Secondary 2 Chemistry Notes

For Secondary 2, Chemistry is grouped under 4 major topics, as follows,

Elements, Atoms and Molecules

Atomic Structure

Writing Formula and Equations

Acids and Bases

The Periodic Table of Elements

1.1 Elements, Atoms and Molecules

An element is a substance that cannot be broken down into simpler substances

by chemical methods. To date, there are 118 known elements, as seen from the

Periodic Table of Elements. 92 of them are found in Earth, in the rocks, soil, air

and water. Others are made artificially by scientists in laboratories, and are

known as Synthetic Elements. One example is Technetium, with an atomic

number of 43. They cannot be found on Earth naturally as they have short half-

lives, which mean that the element will decay away quickly.

Every element is represented by a chemical symbol. The symbol consists of a

capitalized letter, but for cases with symbols having 2 letters, only the first letter

is capitalized.

Capitalized

Not Capitalized

Capitalized

Sources: Google Image

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Classification of elements

Classifying by physical states:

- At room temperature, elements have varied physical state, thus can be

classified in this manner.

- Among 92 naturally occurring elements,

11 are gases (Hydrogen, Helium)

2 are liquids (Bromine, Mercury)

79 are solids (Lithium, Carbon)

Classifying as metals and non-metals:

Elements are classified as metals and non-metals, based on the element’s

electrical conductivity.

Metals are good conductors of electricity while non-metals are very poor

conductors of electricity.

However, there are some elements are classified as metalloids, such as Silicon.

They are semi-conductors of electricity, having properties between those of

metals and non-metals, useful for the computer industry.

1.2 Atoms

Elements are made up of tiny particles called atoms. Atoms differ from one

element to another.

It is the smallest particle in any element, having a diameter of only 0.1

nanometre. It cannot be seen with the naked eye, only to be seen with the

electron microscope, or the scanning tunnelling microscope (STM). Every element

consists of only a specific type of atom. (So Sodium, Na, can only contain Sodium

atoms, not any others)

Not all elements exist as atoms. However, some are. They are under Group 0 in

the Periodic Table of elements, founded by William Randsey. These elements are

also called the noble gases, since they are extremely unreactive with other

elements. Thus, they were once known as inert gases, but this mistake was

corrected when it was realized that the noble gases gets more reactive as it goes

down the group. One example would be Xenon reacting with Fluorine to form

Xenon Tetrafluoride (XeF4)

It is possible for the atoms to exist as atoms as they have a stable octet electron

configuration, and a stable duplet electron configuration in the case of Helium.

(This would be explained further under the chapter of “Atomic Structure”)

Metals are also able to exist as atoms. This is due to metals arrangement to form

giant structures. Most non-metals exist as molecules.

1.3 Molecule

A molecule is a group of two or more atoms that are covalently bonded together.

Molecules of elements consist of a fixed number of the same type of atoms

combined together.

For example, Fluorine molecules are formed from two chemically combined

atoms of Fluorine. Thus, its molecular formula is F2.

Ne (Neon) is an example of a monatomic (mono = 1) element, since it exists as

single atoms.

F2 (Fluorine molecule) is an example of a diatomic (di = 2) molecule, since it is

formed by the combination of 2 atoms

O3 (Ozone), is an example of a triatomic ( tri = 3) molecule, consisting of 3 atoms

chemically combined together.

Any molecules formed by 3 or more molecules, i.e S8 (Sulfur), is known as

polyatomic molecules (poly = many)

1.4 Compounds

Compounds are made up of molecules and ions.

Compounds are made up of 2 or more different elements chemically combined

together.

Molecules of compounds consist of a fixed number of two or more different types

of atoms combined together.

An example would be Ammonia, NH4.

Ions also make up compounds, known as ionic compounds.

One example would be table salt, Sodium Chloride, with the chemical formula of

NaCl. Ions are simply atoms with electric charges as they had gained or lost 1 or

more electrons.

