chemistry module 2 enabling course · this will be discussed in section 2.2 of this module. 1.3...

Chemistry – Module 2

Enabling Course

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May 2011

©2011 Endeavour College of Natural Health MODULE 2 Enabling course: Chemistry

Module 2 The building blocks

1 Elements

1.1 Background

Our knowledge of the quest for what the world was made of can be traced to the various theories of

Ancient Greece. The alchemists built on these theories and traditional knowledge and through

experimentation discovered some substances that could not be decomposed, or broken down further,

by chemical means.

During the 17th century many more substances were tested to find out what they were made of. The

word element was used for the 12 substances already discovered that could not be decomposed

further: gold, iron, silver, sulfur, carbon, lead, mercury, tin, arsenic, bismuth, antimony and copper.

John Dalton redefined the word “element” to include his idea of the atom – the first modern atomic

theory (see Section 2.2).

In antiquity an “element” had to correspond to or be recognizable in what was found around you.

However, most of us will never see or hold many of the substances we call elements today. Another

76 elements were discovered between 1557 and 1925. We now know that 92 elements exist naturally.

Scientists working in laboratories have been able to make another 23 artificial elements. There is

nothing obvious about what makes up an element! (Ball 2002, p. 24).

1.2 The periodic table

Source: http://www.teachnet.ie/macalvey/images/Mendeleev.gif

This image of Dmitri Mendeleyev can be found at the above Internet address (Teachnet 2010).

Mendeleyev observed in 1869 that the properties of the 63 then-known elements arranged in order of

their “weights” would repeat in a regular pattern. His arrangement became known as the periodic

table.

http://www.teachnet.ie/macalvey/images/Mendeleev.gif


There are now over 110 elements into which all other substances can be broken down by chemical

means. As increasing numbers of elements were discovered, scientists measured their weights and

how they reacted. They began to notice families of elements behaved similarly. Groups of three were

observed to be similar, for example lithium, sodium and potassium. The Russian chemist and

professor Dmitri Mendeleyev (1834–1907) observed that elements listed in order of their “weights”

could be arranged so that their properties repeated regularly (or periodically). In 1869 he announced

his Periodic Law (periodicity) and published a list of known elements set out in a table. He left gaps

where the Periodic Law did not seem to fit, predicting that new elements would be found to fill them.

(Newmark 2002, pp. 22–23). If Mendeleyev’s system was correct, the properties of any yet

undiscovered element could be predicted by the Periodic Law. The first “gap” to be filled on

Mendeleyev’s periodic table was by the discovery of the element gallium in 1875.

There are many different ways of presenting the modern periodic table. An example is on the inside

front cover of your textbook.

A periodic table that includes the names of the elements spelled out is “The WebElements™ printable

periodic table” (the WEBELEM2.pdf file). A periodic table is also included with these Modules.

The modern periodic table is like a cipher, or code, for many properties of the elements, far beyond

those known in the time of Mendeleyev. For instance, information of the structure of atoms can be

gathered from the periodic table. This will be discussed in Section 2.2 of this module.

1.3 Element symbols and types

Elements and their symbols

Each element on the periodic table is given a symbol which is shorthand for its name, just as your

initials are shorthand for your name.

The use of upper and lower case letters is specific. The first letter must be a capital letter. If there is a

second one for the element symbol, it must be in lower case. For example, Co is the element symbol

for cobalt, Mn for manganese and Mg for magnesium. However, the letters “CO” together do not mean

the element cobalt. CO is the symbol for the compound carbon monoxide, made of the elements

carbon (C) and oxygen (O) bonded together.

Some element symbols are not related to their names in English, but to their original names. In the

first 20 elements, the symbol for sodium is Na, which comes from the word natrium. The symbol for

potassium is K, from kalium.


Need to know check:

o The periodic table is an arrangement of elements in increasing “weights” that shows

regularly repeating patterns (periodicity). This table allows the properties of elements to be

predicted based on their position on the periodic table.

o You will need to know the correct spelling of the name as well as the symbol for each of

the first 20 elements plus other elements such as manganese, iron, cobalt, nickel, copper,

zinc, selenium, bromine, silver, tin, iodine, gold, mercury and lead. You will learn more

element symbols as they are mentioned in your studies.

Activities: Complete the following activities.

1. Add each of the bolded terms and its explanation to the Glossary you started in Module 1.

Check your definitions in your textbook, and memorise each term.

2. Print a copy of the periodic table.

3. RQ1:** Draw up a table with the headings: “Number”, “Symbol”, “Name” and “Type of

element”.

a. List the number, symbol and name of the first 20 elements. An example is shown

in the first row. Continue to fill in your table with the additional elements listed in

the “Need to know check”. Leave room to list additional elements.

Element

number

Symbol Name Type of element

1 H Hydrogen

Keep this table for later exercises.

**NOTE: Suggested answers to RQ review questions are listed in Appendix A.

Element types: metals, non-metals and metalloids

The elements on the periodic table are divided into two main types: metals and non-metals. Most

elements are metals, which are found on the left-hand side of the periodic table. Only 22 of the

elements are non-metals, and they are found on the right of the periodic table. Metals and non-metals

have different properties.


The main properties of metals and non-metals are summarized in the following table.

