(A) + b% x M(B) + c% x M(C)
Abc% =abundance of isotopes
M abc = isotopic masses
Element
Symbol
Aluminium
Al
Calcium
Ca
Chlorine
Cl
Copper
Cu
Hydrogen
H
Iron
Fe
I. Electronic Arrangement of Atoms
Relative atomic
mass
27.0
40.1
35.5
63.5
1.0
55.8
Element
Symbol
Magnesium
Oxygen
Potassium
Silver
Sodium
Sulphur
Mg
O
K
Ag
Na
S
Relative atomic
mass
24.3
16.0
39.1
107.9
23.0
32.1
a. Electron shells
(i) Electrons in an atom exist in a number of regions (called electron shells) surrounding the central nucleus.
Part II The Microscopic World I/P.6
(ii) Each electron shell is given a number 1, 2, 3, 4 and so on, starting from the one closest to the nucleus (i.e.
the innermost shell). Each shell can hold up to a certain maximum number of electrons.
Shell Number, n
Maximum number of electrons (= 2n2)
1
2
2
8
3
18
4
32
.
.
(iii) Electrons in an atom are arranged into shells.
The distribution of electrons in the various shells is called Electronic Arrangement (or electronic
configuration)
For example, a sodium atom has 11 electrons. The electronic arrangement of a Na atom is
2 , 8, 1
no. of
electrons in:
1st shell 2nd shell 3rd shell
Note: The writing of this notation starts from left to right.
(iv) The electronic configuration of the first 20 elements
Element
Symbol
Atomic number
Hydrogen
H
1
Electronic configuration
1
Helium
He
2
2
Lithium
Li
3
2,1
Beryllium
Be
4
2,2
Boron
B
5
2,3
Carbon
C
6
2,4
Nitrogen
N
7
2,5
Oxygen
O
8
2,6
Fluorine
F
9
2,7
Neon
Ne
10
2,8
Sodium
Na
11
2,8,1
Part II The Microscopic World I/P.7
Magnesium
Mg
12
2,8,2
Aluminium
Al
13
2,8,3
Silicon
Si
14
2,8,4
Phosphorus
P
15
2,8,5
Sulphur
S
16
2,8,6
Chlorine
Cl
17
2,8,7
Argon
Ag
18
2,8,8
Potassium
Ka
19
2,8,8,1
Calcium
Ca
20
2,8,8,2
(v) Electronic diagrams of atoms
Notes:
1. The nucleus is represented by the symbol of the atom.
2. Electronic shells are represented by concentric circles around the nucleus.
3. Electrons are represented by dots or crosses.
Part II The Microscopic World I/P.8
II. Periodic Table
A. Groups and Periods
a. The Periodic Table is an arrangement of elements in an order of increasing atomic number.
b. It is divided into:
(i) Vertical columns called groups. (I,II,III….0)
(ii) Horizontal rows called period. Period 1-7
c. Group number = number of outermost shell electrons of the atoms
Period number = number of occupied electron shells of the atoms
d. Group names
Group
Group name
I
The alkali metals
II
The alkaline earth metals
VII
The halogens
0
The noble gases
e. Elements having similar chemical properties are put together in the same group.
Elements of the same group have similar chemical properties because it depends mainly on the number of
outermost shell electrons.
B. Patterns across the Periodic Table
Some Properties of the Elements in Period 3
Element
Sodium Magnesium Aluminium Silicon Phosphorus Sulphur Chlorine Argon
State at room
temperature and
Solid
Gas
pressure
Melting Point
98
650
660
1410
44
113
-101
-189
(oC)
Boiling Point
890
1120
2450
2680
280
445
-34
-186
(oC)
Electrical
Good
Moderate
Poor
Conductivity
Type of Element
Reactivity
Metal
Reactive
Metalloid
Non-Metal
Moderately Very
Very Extremely
Moderately reactive
reactive unreactive
reactive unreactive
a. The elements change from metals through metalloid to non-metals.
Part II The Microscopic World I/P.9
b. The reactivity of the elements also changes across a period. Apart from the noble gases, the most reactive
elements are near the edges of the periodic table and the least reactive ones are in the center.
c.
The table above shows the different blocks of elements in the periodic table.
Elements near the zig-zag line are metalloids, for example, boron, silicon and germanium. Elements
between Groups II and III are transition metals.
C. Properties of Elements
a. Group I Elements – Alkali Metals → Silvery solids
The six elements in Group I are Lithium, Sodium, Potassium, Rubidium, Caesium and Francium. These
elements react with water to form alkalis. Therefore they are called the alkali metals.
(i)
Similarities of Group I elements
1. relatively low melting points and boiling points when compared with other metals.
2. soft and can be cut with a knife.
3. low densities – they can float on water
4. reactive metals and reacting with air readily, must be stored in paraffin oil
5. react with non-metals to form ionic compounds.
6. react with water to give hydrogen and form an alkaline solution.
(ii) Differences in reactivity of Group I elements
Group I elements are all very reactive. The reactivity increases as we move down the group
Element
Reaction with water
Lithium
Float, produce hydrogen gas
Sodium
Melt→form silvery ball which moves abt quickly on the water, produce hydrogen gas rapidly
Part II The Microscopic World I/P.10
Potassium
Rubidium
Caesium
Melt→form silvery ball which moves abt quickly on the water, produce hydrogen gas which
catches fire itself
Reacts vigorously than potassium
Reacts vigorously than Rubidium
b. Group II Elements – Alkaline Earth Metals→Silvery solids
The six elements in Group II are Beryllium, Magnesium, Calcium, Strontium, Barium and Radium. They
are called as alkaline earth metals.
(i)
Similarities of Group II elements
1. all have relatively low m.p. and b.p. when compared with other metals (except Group I metals)
2. all have low densities but denser than grp I
3. Less reactive than alkali metals
4. Except beryllium react with water less vigorously than grp I
5. all reactive metals and react readily with dilute hydrochloric acid to give hydrogen gas.
6. All react with non-metals to form ionic compounds
(ii) Differences in reactivity of Group II elements
Group II elements are less reactive than Group I elements. The reactivity increases as we move down the
group.