Ions with a positive charge (Lost electron) are known as cations.

Ions with a negative charge (Gained electron) are known as anions.

2.1 Atomic Structure

Relationship between

the two elements, 1 N

atom for every 4 H

atoms.

Symbol for element

The structure of the atoms is quite simple, although the processes going on are

quite complicated. Let’s now look at the sub-atomic particles inside the atom.

As you can see, the atom is made up of 3 fundamental sub-atomic particles, the

Proton, Neutron, and the Electron. The neutrons and the protons are grouped

together in the nucleus, and together called the nucleons. The nucleus is very

dense, since it accounts for all the mass inside the atom, although it takes up little

spaces, but surrounded by electrons. These electrons are attracted to the

protons electrostatically.

Proton Neutron Electron Relative charges +1 0 -1

Relative masses 1 1 1/1836 Location in the atom

Nucleus Nucleus Electron shell

Since every atom is electrically neutral, when an electron, having a negative

charge, is removed, the atom becomes positively-charged, and it becomes

negatively charged when an electron is gained. Electrons are also responsible for

the chemical properties on an atom.

Source: Google Images

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2.2 Expression of the Atom

To represent an atom, we could use the following method of

AZX

Where X is the symbol of the Element, while A is the nucleon/mass number

(Number of protons + neutrons) and Z is the atomic (proton) number.

Since the amount of electrons is equivalent to the amount of protons in an

electrically neutral atom, we can derive the amount of electrons from the atomic

number.

We can also derive the number of neutrons by subtracting the atomic number

from the nucleon number.

2.3 Isotopes

Isotopes are atoms of the same element, but have different number of neutrons.

Isotopes of an element have the same

Proton Number Number of electrons in an atom

Electron configuration

Chemical properties

Isotopes on an element have different

Nucleon Numbers Number of neutrons in an atom

Physical properties, i.e. Boiling/Melting points

One example would be Carbon. It could have the mass number of 12, 13 or even

14.

2.4 Shells

Electrons move around within the atom along their shells. These shells can hold

up to a maximum number of 2n2 electrons, where n is the number of the shell.

The shells are numbered 1, 2, 3, 4, 5...Etc. The 1st is nearest to the nucleus,

followed by the 2nd shell, and then the 3rd shell.

Now, look at the electronic configuration of the first 20 elements in the Periodic

Table of Elements.

Element Electronic Configuration Hydrogen 1

Helium 2 Lithium 2.1

Beryllium 2.2

Boron 2.3 Carbon 2.4

Nitrogen 2.5 Oxygen 2.6

Fluorine 2.7

Neon 2.8 Sodium 2.8.1

Magnesium 2.8.2 Aluminium 2.8.3

Silicon 2.8.4 Phosphorus 2.8.5

Sulfur 2.8.6

Chlorine 2.8.7 Argon 2.8.8

Potassium 2.8.8.1 Calcium 2.8.8.2

Notice that the first shell can only hold up to 2 electrons, while the second shell

can hold up to eight. The third shell, for the first 20 elements, can only hold up to

8 electrons. This is due to the energy levels of the shells that will not be tested in

Secondary 2.

Look at the noble gases, highlighted in red. They all have a stable duplet structure

for Helium; since it is the valence shell (last shell) had been fully occupied by the

electrons and is stable. On the other hand, for both Neon and Argon, they have a

stable octet structure. Thus, few can combine with them to form compounds or

even with each other to form molecules. This explains why they are monatomic.

2.5 Chemical bonds

Chemical bonding is the attraction between atoms or molecules which allows

the formation of chemical compounds, which contain two or more atoms. These

atoms bond to achieve a stable electronic configuration.

Ionic Bonding

Before we talk about Ionic Bonding let’s have an introduction to ions.

Ions are atoms which had either lost or gained electrons. This makes the atoms

not electrically neutral and is given the term, ion.

Ions are separated into two different categories, Cations and Anions. Cations are

ions with a positive charge (lost electrons) while Anions are ions with a negative

charge (gained electrons).