Main properties of metals and non-metals

Metal Non-metal

Solids at room

temperature except

for mercury, which is

a liquid

At room temperature, 11 are

gases, 10 are solid, one

(bromine) is a liquid

Can be polished to

produce a high shine

or lustre

The solids cannot be polished

– they are dull or glassy

Good conductors of

electricity and heat

Usually poor conductors of

electricity and heat

Can be beaten or

bent (malleable) into

a variety of shapes

They cannot be bent into

shape: they are brittle

Can be made into a

wire (ductile)

They are not ductile

Have high melting

points

Have lower melting and

boiling points: many are

gases at room temperature

There is also a third type of element called a metalloid. The metalloids have features of metals and

non-metals, and there are eight: boron, silicon, arsenic, germanium, antimony, polonium, astatine, and

tellurium. Sometimes astatine is classified as a non-metal.

Need to know check:

o The three types of elements (metals, non-metals and metalloids); where to find each type

of element on the periodic table; the properties of metals and non-metals; which of the first

20 elements are metals, non-metals and metalloids.


o Add each of the bolded terms and its explanation to the Glossary you started in Module 1.


o RQ2: Write a list of the main properties of metals. Describe how non-metals are different.

o RQ3: Complete the table in Exercise 4 by adding whether each element is a metal,

metalloid or non-metal in the “Type of element” column.


2 Atomic theory

2.1 Background

After Leucippus and his teacher Democrites in Ancient Greece, there was a break of about 2000 years

before matter was again thought of as consisting of “atoms”. A modern theory about atoms started to

be developed in the late 18th century. By then new gases, metals, and other substances had been

discovered. Many chemical reactions were studied and the weights of substances involved were

measured carefully. John Dalton, the “father” of modern chemistry, brought together the existing ideas

of the nature of matter and put forward a recognizably modern atomic theory, which he first stated in

1803.

Dalton’s basic ideas:

1. The elements are composed of indivisible identical particles called atoms.

2. The atoms of different elements differ in properties.

3. Atoms are the units of chemical change: chemical change merely involves the combination or

the rearrangement of atoms – atoms themselves are not created, destroyed or changed.

4. When atoms combine they do so in fixed ratios of whole numbers forming particles known as

“molecules”. The molecules in a compound are all identical.

Source: <royalsociety.org/Elements-and-Compounds/>

http://royalsociety.org/Elements-and-Compounds/


A page from a letter from John Dalton to the French physicist Gay-Lussac can be found at the above

Internet address. This shows some of the symbols that Dalton used for the elements, and

combinations of them, to represent compounds.

2.2 Modern atomic theory

Refinements to Dalton’s atomic theory followed understanding of the nature of electricity and its link

with the composition of atoms.

Today, atomic physicists continue the work in this field and they have identified more than 30

fundamental subatomic particles. However, in the energy range in which most chemical reactions

occur, only three particles are of interest.

Particles and their charges

The current model of the atom has three main subatomic particles – neutrons, protons and electrons.

These three particles have different charges. Just as we say a live electric wire carries a charge, or

an electric fence carries a charge, the proton and electron carry a charge. Protons are positively

charged, electrons are negatively charged and neutrons carry no charge.

Although a charge might be hard to define, we can easily see the effects of similar and different

charges. If you take two bar magnets you will notice that when two ends are brought close to one

another they will attract one another. If one of the magnets is turned around and the end is then

brought near to the first magnet, the two magnets repel one another. This is very similar to what

happens with electric charges.

Like charges repel and opposite charges attract. (Timberlake 2006).

A positive proton will attract a negative electron. But electrons will repel other electrons and protons

will repel other protons. The charges on the particles in an atom are shown in the following table.


Charges in an atom

Particle name Symbol Charge

Proton p+ One positive charge = 1+

Electron e– One negative charge = 1–

Neutron 0 No charge – it is neutral

Structure of an atom

Atoms are too small to see and so models are used to represent them. The atom is mostly empty

space. In the middle is a positively charged nucleus that contains protons and neutrons. Negatively

charged electrons move around the nucleus.

A model of the structure of an atom. In an atoms, the protons and neutrons are found in the nucleus;

the electrons are located outside the nucleus. (Timberlake 2006).

Protons and neutrons have approximately the same mass and electrons are much smaller, but move

very rapidly and take up most of the space occupied by the atom. Electrons have only a tiny mass

compared with either protons or neutrons. It takes about 1,850 electrons to equal the mass of one

proton or one neutron. The neutron and proton are contained in the nucleus at the centre of the atom.

Because protons are positively charged and repel each other, scientists think that the job of the

neutron is to help the protons in the nucleus to stay close together. The nucleus accounts for most of

the mass of the atom. The positively charged nucleus attracts the negatively charged electrons.

The three main particles in the atom

Particle Charge Mass (g) Relative mass (amu*)

Electron –1 9.11 × 10–28 1/1836

Proton +1 1.67 × 10–24 1

Neutron 0 1.67 × 10–24 1

* amu = atomic mass unit, a relative size of protons, neutrons and electrons, is used for convenience.


The atomic mass unit (amu) is a unit used to describe the very small masses found in single atoms.