Element
Reaction with water
Beryllium
Not react with water or steam
Magnesium
Almost no reaction with cold water but react with steam to produce hydrogen
Calcium
Reacts steadily with cold water to produce hydrogen
Strontium
Reacts vigorously with cold water
Barium
Reacts even more vigorously than strontium does
c. Group VII Elements – Halogens
Group VII elements include Fluorine, Chlorine, Bromine, Iodine and Astatine →Halogens.
(i)
Similarities in properties of Group VII elements
1. All are reactive
2. All react with metals to form ionic compounds
3. All react with non-metals to form covalent compounds
4. All coloured. They become darker in colour down the grp
Fluorine
Pale yellow gas
Chlorine
Greenish-yellow gas
Bromine
Reddish brown liquid
Iodine
Black solid
Astatine
Black solid
(ii) Differences in reactivity of Group VII elements
The reactivity of these elements decreases as we move down the group.
Element
React with hydrogen
Fluorine
Reacts explosively even in the dark
Chlorine
Reacts explosively in sunlight but reaction is slow in the dark
Bromine
Reacts only in sunlight or when heated
Iodine
Virtually no reaction even in direct sunlight or with strong heating
d. Group 0 Elements – Noble Gases→colourless gases
(i) The six elements in Group 0 are Helium, Neon, Argon, Krypton, Xenon and Radon. They are called noble
gases because they rarely react with other substances.
Part II The Microscopic World I/P.11
(ii) Similarities
1.
All are colourless gases at room temperature and pressure
2.
All are very unreactive, they have little or no reaction with other elements
3. The noble gases are stable because their outermost shells are full of electrons.
Helium Very light and does not burn, it is use to fill balloons and airships
Argon Does not react with metals but use to fill light bulbs as it does not react with metal filament inside the
light bulb
(iii)Electron arrangement
Helium (He)
2
Neon (Ne)
2,8
Argon (Ar)
2,8,8
Krypton (Kr)
2,8,18,8
Xenon (Xe)
2,8,18,18,8
Radon (Rn)
2,8,18,32,18,8
Octet →8 electrons in the outermost shell
Duplet →2 electrons in the only one occupied shell
*atoms of elements other than noble gases usually not stable→attain an octet or duplet structure to become stable
D. Predicting the chemical properties of unfamiliar elements (Extension)
Because of the similar chemical properties of elements of the same group, it is possible to predict the chemical
properties of unfamiliar elements.
Example
1. Caesium is a Group I element below potassium in the Periodic Table.
a. How many outermost shell electron(s) is/are there in a caesium atom? Explain your answer.
b. Predict the state of caesium at room temperature and pressure.
c. Predict the observation when caesium reacts with cold water.
d. Which of the metals, potassium or caesium, is more reactive?
e. Suggest ONE method to store caesium safely in the laboratory.
Answer
a. 1 outermost shell electron. Because it is a Group I element. Group I elements have one
outermost shell electron.
b. Solid
c. It floats and moves on the surface of water.
It reacts with water vigorously and gives a colourless gas (hydrogen)
Coloured flame may be seen.
“Hissing” sound is heard.
d. Caesium is more reactive.
e. Caesium should be stored in paraffin oil.
2. Astatine is a Group VII element below iodine in the Periodic Table.
a. How many outermost shell electron(s) is/are there in an atom of astatine?
b. Predict the state of astatine at room temperature and pressure.
c. Predict whether astatine is poisonous.
Part II The Microscopic World I/P.12
Answer
a. 7 outermost shell electrons
b. Solid
c. It is poisonous
Part II The Microscopic World I/P.13
III. Chemical Bonds
A. Investigating electrical conductivity of substances
Conductors
Electrolysis
Electrical Conduct electricity in both solid Conduct electricity only when
properties and liquid states
molten or in aqueous solution and
decomposed by electricity during
conduction
Examples All metals: copper, iron
Compound of metals and nm:
1 non-metals: carbon(graphite)
Sodium chloride
Copper(II) sulphate
Lead(II) bromide
Nm compounds:
Hydrogen chloride
Ethanoic acid
Citric acid
Non-conductors
Not conducting electricity
All non-metals except
Graphite:
Iodine
Sulphur
Carbon (diamond)
Compound of nm:
Distilled water
Oil
ethanol
A chemical bond refers to the strong electrostatic attraction (i.e. attraction between opposite charges) holding
atoms or ions together.
B. Classification of Chemical Bonds
Ionic bonds
Covalent bonds
Metallic bonds
Type of atoms
involved
formed between metal
atoms and non-metal
atoms
formed between nonmetal atoms
formed between metal
atoms
Way to obtain
stability
by transfer of electrons
to form cations and
anions
by sharing of electrons
Nature of the
bonding
electrostatic attraction
between oppositely
charged ions
electrostatic attractions
between nuclei and
shared electrons
metal atoms lose their
outermost electrons to
form sea of electrons and
positive metal ions
electrostatic attraction
between sea of electrons
and metal ions
Part II The Microscopic World I/P.14
IV. Ionic Bond
A. Formation of Ions
a. Positive ions / Cations= lose electrons
(i) Sodium atom - which has an electronic arrangement of 2, 8, 1 - tends to lose one electron from its
outermost shell in order to achieve the stable electronic arrangement of the nearest noble gas, neon (2, 8).
The atom becomes positively charged when the number of protons it possesses is greater than the number
of electrons. Positive ion or cation is thus formed.
(ii) A sodium ion carries 1 positive charge and is represented by the symbol Na+. The “+” sign means 1
positive charge.
(iii) Another example
A calcium ion carries 2 positive charges and is represented by the symbol Ca 2+. The “2+” sign means 2
positive charges.
b. Negative ions/ Anions=gaining electrons
(i) For a chlorine atom with an electronic arrangement of 2, 8, 7, it tends to gain one electron to achieve the
stable electronic arrangement of the nearest noble gas, argon (2, 8, 8)
The atom becomes negatively charged when the number of protons it possesses is smaller than the
number of electrons. Negative ion or anion is thus formed.
(ii) A chloride ion carries 1 negative charge and is represented by the symbol Cl -. The “ - ” sign means 1
negative charge.
(iii) Another example
An oxygen atom has an electronic arrangement of 2,6. It tends to gain two electrons in order to get the
stable electronic arrangement of a neon atom (2,8)
Part II The Microscopic World I/P.15
When an oxygen atom gains two electrons, an oxide ion forms. It carries 2 negative charges and is
represented by the symbol O2-. The “2-“ sign means 2 negative charges.