Ionic bonding takes place due to most atoms not having the stable electronic

configuration of a noble gas. Thus, the atoms would have to either lose or gain

electrons to achieve that configuration.

To remember which ions are cations and which are anions, there is a simple

method. Metals will form cations while non-metals will form anions.

Let’s have some examples.

Atom of chlorine Ion of chlorine

This negative sign tells us that Chlorine only

gained one electron.

This positive sign tells us that Sodium only

lost one electron.

Atom of Sodium Ion of Sodium

Source: Google Images

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Source: http://www.bbc.co.uk/schools/gcsebitesize/sc

ience/add_aqa/atomic/ionicrev3.shtml

If you have ionic compounds such as Calcium Chloride, CaCl2, the Calcium atom

will transfer one electron to one of each Chlorine atom and then all of them will

achieve a stable octet structure.

Covalent Bonding

This type of bonding is only achieved between non-metals. Like ionic bonding, it is

to attain the electronic configuration of a noble gas.

The electrons of the Hydrogen atoms and those of Oxygen’s will share, giving all

of the atoms an electronic configuration of a noble gas.

3.1 Writing Formula and Equations

Some rules/ definitions:

1) For all formulas, the metal is always written before the non-metals. So in

the case of Sodium Chloride, it is always written as NaCl.

2) Cations are usually metals and anions are usually non-metals

3) Diatomic molecules’ names end in –ide, i.e. Germanium Telluride

Hydrogen Electrons

Oxygen Electrons

Source: Google Images

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4) If a compound contains oxygen, its name would end in –ate or –ite

5) A compound contains more oxygen atoms if its name end with –ate instead

of –ite.

3.2 Chemical formula of ionic compounds.

Writing chemical formulas are an important part in the learning of chemistry,

especially when you need to express the composition of a substance.

There is a simple method to write chemical formulas of ionic compounds. Let’s

use an example, Calcium Chloride.

Firstly, we have to determine the amount of electrons each of their ions gain or

lose.

Calcium gains two electrons, while chlorine loses one.

Written in symbols:

Ca2+ Cl-

Next, do a switch.

Ca2+ Cl-

Ca Cl2

The chemical formula for Calcium Chloride is thus CaCl2.

Generally, when Ax+ and By- forms an ionic compound, the resulting chemical

formula will be AyBx.

3.3 Molecules

This chapter had been explained earlier on. However, what’s new is the

nomenclature (naming of compounds).

Assuming we are asked for the chemical composition of Sulfur Trioxide, how do

we find it out?

Look at the table below:

Amount of element in a molecule Prefix

1 mono

2 di 3 tri

4 tetra 5 penta

6 hexa 7 hepta

Sulfur Trioxide

Thus, we can conclude that the chemical composition is SO3.

3.4 Balanced Chemical Equations

We used to write word equations in Secondary 1. However, we now have to

express the equations as symbols.

Let’s say we are asked to write a chemical equation of

Magnesium + Oxygen Magnesium Oxide

Prefix of 3

Step 1: Write down the reactant (Left hand side) and Product (Right hand side).

Mg + O2 MgO

Mg atom: 1 Mg atom: 1

O atoms: 2 O atoms : 2

Step 2: Balancing the equations

Mg + O2 2MgO

Mg atom: 1 Mg atoms: 2

O atoms: 2 O atoms: 2

Step 3: Recheck

We realize that the Mg atoms are not balanced. So, we have to add one more Mg

atom to the side of the reactant.

Therefore,

2Mg + O2 2MgO

Mg atoms: 2 Mg atoms: 2

O atoms: 2 O atoms: 2

The equation is now balanced.

Last step: State symbols (If required)

Sometimes, we will be asked to show the physical state of a substance.

For solids, we use (s)

For liquids, we use (l)

For gas, we use (g)

For aqueous solutions (solutions/ substances dissolved in water), we use (aq)

4.1 Acids and Bases

An acid is a substance which dissociates in water to produce hydrogen ions, H+.