The size is based on the mass of one twelfth the mass of one atom of carbon-12. If we drew the atom

to scale and made protons and neutrons 1 cm in diameter, then the electrons would be less than the

diameter of a hair. The entire diameter of an atom would be greater than the length of 30 football

fields!

Need to know check:

o Names of the three main particles found in an atom; where each particle can be found; the

charge and relative mass of each particle.




o RQ4: Draw a model of an atom that shows where each particle can be found and the

charge of each particle.

o RQ5: Describe the relative size of particles in an atom in amu.

Elements and isotopes – how nuclei differ from each another

Elements

If you look at the periodic table you will see that each atom has a single whole number, starting with

“1” for the element hydrogen. This number is unique for that element and is known as the atomic

number. The atomic number is the same as the number of protons in the nucleus of that atom. The

nucleus of every atom of a given element contains the same number of protons. This number can be

used to identify the element and is given the symbol “Z”. For example, every hydrogen atom has 1

proton, and every chlorine atom has 17 protons.

We can think of an atom as an (extremely small) electrically neutral (neither positive nor negative)

unit of matter. For an atom to be electrically neutral, the total number of negative charges must equal

the total number of positive charges. This means that the number of negatively charged electrons

must be the same as the number of positively charged protons. Once you know the number of

protons, you will also know the number of electrons in any atom of any element.

The smallest atom, hydrogen, has one proton (and therefore one electron). Any atom with more than

one proton is not hydrogen. Carbon has atomic number 6. The element carbon is composed of carbon

atoms each with 6 protons and 6 electrons (and neutrons in the nucleus). Argon has atomic number

18. The element argon is composed of argon atoms with 18 protons and 18 electrons (and neutrons).


Need to know check:

o All atoms of an element have the same number of protons. The number of protons in any

atom is the same as the atomic number on the periodic table. The number of electrons is

the same as the number of protons for any atom because an atom is electrically neutral.




o RQ6: Look at your Table from Exercise 6 in this Module. Which column corresponds to Z,

the number of protons in the nucleus?

o RQ7: How many electrons are found in elements with atomic number 2, 4, 6, 8, and 10?

Explain your answer.

Isotopes

All atoms of an element have the same atomic number, Z, but can have different numbers of neutrons.

The number of neutrons is denoted by N. Atoms with the same Z but different N are called isotopes.

Most of the naturally occurring elements, those up to atomic number 92, have isotopes. The number of

protons plus neutrons is called the atomic mass number, A. So, in the nucleus of an atom, we will

find the total number of particles (protons and neutrons) to be:

A = Z + N (Equation 1)

Each element has its own unique element symbol “X”. These three components – the symbol, the

atomic number and the mass number (also called the atomic mass number) – can be used to

represent an individual atom or isotope of an element. This is sometimes called a nuclear symbol.

“A” appears above the element symbol and “Z” below:

For example, chlorine has 17 protons but some atoms of chlorine contain 18 neutrons and others

contain 20 neutrons. These two isotopes of chlorine can be represented as:

.


You can see that the atomic mass number, A, 35 and 37, is different in these two isotopes of chlorine.

Note that the atomic mass number is a whole number and refers to the exact number of protons and

neutrons in an individual atom.

Hydrogen is unlike any other element. It is number 1 in the Periodic Table because its nucleus

contains one proton. It is the only element that contains a nucleus with no neutrons. The nucleus of

most H atoms contains only one proton and no neutrons (A = 1). However, there are two other

isotopes of hydrogen: deuterium (A = 2) and tritium (A = 3). Tritium is radioactive. Hydrogen is the only

element that has been given separate names for its isotopes.

The term atomic mass unit (amu) refers to the mass of a single isotope of an individual atom. When

you look at the periodic table at the element with atomic number 1 (hydrogen), you would expect to

find that it also has an atomic mass number of 1. However, the mass for hydrogen is about 1.01 amu,

and not 1 amu. This is because the number on the periodic table is the average or relative atomic

mass, which is calculated as a “weighted average” of the masses of all the isotopes of an element. It

is called “weighted” because it is adjusted for the abundance of each isotope of the element. This

explains why relative atomic mass is usually not a whole number.

The relative atomic mass is measured by using a mass spectrometer. The mass of every particle is

measured against the mass of a carbon-12 isotope, which has been assigned a value of exactly 12,.

This instrument measures both the masses and the proportions (abundances) of each isotope in a

naturally occurring sample of the element.

Need to know check:

o Isotopes are atoms with the same number of protons but different numbers of neutrons.

The nuclear symbol is made up of the element symbol, the atomic number and the mass

number. The average or relative atomic mass is found on the periodic table. It is a

“weighted average” of the masses of all the isotopes of an element.




o RQ8: Look up the relative atomic masses for the elements lithium, beryllium, boron,

carbon, nitrogen, oxygen, fluorine and neon. Record your answers in a table.

o RQ9: From your table, can you describe a “trend” in the relative atomic masses? Explain

this trend.

o RQ10: You have found that the relative atomic mass of oxygen is 16.00. Explain what this

means.


Electron energy levels

In chemistry we are interested in the electrons and where they are found because these tiny particles

are responsible for how atoms combine together (or do not) to form molecules and compounds!