Notes:
1. Metals usually have one, two or three outermost shell electrons and they usually form ions of charge +1, +2
and +3 respectively.
2. Non-metals such as Group V, VI or VII usually form ions of charge -3, -2 and -1 respectively.
3. Ions may be formed from simple atoms such as Na+, K+, Cl- and I- which are called simple ions. Those formed
from two or more atoms such as OH-, NO3-, NH4+ are called polyatomic ions.
B. Ionic Bonds
a. When sodium and chlorine react together, the sodium atom loses one electron to the chlorine atom. This
transfer of electron results in the formation of two ions, Na+ and Cl-.
b. The electronic diagram ("dot and cross" diagram) shows the transfer of electrons. Ions are put inside square
brackets with the charge written at the right hand corner.
You should note that all electrons are identical. The dots and crosses are symbols only.
c. The opposite charges of sodium ion and chloride ion attract each other strongly. This type of attractive force is
called ionic bond.
Part II The Microscopic World I/P.16
Ionic bond is a strong electrostatic attraction between oppositely charged ions which are formed by
transfer of electrons from one atom (or group of atoms) to another.
d. A simplified electronic diagram (showing the outermost shell electrons only) of a compound is as follows.
xx
Na
Cl
x
+
x
x
Na
-
xx
x
xx
x
x
Cl
xx
Classwork
Draw electronic diagrams to show bond formation between the following elements
1. potassium and sulphur
2. aluminium and oxygen
3. lithium and oxygen
C. Names and formulae of common ions
Positive Ions
+1 ions
Name
+2 ions
Symbol
Name
+3 ions
Symbol
+
Magnesium ion
Mg
+
Calcium ion
Ca
Barium ion
Ba
+
lead(II) ion
Pb
+
copper(II) ion
Cu
+
Zinc ion
Zn
iron(II) ion
Fe
mercury(II) ion
Hg
manganese(II) ion
Mn
Lithium ion
Li
Sodium ion
Na
Potassium ion
K
Silver ion
Ag
copper(I) ion
Cu
mercury(I) ion
Hg
Hydrogen ion
H
Ammonium ion
NH4
+
+
+
Name
Symbol
2+
aluminium ion
Al
2+
iron(III) ion
Fe
2+
chromium(III) ion
Cr
2+
2+
2+
2+
2+
2+
3+
3+
3+
Part II The Microscopic World I/P.17
cobalt(II) ion
Co
nickel(II) ion
Ni
2+
2+
Negative ions
-1 ions
-2 ions
-3 ions
Name
Symbol
Name
Symbol
Fluoride ion
FCl-
Oxide ion
O
Sulphide ion
S
Sulphate ion
SO4
Iodide ion
BrI-
Sulphite ion
SO3
Hydroxide ion
OH-
Thiosulphate ion
S2O3
Nitrate ion
NO3
Chromate ion
CrO4
Hydrogencarbonate
ion
Hydrogensulphate
ion
Permanganate ion
HCO3
Dichromate ion
Cr2O7
HSO4-
Carbonate ion
CO3
MnO4ClO-
Silicate ion
SiO3
Chloride ion
Bromide ion
-
Hypochlorite ion
-
Name
Symbol
2-
Nitride ion
N
2-
Phosphide ion
P
Phosphate ion
PO4
2-
3-
33-
2222-
2-
2-
D. Chemical names and Chemical formulae of ionic compounds
a. Writing chemical formulae for ionic compounds
(i) When sodium and chlorine react to form a compound, there must be one sodium ion (Na+) for each
chloride ion (Cl-).
+
-
Therefore, we can represent the compound by the chemical formula Na Cl or simply NaCl.
(ii) NaCl is the simplest chemical formula which gives the simplest ratio of the number of atoms or ions
present in the compound. It is also known as the empirical formula of the compound.
(iii) In writing the chemical formula of an ionic compound by combining the positive and negative ions. The
net charge for the compound must be zero.
(iv) Examples
1. Calcium chloride
Calcium ions (Ca2+) carries 2 positive charges. Chloride ion (Cl-) carries 1 negative charge. The net
charge for the compound must be zero.
Therefore, the simplest ratio of Ca2+: Cl- in the compound should be 1:2. The chemical formula of
calcium chloride is CaCl2.
2. Magnesium hydroxide
Magnesium ion (Mg2+) carries 2 positive charges and hydroxide ion (OH-) carries 1 negative charge.
The net charge for the compound must be zero.
Therefore, the simplest ratio of Mg2+:OH- in the compound should be 1:2. The chemical formula of
magnesium hydroxide is Mg(OH)2.
Notice that brackets should be used for a polyatomic ion if the number of that ion in the chemical
formula is 2 or more.
e.g.
calcium hydroxide Ca(OH)2,
Part II The Microscopic World I/P.18
aluminium hydroxide
ammonium sulphate
Al(OH)3,
(NH4)2SO4
For simple ions, there is no need to use brackets, even when the number of that ion in the
formula is 2 or more.
sodium sulphide Na2S,
e.g.
aluminium oxide
Al2O3,
magnesium chloride
MgCl2
Classwork
Write down the chemical formulae of the following ionic compounds
Ionic Compound
Chemical formula
Ionic Compound
sodium hydroxide
copper(II) sulphate
potassium sulphide
potassium permanganate
calcium chloride
Chemical formula
sodium
hydrogencarbonate
aluminium oxide
iron(III) chloride
calcium oxide
magnesium nitrate
magnesium sulphate
aluminium sulphate
zinc nitrate
ammonium sulphate
potassium phosphate
copper(I) oxide
calcium carbonate
lead(II) carbonate
b. Naming of Ionic Compounds
In naming of ionic compounds, the positive ion is named first, followed by the negative ion.