So, acids only show their properties when dissolved in water.

Due to that property, it is known as a proton donor, since H+ ions contains no

electrons.

Also, the hydrogen ion is responsible for all acidic properties.

Organic Acids: Naturally occurring acids i.e. citric acid, ethanoic acid

Inorganic Acids (mineral acids): Mostly man-made acids i.e. hydrochloric acid

4.2 Strength of Acid

Strength (not to be confused with concentration) of an acid refers to how easily

an acid dissociates into its ions when dissolved in water.

HCl (aq) H+ (aq) + Cl-

Hydrochloric acid, being very strong, completely dissociates in water to give

hydrogen ions and chloride ions.

Acids

Mineral/ Inorganic acids Organic acids

CH3COOH (l) CH3COO- (aq) + H+ (aq)

Ethanoic acid, being a weak acid, dissociates in water partially to give hydrogen

ions and ethanoate ions.

4.3 Properties of Acids

Acids

- Tastes sour

- Turn blue litmus paper red

- pH values below 7

- Dissolves in water to form solutions to conduct electricity (due to the

mobile ions) in aqueous form.

4.4 General equations of acids:

Metal + Acid Salt + Hydrogen

The next page will feature the reactivity series.

Metals placed above Hydrogen in the series will react with acids to form

hydrogen (fulfilling the above equation).

Metals placed below Hydrogen in the series will not react with acids to form

hydrogen.

Example: Mg(s) + 2HCl (aq) = MgCl2(s) + 2H(g)

Another equation:

Carbonate + Acid CO2 + H2O + Salt

Example: (Chemical Equation)

MgCO3 + H2SO4 CO2 + H2O + MgSO4

Metal Acid Salt Hydrogen

Source: Google Images

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Example: (Word Equation)

Magnesium Carbonate + Sulfuric Acid Carbon Dioxide + Water + Magnesium

Sulfate

Metal oxide + Acid Salt + Water

Metal hydroxide + Acid Salt + Water

They are neutralisation reactions.

Example:

2NaOH + H2SO4 NaSO4 + 2H2O

Sodium Hydroxide + Sulfuric Acid Sodium Sulfate + Water

4.5 Bases and Alkalis

A base is a substance that reacts with acids to produce salts and water

There are ionic bases and covalent bases.

Ionic bases consists of metal hydroxides and metal oxides

There are many covalent bases, but we need to know about NH3 (aq), known

aqueous ammonia.

Alkalis are bases soluble in water

Like acids, alkalis dissociates in water, but to form hydroxide ions (OH-) instead

of hydrogen ions.

Also, they only show their alkaline properties after being dissolved in water.

Also, hydroxide ions are responsible for all alkaline properties.

Example:

KOH K+ + OH-

Potassium Hydroxide Potassium ion + Hydroxide ion

4.6 Strength of Alkalis

The strength of alkalis refers to how easily an alkali dissociates into its ions when

dissolved in water.

LiOH (s) Li+ (s) + OH- (aq)

Lithium Hydroxide, being very strong, completely dissociates in water to give

lithium ions and hydroxide ions.

NH3(aq) OH- (aq) + NH4+ (aq)

Aqueous ammonia, being a weak alkali, dissociates in water partially to give

hydroxide ions and ammonium ions.

4.7 Properties of Alkalis

Alkalis

- have a soapy feel

- tastes bitter

- turn red litmus paper blue

- have >pH7

- Dissolves in water to form solutions to conduct electricity (due to the

mobile ions) in aqueous form.

4.8 General equations of alkalis

Alkali + Acid Salt + H2O

The reaction of alkali and acid is known as neutralisation.

Example:

KOH + HCl KCl + H20

Potassium Hydroxide Potassium Chloride + Water

Alkali + Ammonium salt Salt + NH3 + H2O

Example:

LiOH + NH4I LiI+NH3 + H2O

Lithium Hydroxide + Ammonium Iodide Lithium Iodide + Ammonia + Water

Alkali (A) + Salt (B) B Hydroxide + Salt (A) Where A and B are metals.