Our idea or model of the atom has changed from J.J. Thomson’s “plum pudding” model (1898), to

Ernest Rutherford’s planetary model (1909–11), to Neils Bohr’s electron orbital model (1913) through

to the quantum model as more information has been discovered. A more detailed history can be found

at a history of the atom website such as Buescher 2010 Atomic Structure Timeline.

The most modern model of the atom includes the idea that electrons are constantly moving but are

more likely to be found in particular energy levels or “shells”. Increasing numbers of electrons are

found in the energy levels further from the nucleus. However, electrons do not follow regular paths like

cars on a race track or sit in neatly arranged circular shells.

You will learn how to work out how many electrons can be found in each main energy level (or shell).

Electrons repel each other and the number of electrons that each shell can hold is limited to a

maximum number. If we start with the shell closest to the nucleus:

o The first shell (number 1, or n = 1) can hold only 2 electrons

o The second shell (number 2, or n = 2), can hold only 8

o The third shell (number 3, or n = 3), can hold only 18

o The fourth shell (number 4, or n = 4), can hold only 32 …

We can work out the maximum number of electrons in any shell or level by the formula:

Maximum number of electrons = 2 × n2 (Equation 2)

where n is the shell number. In the fifth shell, the maximum number of electrons possible is: 2 × 52 (50

electrons).

Remember that in an atom, the total number of electrons is the same as the total number of protons. If

we look at the diagram of the periodic table below for elements 1 to 20 (from left to right, top row first),

as the number of protons increases by one, the number of electrons will increase by one and fill the

shells in order.

Filling up the energy levels

A position of an element on the Periodic Table tells us information about the electronic configuration of

that element:

If we look at the lithium atom, atomic number 3, it has 3 protons and 3 electrons. The 3 protons are

found in the nucleus, but how do we arrange the 3 electrons? Only 2 electrons can go into the first

shell. That still leaves one electron to go in to the second shell (after Atkin et al. 2009, p. 32).


For the element with atomic number 4, beryllium, two electrons can fit into the second shell.

Each horizontal row of the Periodic Table is called a period. If we start with the first period, the one

closest to the nucleus, you can see that there is only enough space for 2 electrons before the end of

the period is reached. In the second shell, as the atomic number increases by one, the second shell

increases by one electron, until the second shell becomes full with 8 electrons. This corresponds to

element number 10, neon. The third shell (n = 3) can hold a maximum of 18 electrons. However, in the

third period, only eight electrons will fit before we reach the end of the period. For the first 20

elements, only 8 electrons will fit into the second and third energy levels. This is represented below.

Source: <http://sallyann.free.fr/nav011/teaching/teaching-iframe.htm>

The electronic configuration of the first 20 elements in shells or energy levels is shown as a diagram in

the powerpoint presentation “Electron Configuration and the Periodic Table” at the above Internet

address (Parker 2004).

A shorthand way to write this is shown in the diagram below for the first 20 elements. For example,

oxygen (O), atomic number 8, has 8 electrons to arrange around the nucleus. There are 2 electrons in

the first shell and 6 electrons in the second shell of the oxygen atom. This is written as O: 2, 6.



The electronic configuration of the first 20 elements in shells or energy levels is written as shown in

the powerpoint presentation “Electron Configuration and the Periodic Table” at the above Internet

address (Parker 2004).

This is the simplest way of writing the electronic configuration. It is based on Bohr’s model of the atom.

It is a useful starting point for explaining the different energy levels of electrons around the nucleus.

There is a different system for writing the electronic configuration according to energy levels and

sublevels which is based on the later quantum model of the atom. This is explained in your textbook.

What happens when atoms have an atomic number greater than 20? Think about bromine (Br) with

atomic number 35. The electronic configuration can be written as Br: 2, 8, 18, 7. Here the third shell

can fit in the maximum number of electrons (18).

The period number is the same as the number of energy levels the electrons occupy. This affects

certain properties of the elements. However, when we want to know overall about how an element is

likely to react, chemists are mainly interested in the electrons in the outer shell.

The outermost shell of an atom is known as the valence shell.

Going down the groups

The elements are also arranged vertically in columns. The columns are called groups, and are

numbered from left to right. Look at the diagram of the periodic table above. For example, the

configuration of element 11, sodium, is Na 2, 8, 1. Na is found in group 1 of the periodic table. Na also

has 1 electron in its outermost shell. The group number is the same as the number of electrons in

the valence shell. For example, in Group VII (or 17), the elements have the following electron

configurations:

F 2, 7

Cl 2, 8, 7

Br 2, 8, 18, 7

I 2, 8, 18, 18, 7

1 2

2,1 2,2 2,3 2,4 2,5 2,6 2,7 2,8

2,8,1 2,8,2 2,8,3 2,8,4 2,8,5 2,8,6 2,8,7 2,8,8

2,8,8,12,8,8,2

1 2

2,1 2,2 2,3 2,4 2,5 2,6 2,7 2,8

2,8,1 2,8,2 2,8,3 2,8,4 2,8,5 2,8,6 2,8,7 2,8,8

2,8,8,12,8,8,2


Can you tell how many electrons each of these atoms has in its valence shell? Seven!