For example, a compound consists of sodium ions and chloride ions is named as sodium chloride. Further
examples are given below:
PbBr2
MgF2
CuCO3
lead(II) bromide
magnesium fluoride
copper(II) carbonate
Name the following compounds
Formula
Chemical name
Formula
KOH
Fe2O3
LiF
CuI
Al(NO3)3
NaHCO3
CuCO3
KHSO4
K2Cr2O7
NH4Cl
Chemical name
Part II The Microscopic World I/P.19
AgCl
Na2SO3
FeSO4
MgBr2
NiCO3
CoCl2
Zn(NO3)2
BaSO4
Note: It is important to determine the names of the positive ions and negative ions of the ionic compounds.
E. Colours of ions and ionic compounds
a. Many ions are colourless. However, some ions are coloured.
b. The colour of an ion may be deduced by observing the colour of solutions of a series of compounds.
Activity - To observe solutions of a series of compounds and deduce colours of some ions
Some common compounds are shown. Record the colour of every compound in aqueous state under its
formula in the following table and deduce the colour of the ions.
Example:
(i) As the aqueous (NH4)2CO3 is colourless, therefore, the colour of NH4+ and CO32- are both colourless.
(ii) As the aqueous CuCO3 is blue, and the colour of CO32- is colourless, therefore, the colour of Cu2+ is blue.
Ammonium
ion
(colourless)
Copper(II)
ion
(blue)
Iron(II) ion
(
)
Iron(III)
ion
(
)
Potassium
ion
(
)
Sodium ion
Ion
(
)
Nickel(II)
Ion
(
)
Carbonate
ion
(colourless)
Chloride
Ion
(
)
Sulphate
Ion
(
)
Nitrate
Ion
(
)
Dichromate
ion
(
)
(NH4)2CO3
(colourless)
NH4Cl
(colourless)
(NH4)2SO4
(
)
NH4NO3
(
)
(NH4)2Cr2O7
(
)
CuCO3
(blue)
CuCl2
(blue)
CuSO4
(
)
Cu(NO3)2
(
)
Permanganate
ion
(
)
FeSO4
(
)
FeCl3
(
)
K2CO3
(
)
(
Na2CO3
(
)
(
Fe(NO3)3
(
)
)
K2SO4
(
)
KNO3
(
)
K2Cr2O7
(
)
)
Na2SO4
(
)
NaNO3
(
)
Na2Cr2O7
(
)
KCl
NaCl
NiSO4
(
)
c. Colours of some common ions in aqueous solutions:
Name
Symbol for ion
copper(II) ion
Cu2+
iron(II) ion
Fe2+
iron(III) ion
Fe3+
cobalt(II) ion
Co2+
nickel(II) ion
Ni2+
chromium(III) ion
Cr3+
chromate ion
CrO42dichromate ion
Cr2O72-
Colour
blue or green
pale green
brown or yellow
Pink
green
green
yellow
orange
KMnO4
(
)
Part II The Microscopic World I/P.20
manganese(II) ion
permanganate ion
Mn2+
MnO4-
very pale pink
purple
d. Notice that the transition metals usually form coloured ions, which may be cations (e.g. Cu 2+ ion) or
polyatomic anions (e.g. permanganate ion MnO4-).
On the other hand, elements in main groups (Gp 1 to Gp 0) in the Periodic Table form colourless ions.
e. Colour of ions in gemstones
Gemstone
Amethyst
Colour
Purple
Ion responsible
Manganese(III) ion, Mn3+
for colour
Gemstone
Colour
Ion responsible
for colour
Peridot
Light green
Iron(II) ion, Fe2+
Emerald
Green
Chromium(III) ion, Cr3+
Topaz
Yellow
Iron(III) ion, Fe3+
Jade
Green
Chromium(III) ion, Cr3+
Turquoise
Bluish green
Copper(II) ion, Cu2+
f. Example
A student used the following set-up to study the movement of ions.
filter paper moistened with tap water
microscope slide
B
A
+
C
-
d.c. power supply
The student placed a drop of copper(II) sulphate solution at A and a drop of orange solution at C. The two
solutions do not react.
a. The orange colour of the solution at C is due to the anion present. Name the ion responsible for the colour.
b. Electricity was passed through for some time.
(i) What would be the colour change at A? Explain your answer.
(ii) What would be the colour change at B? Explain your answer.
Answer
a. Dichromate ion Cr2O72b. (i) The blue colour fades. The blue copper(II) ions move towards the negative electrode.
(ii) A green colour appeared. The orange negative ions move towards the positive electrode
while the blue positive ions move towards the negative electrode. They mix to give a
green colourn at B.
c. Migration of copper(II) ions and dichromate ions
*The gel slows down the mixing of the bottom layer in the U-tube with the top layer (dilute hydrochloric
acid)
Observation:
Part II The Microscopic World I/P.21
The bottom of the U-tube is filled with a gel containing copper(II) ions and dichromate ions. After the
electric current has passed for some time, a blue colour gradually appears in the solution around the
negative electrode. An orange colour gradually appears in the solution around the positive electrode.
The migration of ions under the influence of electric field provides further evidence that ions exist in
some compounds.
F. Ionic bonding and ionic compounds
Sodium lose 1 eChlorine gain 1 e-
V. Covalent Bonds
A. Formation of Covalent Bonds by Sharing of electrons
a. Let us take chlorine as an example.
The chlorine atom, Cl, is very unstable. Its outermost shell contains only 7 electrons. Electron transfer between
chlorine atoms cannot occur here, as they all tend to gain electron, and no one is willing to lose it.
However, by sharing of electrons, a chlorine molecule is formed, in which each chlorine atom gas a stable
octet in the outermost shell. Cl2
Part II The Microscopic World I/P.22
b. Electronic diagram showing sharing of two electrons in the formation of a chlorine molecule* (only outermost
shell electrons are shown):
*Non-metal atoms join together by sharing of electrons to form a group which is called as a Molecule 分子.
c. Definition:
Covalent bond is the strong directional electrostatic attraction between shared electrons (negatively
charged) and two nuclei (positively charged) of the bonded atoms.
d. Molecular formulae and Structural Formulae
(i) A shared pair of electrons makes a single covalent bond.
It is often denoted by a stroke (⎯) between the atomic symbols.
e.g. A chlorine molecule Cl2 can be written as “Cl─Cl ”
(ii) Cl2 is the molecular formula of chlorine, while Cl⎯Cl is the structural formula of chlorine.
1. The molecular formula of a molecular substance is the formula which shows the actual number of each
kind of atom(s) in one molecule of the substance.