Example:

2NaOH +NiCl2 Ni(OH)2 + 2NaCl

However, the metal hydroxide produced mustn’t be soluble in water.

Enough about the equations, if we can’t apply them in real life, the knowledge

wouldn’t help us much.

The following table will tell you more about the bases/ alkalis and their uses.

(Continued in the next page)

Bases and Alkalis Uses

Sodium hydroxide Making soaps and detergents Involved in the production of paper

Magnesium hydroxide Used in toothpastes Used in antacids to aid gastric problems

and indigestion Calcium oxide Make iron, concrete and cement

Ammonia solution Used in fertilisers

Used in solutions for cleaning windows Potassium hydroxide Used to neutralise acidic soil

4.9 pH

Source: Google Image

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Above you see the pH scale. It compares the strength of acids and how acidity or

alkaline a solution is.

As you can see, pH of 7 is neutral. They are neither acidic nor are they alkaline i.e.

Deionised water

pH < 7 are acidic i.e. Hydrochloric acid

pH > 7 are alkaline i.e. Sodium hydroxide

Also, the lowest value is 1 and the highest is 14

The pH scale measures a solutions based on the number of hydrogen or hydroxide

ions it has.

The more hydrogen ions a solution have, the more concentrated the acid and

the more hydroxide ions a solution have, the more concentrated the alkaline.

When there is a neutralisation reaction, the hydroxide ions (OH-) and hydrogen

ions (H+) will react.

H+ + OH- H2O

This is why water is one of the products of a neutralisation reaction.

4.10 Acid- Base Indicators

It was learnt in primary school that CO2 reacts with limewater (Ca(OH)2)

Limewater is an alkaline and acts as an indicator for CO2.

So what are indicators?

Indicators are substances that changes colour when added to an alkaline or acidic solution.

There is an indicator known as “Universal Indicator”. It contains a mixture of

several different dyes and can determine the pH of a solution. We will be able to

recognise its pH from the colour the dye turns into.

The colours can be seen from the picture on the next page

The colour changes when the acidity/ basicity changes.

Two other very common indicators are methyl orange and phenolphthalein.

- Phenolphthalein is colorless in acidic solutions and pink in basic

solutions.

- Methyl orange is red in acidic solutions and yellow in basic solutions. At

the end point of the titration, methyl orange appears orange.

4. 11 Oxides

This is the last section to be studied at Secondary 2 level (although this may vary

from school to school)

Oxides can be branched out into 4 different types shown below

Oxides

Acidic Oxides Basic Oxides Amphoteric Oxides Neutral Oxides

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Acidic oxides are oxides of non-metals that dissolve in water to form an acid

However, they also react with bases to form salt and water (like acids)

Basic oxides are oxides of metals that react with acids the same way bases do.

(Form salt and water).

Oxides of reactive metals will dissolve in water to form alkalis.

Basic oxides, as the name suggests, are insoluble in water (bases). However, some

are alkaline, dissolving in water readily e.g. NaOH

Note: Basic oxides are solids at room temperature.

Amphoteric oxides are metallic oxides which react in both acids and alkalis to

form salts and water, displaying properties of both acidic and basic oxides.

Neutral oxides are usually non-metals, showing neither acidic nor basic

properties. They are insoluble in water.

Some reactions of oxides

ZnO + H2SO4 ZnSO4 + H2O (Reaction of an amphoteric oxide)

MgO + H2SO4 MgSO4 + H2O (Reaction of a basic oxide)

K2O + H2O K(OH)2 (Reaction of a basic oxide)

SO3 + H2O H2SO4 (Reaction of an acidic oxide)

5.1 The Periodic Table

The Periodic Table of Elements that we are using today is the brainchild of a

Russian chemist Dmitri Mendeleev. At his time, in the 19th Century, more and

more elements are discovered and it became almost impossible for anyone to

remember the elements by heart. Organisation of the elements was needed by

chemists all around the world. Dmitri, while preparing a chemistry textbook, saw

patterns among the elements and created the Periodic Table of Elements.