Electron dot formulas (Lewis dot diagrams)

The electrons in the valence shell determine how an atom will react and which compounds will be

formed. A chemist, G.N. Lewis (1875–1946), devised a shorthand way of writing a formula for an atom

or ion using dots or crosses for the valence electrons. You will find that this way of representing

electrons can be used to show the arrangement of electrons in atoms, ions and compounds (very

handy!).

To write an electron dot formula:

1. draw an imaginary square around the symbol for the atom

2. place electrons one at a time on the four faces of the imaginary square

3. pair up the electrons until the full number has been used up

4. use dots (•) or crosses (x) for electrons but not + signs because these can be confused with

positive charges.

For example, the atom chlorine (Cl) is in group 7 of the Periodic Table and has 7 electrons in the

valence shell. Its dot formula can be written as:

For a step-by-step guide to electron dot diagrams, see “Electron Dot Structure: Shell Diagrams For

The First 20 Elements” (Green Planet Solar Energy 2010).

Need to know check:

o Electrons are found in different energy levels or shells around the nucleus. Each shell can

hold a theoretical maximum number of electrons, but for the first 20 elements, the third

shell can fit only 8 electrons. The arrangement of electrons is called the electron

configuration. You can work out the electron configuration based on how many electrons

will fit into each level (or sublevel). The group number is the same as the number of

electrons in the valence shell. You will be need to be able to draw the electron dot

structures for the first 20 elements.




Cl


o RQ11: Look at the periodic table to find how many electrons there are in the outermost

shell of: (a) lithium, (b) krypton, (c) tin, (d) phosphorus, (e) potassium, (f) barium, (g)

calcium, (h) caesium, (i) silicon, (j) carbon, (k) radium and (l) neon.

o RQ12: Arsenic has electronic configuration: As 2, 8, 18, 5. To which group of the periodic

table does arsenic belong? Check your answer by finding arsenic on your periodic table.

o RQ13: Look at the electronic configuration diagrams above. Write out the electronic

configuration of: (a) sodium, (b) aluminium, (c) helium, (d) beryllium, (e) chlorine, (f)

fluorine, (g) boron, (h) oxygen and (i) potassium.

o RQ14: Explain the difference between an electron dot diagram and the electron

configuration of an atom. Use bromine as an example: Br atomic number 35: Br 2, 8, 18,

7.

3 Bonding

3.1 The periodic table and the Octet rule

Most atoms will bond with other atoms. This tendency to bond is what gives chemicals some of their

properties and drives chemical reactions. Some elements such as sodium or chlorine are highly

reactive and are not found in their pure state in nature. Other elements have very little tendency to

bond with others. These differences can be explained by looking at the group number of the element

in the periodic table and understanding the Octet rule.

All groups of the periodic table have numbers that correspond with the number of electrons in the

valence shell. Some groups also have particular names:

Group 1 alkali metals (except for hydrogen)

Group 2 alkaline earth metals

Groups 3–12 transition metals

Group 17 (or 7) halogens (means “salt-forming)

Group 18 (or 8) noble gases.

It was noticed that the elements in Group 18 (or 8) are not only rare but also tend not to combine with

other elements. In the late 19th century they were named “noble gases”. Unlike other elements, the

noble gases have a full outer shell of electrons. This is the most stable arrangement possible for an

atom. Other atoms will gain or lose electrons to attempt to achieve the most stable arrangement.

The Octet rule helps describe the driving force for chemical reactions and properties. It says that: All

elements gain or lose electrons so they end up with the same electron configuration as the nearest

noble gas on the periodic table. (Guch 2003, p. 80). According to this rule of thumb: How many

electrons would a fluorine atom, F 2, 7, tend to gain to become more stable? (Answer: 1). How many


electrons would an oxygen atom, O 2, 6, tend to gain to become more stable (Answer: 2). How many

electrons would a lithium atom, Li 2, 1, have to lose to become more stable? (Answer: 1).

3.2. Moving electrons to form ions

When an atom gains or loses electrons to become more stable, it no longer has a neutral charge. It

becomes positive or negative and is no longer called an atom. An ion is an atom (or group of atoms)

with a net positive or negative charge.

If an atom loses one electron, it becomes a positive ion. This is because there is one more proton

(positive charge) than the number of electrons (negative charges)

If an atom gains one electron, it becomes a negative ion. This is because there is one more electron

(negative charge) than the number of protons (positive charges)

The following diagram shows the process of atoms either gaining or losing electrons to form ions.

Metals will lose electrons to form positive ions (cations). Non-metals will gain electrons to form

negative ions (anions).


The process of forming ions from atoms is shown as a diagram in the powerpoint presentation

“Structure of the atom” at the above Internet address (Parker 2004).

3.3 Ions – their symbols and names

Chemists need to tell the difference between the element and the ion. This is done by giving ions

different symbols and different names to the atom. The symbol for an ion needs to show the charge on

that ion.

_

+

+

+

_

_

Loss of electron

Overall positive charge on ion

Gain of electron

Overall negative charge on ion

_

+

__

++

+++

+

_

_++

__

__

Loss of electron

Overall positive charge on ion

Gain of electron

Overall negative charge on ion


Chemical shorthand:

The charge on an ion is written above and to the right of its element symbol.