2. The structural formula of a molecular substance is the formula which shows how the atoms are joined up
in one molecule of the substance.
(iii) Generally, when we say the formula of a molecular substance, we mean its molecular formula.
B. Covalent bond formation in some molecules
a. Hydrogen molecule
In the molecule, each hydrogen atom forms a duplet (not an octet) which is the stable electronic configuration
of the noble gas helium.
A single covalent bond is formed with the sharing of 1 pair of electrons between the two atoms.
Molecular formula:
H2
Structural formula:
H－H
b. Oxygen molecule
An oxygen atom contains 6 electrons in outermost shell→ need 2 more electrons to be stable neon atom. A
Part II The Microscopic World I/P.23
double covalent bond is formed with the sharing of 2 pairs of electrons between the two atoms. Each oxygen
atom has a stable octet.
An oxygen molecule can be shown as O = O.
Molecular formula:
Structural formula:
O=O
O2
c. Nitrogen molecule
A nitrogen atom contains 5 electrons in its outermost shell→ need 3 more electrons to get stable neon atom. A
triple covalent bond is formed with the sharing of 3 pairs of electrons between the two atoms.
A nitrogen molecule can be shown as N≡N. Each nitrogen atom has a stable octet. N3
Molecular formula:
Structural formula:
N2
NN
d. Hydrogen chloride molecule
After reaction, the hydrogen atom forms a duplet (not an octet) which is the stable electronic configuration of helium.
Molecular formula:
HCl
Structural formula:
H－Cl
e. Water molecule
There are two unshared pairs of electrons in the valence shell of oxygen. The unshared pairs also called lone
pairs.
Molecular formula:
H2O
f. Ammonia molecule
H－O－H
Structural formula:
Part II The Microscopic World I/P.24
Molecular formula:
Structural formula:
g. Tetrachloromethane molecule 四氯甲烷
Molecular formula:
Structural formula:
h. Carbon dioxide molecule 二氧化碳
Molecular formula:
Structural formula:
C. Writing chemical formulae of covalent compounds
We can use the following steps to work out the chemical formulae (molecular formulae) of covalent compounds.
Example 1
Step
1. Write down the electronic configurations of the atoms
involved
Compound formed from
hydrogen and oxygen
H
1
O
2,6
Part II The Microscopic World I/P.25
2. Decide the number of electrons needed to get a stable
electronic arrangement.
hydrogen atom needs 1 electron, while
oxygen atom needs 2 electrons
1
H
2
O
3. Decide the number of each type of atoms in one molecule 1
(cross multiply the numbers and the symbols).
H
2
O
=H2
4. Combine the symbols and simplify the ratio if necessary.
Example 2
Step
=O1
H2O
(Omit the number of 1 for oxygen)
Compound formed from
carbon and hydrogen
1. Write down the electronic configurations of the atoms
involved
C
2,4
2. Decide the number of electrons needed to get a stable
electronic arrangement.
carbon atom needs 4 electrons, while
hydrogen atom needs 1 electron
H
1
4
C
3. Decide the number of each type of atoms in one molecule 4
(cross multiply the numbers and the symbols).
C
=C1
4. Combine the symbols and simplify the ratio if necessary.
1
H
1
H
=H4
CH4
(Omit the number of 1 for carbon)
Classwork
Draw the electronic diagrams (showing the valence shell only) and give the molecular formulae, structural
formulae for the molecules formed by
a. F atoms
d. H and Si atoms
g. H and Br atoms
j. N and F atoms
b. Br atoms
e. H and P atoms
h. C and S atoms
k P and Cl atoms
c. I atoms
f. H and S atoms
i. C and Cl atoms
l. P and Br atoms
Part II The Microscopic World I/P.26
D. Dative Covalent Bond (Coordinate bond) 配價鍵
The two electrons which form a covalent bond do not necessarily have to come from each atom; both
may originate from one of the atoms.
Dative covalent bond occurs when one atom provides both of the electrons necessary for the
formation of a single covalent bond. Once the dative bond is formed it is indistinguishable from other
covalent bonds of the same type.
a. Ammonium ion NH4+
When ammonia reacts with hydrogen chloride to form ammonium chloride, a dative covalent
bond is formed between the lone pair of electrons on the N atom in NH3 and a H+ ion from HCl.
The symbol “→” is used to represent the dative covalent bond.
Part II The Microscopic World I/P.27
b. Hydronium ion (H3O+)
When an acid is dissolved in water, hydrogen ions H+ are formed.
The H+ ion is attracted to the unsharded electrons of oxygen atom of a water molecule, forming a
dative covalent bond. A more stable ion, hydronium ion H3O+, is obtained.
Example
HKCEE 1995 Q4 (Essay Question)
“When atoms combine, they tend to attain noble gas electronic structures.”
Discuss how atoms can attain the noble gas electronic structure. In your answer, you should give suitable
examples and the electronic structures of the products formed.
(8 marks)
Part II The Microscopic World I/P.28
Answer
Chemical Knowledge (5 marks)
Atoms become stable by attaining noble has electronic structures. They can achieve that by losing,
gaining or sharing of electrons.
Consider the formation of chlorine molecule from two chlorine atoms. A chlorine atom has an
electronic configuration of 2,8,7. It needs one more electron to obtain the noble gas electronic structure.
In order to obtain stable electronic structure, each chlorine atom will give one electron for sharing.
xx
x
x
Cl
xx
+
x
Cl
x
x
xx
Cl
x
Cl
xx
Consider the formation of sodium chloride from sodium atom and chlorine atom. A sodium atom has an
electronic configuration of 2,8,1. It can obtain a noble gas electronic structure by losing one electron. A
chlorine atom has an electronic configuration of 2,8,7. It can obtain a stable electronic structure by
gaining one electron.
When sodium and chlorine react, one sodium atom will transfer one electron to the chlorine atom to
form sodium ion (a positive ion) and chloride ion (a negative ion).
+
xx
Na
x
+
Cl
x
x
Na
x
x
-
x
Cl
xx
Effective Communication (3 marks)
A. The ability to present ideas in a precise manner (this mark should not be awarded to answers which
contains a lot of incorrect / superfluous materials)
B. The ability to present ideas in a systematic manner (i.e. the answer is easy to follow)
C. The ability to present answer in paragraph form and to express ideas using full sentences.
Part II The Microscopic World I/P.29
VI. Metallic Bond
a. Metallic structure and Metallic bond
(i) A metal consists of atoms packed closely together.