Chemists soon accepted it and over the years, some minor changes were made to

the table we know today.

We have to know a few groups of elements, as followed

- Alkaline metals

- Alkaline earth metals

- Metalloids

- Halogens

- Noble Gases

- Transition elements

5.2 Alkaline Metals

Alkaline metals are found in Group I of the Periodic Table, excluding the element

hydrogen. As their names suggests, alkaline metals form compounds with oxygen

and give alkaline solutions when dissolved in water. These metals are also highly

reactive with water.

5.3 Alkaline earth metals

This group of metals are found in Group II. They form compounds with oxygen

and give alkaline solutions after dissolving in water, however, the charge of their

ions will be different from the alkaline metals. This will be elaborated further

afterwards.

5.4 Metalloids

Metalloids have properties of non-metals and metals. They are semi-conductors

and so, they do not conduct electricity as well as metals do, but better than non-

metals. Some examples are silicon and germanium. They are modified to become

transistors, allowing the use of electronic devices and are responsible for solid-

state electronics’ rapid growth.

5.5 Halogens

Halogens are found in group VII of the Periodic Table. The name is derived from

the two Greek words, halas and gennau, which means salt and generate

respectively. Fluorine from this group is highly reactive, being able to form bonds

with noble gases. However, it is less reactive as you go down the group.

5.6 Noble Gases

Noble gases were once thought to be inert gas, which means that they show no

sign of any chemical reactivity. However, it was found that some noble gases like

Xenon have some chemical reactivity with highly reactive elements like Fluorine.

This term, “Noble” doesn’t restrict to only gases. Gold is commonly referred to as

a “Noble metal” due to its limited reactivity with other elements.

5.7 Transition metals

Transition metals are just the block of elements found in the middle of the

Periodic Table, with no group numbers. They are metals and thus have all the

qualities that a metal should possess i.e. ductility and malleability.

5.8 Ion charges

Aside for the ability for us to recognise the properties of the elements, the

Periodic Table allows us to determine the charges of an element so that we can

know the charges of a polyatomic compound and help us write chemical

equations.

The alkaline metals have a +1 charge when it becomes an ion, a.k.a. it loses one

electron. The halogens have a -1 charge when it becomes an ion a.k.a it gains an

electron. This is why when alkaline metals and halogens form an ionic bond; it is

always in the ratio of 1:1.

The alkaline earth metals would then have a +2 charge when it becomes an ion

and so on and so forth.

Group III elements would have a +3 charge when it becomes an ion

Group IV elements will have a -4 charge when it becomes an ion

Group V elements will have a -3 charge when it becomes an ion

Group IV elements will have a -2 charge when it becomes an ion

A few examples:

NaCl

SiO2

H2O2

You would realise that it’s either the positive charges and negative charges cancel

each other out or that the amount of electrons will be 8, allowing the compound

to obtain a stable duplet (for the case of hydrogen molecules, and monatomic

helium) or a stable octet configuration.

For transition elements, there is no pattern for us to recognise the charges. Below

will list some common charges transition metals when they are ions and some

compounds associated with transition elements.

Iron (II), Iron (III) Fe2+, Fe3

+ Zinc Zn2

+

Copper Cu+, Cu2+

Chromium (II), Chromium (III) Cr2+, Cr3

+ Silver Ag+

Barium Barium2+ Potassium Dichromate K2Cr2O7

Permanganate ion MnO4-

Some transition elements will have a Roman numeral behind its name, like

Chromium (II). This shows that Chromium has a charge of +3.

Credits

O-Level Chemistry Guide by Bob Ryan

Chemistry Matter and its Changes, Brady, Senese

Secondary 2 Chemistry Notes