Metal elements lose electrons to form an ion. For example:

Sodium (Na) forms a sodium ion Na+

Magnesium (Mg) forms a magnesium ion Mg2+

Aluminium (Al) forms an aluminium ion Al3+

Non-metals gain electrons to form ions. For example:

Fluorine (F) forms a fluoride ion F–

Oxygen (O) forms an oxide ion O2–

Nitrogen (N) forms a nitride ion N3–

The names for the metal ions are also different to the names of the non-metal ions. The name for a

metal ion is the name of the metal plus the word “ion”. However, when an atom gains electrons to

form a negative ion, its name is changed to end in “– ide”.

Metal atoms, with only a few valence electrons, can achieve a more stable configuration by releasing

their valence electrons and forming positively charged ions. For example, sodium (electron

configuration Na: 2, 8, 1) can achieve a noble gas configuration (2, 8) by releasing its valence

electron. As the atom now has one less electron than it has protons, it has an overall positive charge,

and is written Na+.

Non-metal atoms can achieve a more stable configuration by accepting electrons into their valence

shells. For example, chlorine (Cl: 2, 8, 7) can achieve a noble gas configuration (2, 8, 8) by accepting

one electron to form a negatively charged chloride ion, Cl–.

You can see from the list of ions above, that different elements will lose or gain different numbers of

electrons. For the first 20 elements it is easy to tell how many electrons that will be.

1. Find the atom or element on the Periodic Table.

2. Decide whether the element is a metal or a non-metal.

3. Check the group number on the top of the Table (use the numbers from 1 to 8).

4. If the element is a metal, count down by the group number to find the nearest noble gas. This

number will tell you how many electrons the metal atom will lose to become an ion. For

example, lithium (Li) is in Group 1 of the Periodic Table. It is a metal. Subtract “1” to reach an

element in Group 8, the noble gas helium. The lithium atom has lost one negative charge (an

electron), so the lithium ion has a charge of “+1”.

5. If the element is a non-metal, count up from its group number to 8 to find the nearest noble

gas. This will tell you how many electrons the non-metal atom will gain to become an ion. For


example, oxygen is in group 6 of the Periodic Table. To reach an element in Group 8, two

electrons must be added. The oxygen atom has gained two negative charges, so its ion has a

charge of “–2”.

What about carbon, in group 4 of the periodic table? Carbon has four electrons in its outer shell and it

tends to share those with other atoms rather than losing four or gaining another four.

There are some metals that can form more than one stable ion. These are found in the transition

metals. The charges of the stable ions can not be worked out from their position on the periodic table

and need to be memorized. Examples of elements that can form more than one stable ion are iron

(Fe) and copper (Cu). Iron forms two stable ions, Fe2+ and Fe3+. Copper forms two stable ions, Cu+

and Cu2+. To tell the difference between the ions, the element name is followed by a roman number:

Fe2+ is named the iron (II) ion; Fe3+ is named the iron (III) ion, Cu+, the copper (I) ion, and Cu2+, copper

(II) ion.

Chemical shorthand:

How to show how an atom becomes an ion.

To show how an atom becomes an ion we need:

o the element symbol

o the ion symbol

o a way to write an electron (e–)

o a way to write “plus” (+, written on the line between two items)

o a way to write “minus” (–, written on the line between two items)

o a way to write “goes to make” (→)

If you want to write the sentence, “one chlorine atom gains one electron to make one chloride ion”, you

can write:

Cl + e– → Cl–

Or for the sentence “one magnesium atom loses two electrons to become a magnesium ion”, you can

write:

Mg – 2e– → Mg2+

However, it is preferable to write:

Mg → Mg2+ + 2e–

Copy these sentences in chemical shorthand into your notebook. Congratulations! you have just

written your first chemical equation. This type of equation shows what is happening to the electrons,


and is known as an ionic equation. It can also be called a half equation. Most often you will be writing

net equations that show what happens overall to all the atoms in the reaction. Writing chemical

equations is discussed in Module 3.

Need to know check:

o Atoms of non-metal elements gain electrons to become negative ions and atoms of metal

elements lose electrons to become positive ions. The charge of an ion of the first 20

elements can be worked out by the position of the atom on the periodic table. The process

of forming an ion can be written as an equation.




o RQ15: Work out the names for the negative ions formed from the following atoms: (a)

sulfur, (b) bromine, (c) chlorine, (d) fluorine, and (e) oxygen.

o Describe how each atom either loses or gains electrons to form an ion. Show the charge

on the ion and name the ion formed: (a) barium, (b) fluorine, (c) oxygen, (d) aluminium, (e)

iodine, (f) chlorine, (g) potassium, (h) sodium, (i) calcium, (j) lithium, (k) caesium, (l)

magnesium, (m) sulfur, and (n) bromine. You might like to use an electron dot diagram to

illustrate the process before writing the equation, as shown in Atkin et al. (2009, p. 41); for

example, for barium

which as an equation is written:

Ba → Ba2+ + 2e–

Would you expect the atoms of Group 8 to form ions? Explain your answer.

3.4. From atoms to compounds

Atoms do not react in isolation to gain or lose electrons at random! The Octet rule is a rule of thumb

that can help describe what happens to the electron arrangement in the formation of compounds. For

an atom like sodium to lose an electron, another atom must be present to accept an electron. The

trading of electrons and formation of the new compound has to be energetically favourable, overall

(Atkin et al. 2009, p. 35).