The loosely held outermost shell electrons get separated from their atoms.
b. The electrostatic attraction between a ‘sea’ of negatively charged electrons (negatively charged) and
positively charged metals ions (positively charged) is called metallic bond. They are in non-directional but
same in all directions.
c. Metals conduct electricity
1. metals are good conductors of electricity
2. the delocalized electrons move freely in all directions after connected to a battery→the delocalized electrons
move in one direction only
3. the delocalized electrons move towards the positive pole of the battery, leaving the metal
4. an equal number of electrons move into the other end of the metal from the negative pole
5. any moment, the no. of e- in the metal remains unchanged when conduct electricity there is no chemical change
Part II The Microscopic World I/P.30
VII. Relative Molecular Mass and Formula Mass
A. Relative Molecular Mass
a. Just as the relative atomic mass is used to describe the relative masses of atoms, the relative molecular mass
is used to describe the relative masses of molecules.
b. Relative molecular mass of an element or compound
= Sum of relative atomic masses of all atoms present in a molecule of the element or compound
c. Example
Given: Relative atomic mass of H = 1.0, C=12.0, N = 14.0, O = 16.0, Cl = 35.5
1. Relative molecular mass of nitrogen molecule (N2)
= 14.0  2
= 28.0
2. Relative molecular mass of water molecule (H2O)
= 1.0 + 16.0  2
= 18.0
3. Relative molecular mass of carbon dioxide (CO2)
=
4. Relative molecular mass of chlorine molecule (
)
=
5. Relative molecular mass of hydrogen chloride (
)
=
6. Relative molecular mass of ammonia (
)
=
B. Formula Mass
a. Ionic compounds consists of anions and cations. Since ionic compounds do not contain molecules, we use
Formula mass to describe the relative masses of ionic compounds.
b. Example
Given: Relative atomic masses of Na = 23.0, K = 39.0, Ca = 40.0, Cl = 35.5, C = 12.0, O = 16.0,
H = 1.0, Cu = 63.5, N = 14.0
1. Formula mass of potassium chloride (KCl)
= 39.0 + 35.5
= 74.5
2. Formula mass of sodium carbonate (Na2CO3)
= 23.0  2 + 12.0 + 16.0  3
= 106.0
3. Formula mass of sodium hydroxide (NaOH)
=
4. Formula mass of calcium oxide (
=
5. Formula mass of calcium hydroxide (
=
)
)
Part II The Microscopic World I/P.31
6. Formula mass of copper(II) nitrate (
=
7. Formula mass of ammonium carbonate (
=
)
)
Name of covalent compounds
1. B, Si, C, P, N, H, S, I, Br, Cl, O, F
2. The name of the second element should end up with -ide
3. A prefix (mono, di, tri, tetra)
4. Example: CIF3
Chlorine
trifluoride
N2O4
Dinitrogen
CH4
NH3
H2O
H2O2
tetraoxide= tetroxide
Methane
Ammonia
Water
Hydrogen peroxide
Part II The Microscopic World I/P.32
VIII. Structure, Bonding and Properties
A. Structures of Substances
Structure= what its constituent particles are, and how they are arranged and packed together
All substances exist as either giant structures or molecular structures.
1. Giant structures
a. Giant structures include:
•
giant ionic structures 巨型離子結構 e.g. sodium chloride
•
giant covalent structures 巨型共價結構 e.g. diamond, quartz
•
giant metallic structures 巨型金屬結構 e.g. copper
b. In a giant structure, millions of particles (atoms or ions) are joined together by strong chemical bonds. A
huge network is formed and the structure is difficult to break. A continuous giant lattice forms, particles
are packed in regular pattern and no discrete molecules exist. Almost all substances with giant structures
are solids under room conditions.
2. Molecular structures
a. Molecular structures include:
•
simple molecular structures e.g. hydrogen, chlorine, carbon dioxide, water
•
macromolecules 巨大分子 e.g. polyethene (plastic)
b. Simple molecular structures consist of separate molecules. The atoms within the molecules are strongly
bonded together by covalent bonds. The intermolecular forces between the molecules are weak.
c. Macromolecules are very big molecules containing thousands of atoms joined together by covalent bonds
and usually solids.
B. Structure and Properties
1. Simple Molecular Structures
Most non-metal elements (e.g. hydrogen H2, oxygen O2, chlorine Cl2, iodine I2 etc.) and covalent compounds
(carbon dioxide CO2, water H2O, methane CH4, ammonia NH3) are composed of simple molecules.
In simple molecular substances, the atoms are joined together within the molecule by strong covalent bonds,
but the separate molecules are attracted to each other by much weaker intermolecular forces (e.g. van der
Waals' forces).
a. Carbon dioxide
Dry ice consists of separate carbon dioxide molecules.
In each molecule, strong covalent bonds hold the carbon and oxygen atoms together. The carbon dioxide
molecules are packed closely to one another in a regular pattern. Weak van der Waals’ forces hold the
molecules together.
b. Iodine
Part II The Microscopic World I/P.33
The iodine molecules are packed closely together in a regular pattern. Weak van der Waals’ forces hold the
molecules together. The pattern is repeated millions of times, and the result is a crystal.
Properties of Simple Molecular Substances
(i) Melting points and boiling points
Simple molecular substances have low melting points and boiling points. This is because the molecules
are held together by weak intermolecular forces and can be separate easily by little heat energy.
e.g. melting point of oxygen is -218oC and the boiling point is -183oC.
(ii) Solubility in water
Intermolecular forces are weak. It is easy to separate the molecules and break down the crystal structure.
Simple molecular substances are usually slightly soluble or insoluble in water but very soluble in nonaqueous solvents.
Let us take iodine as an example. Iodine is slightly soluble in water but very soluble in non-aqueous
solvents. Water molecules have strong intermolecular forces. The weak attractive forces between
iodine and water molecules are not strong enough to separate the water molecules. Thus, iodine does not
dissolve readily in water.
In non-aqueous solvents, the molecules are held together by weak attractive forces. Hence iodine and nonaqueous solvent molecules can mix together easily.