Let’s now think about how particular atoms are likely to react in the formation of compounds.

Previously we have looked at the structure of isolated atoms. However, isolated atoms are rare. Most

matter is made up of atoms joined together to form molecules or compounds. Each compound

represents a unique combination of atoms of the elements held together by attractive forces called

chemical bonds.

Ionic compounds, names and formulas

Many atoms form ions by losing or gaining electrons. But where do these electrons come from? Ionic

compounds are formed from metals (elements with few valence electrons) and non-metals (elements

with nearly full valence shells). The metal atom transfers its electrons to a non-metal atom, which

gains electrons.

Electron dot diagrams showing how one electron is transferred from sodium to chlorine to make the

salt, sodium chloride. (Timberlake 2006).

In the above diagram, one sodium atom (electron configuration 2, 8, 1) loses its valence electron to

form a Na+ ion. One chlorine atom (2, 8, 7) “accepts” the valence electron from the sodium atom to

form a chloride ion (Cl–). The electron is transferred from the sodium atom to the chlorine atom. Both

the Na+ ion and the Cl– ion now have stable “noble gas” electron arrangements.

The compound formed is common table salt, called sodium chloride. The ions of sodium and chlorine

stay together to form a compound because the sodium ion has a positive charge and the chloride ion

has a negative charge. Remember how opposite charges attract each other? These oppositely

charged ions are strongly attracted to each other and thus stay together, forming the electrically

neutral compound sodium chloride. The attraction between the positive and negative ions is referred

to as an ionic bond.

The compound formed between a metal ion and a non-metal ion is called a salt. There are thousands

of salts formed in the process of ionic bonding, and each compound has a unique make-up


(composition). Salts tend to be solids at room temperature, have high melting and boiling points and

many are soluble in water.

When negative and positive ions combine they attract each other and form a regular arrangement of

ions called a crystal lattice. Part of the crystal lattice of sodium chloride is shown below. This

diagram is representing the ions as particles.

Source: http://sallyann.free.fr/nav011/teaching/teaching-iframe.htm

A diagram of the crystal lattice of sodium chloride is shown in the powerpoint presentation “Bonding”

at the above Internet address (Parker 2004).

When we write down the formula for a compound we do not attempt to say how many ions there are

all together. From the diagram above you can see that to write down all the ions would be nearly

impossible. Instead when we write a formula, it tells us the ratio of the numbers of each type of ion in

the compound. Thus, for sodium chloride, there is one sodium ion for each chloride ion. We write the

formula as: NaCl. This is the only correct symbol for sodium chloride, and it is recognised worldwide

as the chemical symbol for this compound. Any other combination, such as “Na2Cl”, meaning that

there are 2 sodium for every chloride, is not only incorrect, but does not exist! The formula of a

compound is a unique combination made up as a result of the properties of the original atoms.

In the above example, sodium chloride, the number of positive ions is equal to the number of negative

ions (a ratio of one to one). However, not all compounds are made of ions that have the same number

of positive and negative charges. What happens when magnesium, which only forms the ion Mg2+ ,

makes a compound with chlorine, which can only form the chloride ion, Cl– ? Magnesium has two

electrons to give away, but chlorine can only take up one. To solve this problem, two chlorine atoms

each take up one electron from one magnesium atom.

Cl– Cl– Cl– Cl– Cl–



Na+ Na+ Na+ Na+

Na+ Na+ Na+ Na+

Na+ Na+ Na+ Na+

http://sallyann.free.fr/nav011/teaching/teaching-iframe.htm


Electron dot diagrams showing how two electrons are transferred from magnesium to two chlorine

atoms to make the salt, magnesium chloride. (Timberlake 2006).

This compound is called magnesium chloride. Now we need to show how many of each atom occur in

the new compound. This number is written to the right of the element symbol but below the line.

Numbers or letters below the line are called subscripts. Here the subscript “2” shows that there are 2

chloride ions for every magnesium ion: MgCl2.

All compounds are electrically neutral; that is, the sum of all the positive charges and all of the

negative charges must be zero. When you are writing formulas, the overall total of positive charges

must equal the overall number of negative charges. If you are having trouble understanding how this

works, get a number of buttons (or something similar) to represent electrons and larger circles to

represent the atoms and make models for the electron dot formulas. Then physically move “electrons”

around.

Chemical shorthand:

How to write and name formulas for ionic compounds

The positive ion is written first in both the name and the formula. The negative ion is written second.

The formula must show the ratio of the ions bonded together. For example, K2O means that there are

2 potassium ions for every oxide ion. (Remember that potassium, in group 1 of the periodic table, can

only form the ion K+). Therefore, 2 potassium ions are needed to each donate one electron to oxygen,

which must accept 2 electrons to form the oxide ion, O2–.

The name does not show the ratio of the ions bonded together. For example, K2O is called “potassium

oxide”. This assumes that you know how to write the correct formula!


When naming compounds that contain transition metal elements, the charge of the ion must be

included. For example FeCl2 is named “iron (II) chloride” and FeCl3 is “iron (III) chloride”.

The atoms of Groups 1 and 2 generally form ionic compounds because they form positive ions in

losing electrons to have the same electron configuration as their closest noble gas.