Part II The Microscopic World I/P.34
(iii) Electrical conductivity
Simple molecular substances do not conduct electricity no matter in solid state or liquid state because they
do not contain mobile ions or delocalized electrons to conduct electricity.
Note: Aqueous solutions of some substances with simple molecular structures conduct electricity.
This is because mobile ions are formed when they dissolve in water. Examples include sulphur
dioxide, hydrogen chloride, ammonia, etc.
(iv) Hardness
Solid simple molecular substances are usually soft because the forces of attraction between the molecules
are weak (i.e. weak intermolecular forces i.e. van der Waals’ forces).
2. Giant Covalent Structures
These are substances made from millions of atoms joined by strong covalent bonds to form a giant network.
Common examples are quartz, diamond and graphite.
Diamond and graphite are different forms of the same element carbon, i.e. carbon is said to exhibit allotropy,
diamond and graphite are the allotropes 同素異構體 of carbon.
a. Diamond
(i) Diamond is a form of carbon.
(ii) In diamond, each carbon atom is surrounded by 4 other carbon atoms in the form of a three-dimensional
structure. To break the structure, must have extreme hardness and very high melting point. Diamond cannot
conduct electricity because it does not contain delocalized electrons.
b. Graphite
(i) Graphite is another form of carbon. It has a giant covalent network, but even so its properties are very
different. The carbon atoms are arranged in flat, parallel layers. Each layer contains many six-membered
carbon rings. Graphite is soft and slippery, and it conducts electricity.
(ii) Structure of Graphite
In graphite, each layer contains millions of carbon atoms. Within each layer, every carbon atom is joined to
three others by strong covalent bonds, one outer electron of each carbon atom is delocalized. They are free to
move from con carbon ring to the next within a layer. Thus graphite can conduct electricity. The carbon atoms
are difficult to separate from one another, so graphite, like diamond, has a high melting point.
Part II The Microscopic World I/P.35
(iii) Physical Properties of Graphite
1. Graphite is a good electrical conductor
Each carbon atom in graphite forms covalent bonds with three other carbon atoms. Since each carbon atom
has 4 outermost shell electrons, 1 electron is "free". The free electrons of each carbon atoms can move
between the layers. Graphite conducts electricity because of these free electrons.
2. Graphite is soft and slippery (other are hard and non-conductor)
The forces of attraction between layers are the weak van der Waals' forces. They are able to slide easily
over one another, rather like a pack of cards. This makes graphite soft and slippery. When you write with a
pencil, layers of graphite flake off and stick to the paper.
c. Quartz
(i) Quartz is crystalline form of silicon dioxide, SiO2.
(ii) Structure of Quartz
In quartz, every silicon atom is joined to four oxygen atoms by strong covalent bonds. Every oxygen atom is
also joined to two silicon atoms. This arrangement goes on continuously. The structure consists of a network
of covalent bonds.
(iii) Physical Properties of Substances with Giant Covalent Structures
1. Hardness
Substances with giant covalent structures consists of a network of covalent bonds. This makes them very
hard (except graphite). For example, instruments for cutting glass contain diamond.
2. Melting points and boiling points
Substances with giant covalent structures have high melting points and boiling points. It is because the
forces of attraction between the atoms are strong (the covalent bonds). Much energy is needed to overcome
the large number of bonds with great attractive forces.
3. Solubility in water
They are insoluble in water. It is because the atoms are held together strongly and it is very difficult to
separate the atoms.
4. Electrical conductivity
Since the bonding electrons in substances with giant covalent structure cannot move (i.e. no free ions or
free electrons), they do not conduct electricity in solid state or molten state (except graphite).
Part II The Microscopic World I/P.36
d.
Some properties of quartz, diamond and graphite
Substance State at room temperature and Melting Point
pressure
(oC)
Quartz
Solid
1610
Diamond
Solid
3500
Graphite
Solid
3730
Solubility in
water
Insoluble
Insoluble
Insoluble
Electrical
conductivity
Non-conductor
Non-conductor
Conductor
3. Giant Ionic Structures
a. Ionic compounds are made from the regular packing of positively and negatively charged ions. Because
of this, ionic compounds are described as three-dimensional giant ionic structures.
Structure of sodium chloride
NaCl
Positively charged (Na+) and
negatively charged (Cl-) held tgt
by ionic bonds
Each Na+ is surrounded
by 6 Cl-
Each Cl- is surrounded
by 6 Na+
Structure of caesium chloride CsCl
Since caesium ion is larger in size than the sodium ion, each Cs+ ion is surrounded by 8 Cl- ions and
each Cl- ion is in turn surrounded by 8 Cs+ ions.
b.
We can also show the structure of sodium chloride by using a ball-and-stick model.
c. Properties of Ionic Compounds
(i) Melting points and boiling points
➢
Ionic compounds have high melting points and boiling points.
Part II The Microscopic World I/P.37
It is because the ions are packed in a giant lattice and the force of attraction between the ions (i.e.
ionic bond) is strong. Much heat energy is needed to overcome the attraction.
➢
For example, the melting point of sodium chloride is 808oC and the boiling point is 1465oC;
the melting point of magnesium oxide is 2852 oC and the boiling point is 3600 oC.
(ii) They are crystalline in solid state
➢
The crystals of ionic compounds have flat sides and regular shape.
(iii) Solubility in water
➢
Many ionic compounds are soluble in water but insoluble in non-aqueous solvents.
➢
When sodium chloride is put into water, there is attraction between the ions of sodium chloride and
the water molecules. This cause the sodium and chloride ions to be removed from the lattice. The
ions then move into the solution.
➢
No such attraction exists between the ions of sodium chloride and the non-aqueous solvents
molecules, like tetrachloromethane.
Some ionic compounds such as calcium sulphate are insoluble in water. This is because the attractive
forces between ions in the solid are stronger than those between the ions and the water molecules.
(iii) Electrical conductivity
➢
Ionic compounds conduct electricity when molten or in aqueous solution. They do not conduct
electricity when in solid state because no mobile ions.
➢
It is because the ions become mobile when molten or in aqueous solution.
(iv) Hardness
Ionic compounds like sodium chloride are hard. This is due to the strong ionic bonds between oppositely
charged ions. Ionic compounds are hard but brittle.
4. Giant Metallic Structure
In a metal, atoms are packed tightly together in a regular pattern to form a giant structure. The metal ions are
surrounded by a ‘sea’ of delocalized electrons and it is called giant metallic structure.
a. Properties of metals
(i) Density
Metals generally have high densities due to the close packing of their atoms.
Part II The Microscopic World I/P.38
(ii) Melting point
The atoms in metals are packed closely and the metallic bonds holding them together are very strong. To
melt a piece of metal, a lot of heat energy is needed to overcome the strong attraction forces.
(iii) Electrical and heat conductivities
Metals are good conductors of electricity and conductors of heat due to the movement of mobile electrons
in a metal. The delocalized electrons there get more energy. They move faster and collide with the
neighbouring electrons. Heat is transferred in the collisions. The whole piece of metal becomes hot.
(iv) Malleability 展性 and ductility 延性
Metals are both malleable and ductile. The ions are packed in layers. When apply force to a piece of metal,
the layers of ions can slide over one another. Then the ions can settle into new shape. The metal piece does
not break easily because the non-directional metallic bond to hold the metal ions together.
b. The strength of metallic bond of sodium, magnesium and aluminium
(i)
Metal
Sodium
Magnesium
Aluminium
Melting point
98oC
650 oC
660 oC
(ii) Sodium, magnesium and aluminium are metals. The strength of the metallic bond depends on the number of
delocalized electrons in the metal structure.
Sodium has one outermost shell electron, magnesium has two, while aluminium has three. The strength of
metallic bond and hence the melting point increases from sodium to aluminium.
C. Predicting the properties of substances
Part II The Microscopic World I/P.39
Example
1. A compound Z is formed from the reaction between two elements X and Y. The electronic arrangements of
atoms of the two elements are given below.
Element
Electronic arrangement
X
2,8,8,1
Y
2,8,7
a. Predict the type of bonding present in Z.
b. What type of structure does Z have?
c. Predict the following properties of Z:
(i)
melting point and boiling point; and
(ii) electrical conductivity
Answer
a. Ionic bond
b. Giant ionic structure
c. (i) high melting point and boiling point
(ii) conduct electricity only when in molten state or liquid state.
2.
Structure A
Structure B
Structure C
a. Name three substances which have the same structures as A, B and C respectively.
b. Name the type of bonding between particles in
(1) B,
(2) C.
c. In terms of the forces between particles, explain why there is a large difference in melting point between
Part II The Microscopic World I/P.40
(1) A and B,
(2) A and C.
In each case, state which of the two solids is expected to have the higher melting point.
Answer
a. Structure A: Iodine
Structure B: Sodium chloride
Structure C: Diamond
b. (1) ionic bond
(2) covalent bond
c. (1) B has a higher melting point.
The attraction between the ions in structure B is ionic bond. The attraction between the
molecules in structure A is van der Waals’ forces. The ionic bond is much stronger than
the van der Waals’ forces. Therefore, the melting point of B is much higher than that of
A.
(2) C has a higher melting point.
The attraction between the atoms in structure C is covalent bond. The attraction between
the molecules in structure A is van der Waals’ forces. The covalent bond is much
stronger than the van der Waals’ forces. Therefore, the melting point of C is much
higher than that of A.
3. The table below shows some properties of four substances.
Substance
Melting point (oC) Boiling point (oC)
A
776
1500
Electrical conductivity
Solid state
Molten state
Nil
Good
B
C
D
961
3500
-7
2160
4827
59
Good
Nil
-
Good
Nil
-
Explain which substance is likely to have:
a. a simple molecular structure;
b. a giant metallic structure;
c. a giant ionic structure;
d. a giant covalent structure
Answer
a. D. Because D has a low melting point and boiling point.
b. B. Because B has a high melting point and boiling point. B also can conduct electricity in
solid state.
c. A. Because A has a high melting point and boiling point. A only can conduct electricity in
molten state, but does not conduct electricity in solid state.
d. C. Because C has a very high melting point and boiling point. However, C does not conduct
electricity in both solid state and molten state.
Part II The Microscopic World I/P.41
4. HKCEE 1998 Q7a
Both carbon and silicon are Group IV elements in the Periodic Table. The diagram below show the structures
of dry ice (solid carbon dioxide) and quartz (a form of silicon dioxide):
(i) With reference to the structures of the two substances, explain why quartz is a solid which melts at a high
temperature, while carbon dioxide is a gas at room temperature.
(ii) With the help of a labeled diagram, suggest how to show experimentally that dry ice sublimes to give
gaseous carbon dioxide.
(iii) Sand (an impure form of quartz) and limestone are raw materials used for making glass.
(1) Name the main chemical constituent of limestone.
(2) Suggest ONE reason why glass had been used by mankind for a long time.
(3) Suggest ONE reason why glass bottles are preferred to plastic bottles for the storage of champagne.
(9 marks)
Part II The Microscopic World I/P.42
Answer
Part II The Microscopic World I/P.43
Structure
Giant Metallic
Structure
Giant ionic
Structure
Giant covalent
Structure
Bonding
Strong metallic bond
between delocalised
electrons and metal
ions
Strong ionic bond
holding the ions
Strong covalent
bond holding
atoms
High m.p & b.p
It has giant metallic
structure with strong
metallic bond
between delocalized
electrons and mobile
ions
Electrical
Conduct electricity
conductivity There are presence of
delocalised electrons
and mobile ions
High m.p & b.p
It has giant ionic
structure with
strong ionic bond
between cations
and anions
High m.p & b.p
It has giant
covalent
structure with
strong covalent
bond between the
atoms
Graphite only
There are
presence of
delocalised
electrons
Diamond X
SiO2 X
Insoluble in
water
m.p/ b.p
Solubility
Most are insoluble in
water
Conduct electricity
only in molten or
aqueous state
The ions become
mobile when is
molten or in
aqueous solution
Most are soluble in
water
Simple
Molecular
Structure
Within the
molecules:
covalent bond
Between the
molecules: weak
van der waals’
force
Low
It has weak van
der waals’ force
between
molecules
Do not conduct
electricity
The absent of
delocalised
electrons and
mobile ions
Insoluble in
water and nonaqueous solution