Need to know check:

o How an ionic bond is formed to make an ionic compound; the properties and general

composition of a salt; what is meant by a crystal lattice; how to name and write the correct

formula of an ionic compound.




o RQ16: What is the formula of the compound formed between the following sets of atoms:

(a) sodium and bromine, (b) potassium and sulfur, and (c) magnesium and nitrogen?

o RQ17: Write the formulas for the following compounds: (a) sodium oxide, (b) lithium

bromide, (c) calcium chloride, (d) aluminium oxide, (e) sodium chloride, (f) magnesium

chloride, (g) potassium iodide, (h) barium sulfide, (i) aluminium sulfide, (j) sodium fluoride.

Molecular compounds, names and formulas

When two non-metal atoms both need to gain electrons to become like the nearest noble gas, they

can achieve this efficiently by sharing electrons. The group of atoms held together by the sharing of

electrons is called a molecule. Non-metals bond together by forming covalent bonds.

The atoms in a covalent bond are held together by the electrostatic attraction between the positively

charged nuclei and the shared electrons. The force of attraction that results from sharing electrons is

called covalent bonding.

We know that hydrogen exists in the air in a fairly stable form. But it can’t be present in the air as H

atoms as they do not have a full outer shell. How can hydrogen combine with more of itself to form a

more stable molecule? Hydrogen is formed from two hydrogen atoms joined together by sharing two

electrons. The sharing of two electrons is called a single covalent bond. This is called a diatomic

molecule (“two atoms”).

.



A diagram of the molecule of hydrogen is shown in the powerpoint presentation “Bonding” at the

above Internet address (Parker 2004).

The hydrogen atoms have joined together. By sharing their electrons, each hydrogen atom has now

obtained 2 electrons in its outer shell. Each hydrogen is now like the noble gas helium and now

effectively has two electrons in its outer shell. The formula for hydrogen is H2. This molecular formula

gives information about the actual number of atoms present in each molecule of the substance.

Oxygen also exists as a diatomic molecule. Each oxygen atom has 6 valence electrons. By sharing

two pairs of electrons, each oxygen atom has gained 8 electrons in its outer shell. Since two pairs of

electrons are shared between the two atoms, a double covalent bond is formed.


A diagram of the molecule of oxygen is shown in the powerpoint presentation “Bonding” at the above

Internet address (Parker 2004).

Both oxygen atoms now have a total of 8 electrons in their outer shells. The sharing of the pair of

electrons between the two oxygen atoms binds the two atoms together. The formula for the oxygen

molecule is written as: O2. This formula tells us that two oxygen atoms form a molecule.

One, two or three pairs of electrons can be shared between two atoms. A single covalent bond is

formed by two shared electrons. A double covalent bond is formed by four shared electrons, and a

triple covalent bond is formed by six shared electrons.

11

O O




Other diatomic molecules

Under room temperature and pressure other common non metals also exist as diatomic

molecules:

Nitrogen N2(g)

Fluorine F2(g)

Chlorine Cl2(g)

Bromine Br2(l)

Iodine I2(s)

Covalent bonding can occur between atoms of different elements. When hydrogen and nitrogen

combine to form ammonia, each hydrogen atom requires 1 electron to complete its outer shell and

each nitrogen atom requires three electrons. Each nitrogen atom shares a one electron with three

hydrogen atoms to form an ammonia (NH3) molecule.


A diagram of the molecule of ammonia is shown in the powerpoint presentation “Bonding” at the

above Internet address (Parker 2004).

When writing a formula for compound, which element comes first? When writing the formula for an

ionic compound, we said that the metal was written first. This is easy to follow for an ionically bonded

compound. But molecular compounds are formed between non-metals. Writing and naming all

formulas follows the rules of electronegativity.

Electronegativity is the ability of an element to attract an electron. Non-metals are always more

electronegative than metals because they gain the electron in ionic bonding. Metals can be thought of

as being the opposite of electronegative, electropositive.

When you look at the periodic table there are some trends in electronegativity. First, electronegativity

increases from left to right across the periodic table (from metals to non-metals). For any two elements

in a period, the one further to the right is the more electronegative. For example, oxygen is more

electronegative than nitrogen. Second, electronegativity decreases from the top of the group to the

bottom. For example, sodium is less electronegative (or more electropositive) than lithium. These two

H N

H

H



trends combine to give us the overall picture: the most electronegative element is fluorine (upper right)

and the least is caesium (Cs) (lower left) (francium is radioactive).

Electronegativity of representative elements showing the ability of atoms to attract electrons.

(Timberlake 2006).

Chemical shorthand:

How to write and name formulas for molecular compounds

A molecule is formed when two or more atoms share electrons. A molecule has no net charge.

A covalent bond is the sharing of electrons between atoms. In covalent bond formation, atoms go as

far as possible toward completing their octets by sharing electrons.

Usually the same number of electrons is donated from each atom. In a molecule of hydrogen gas, one

electron comes from each H, for example.

In forming molecules and covalent bonds, there is a limit to the number of electrons an atom can have

in its bonding shell. For example, F cannot take on more than 8 in its outer shell. H cannot take on

more than 2 in its outer shell because this is the maximum number of electrons that can “fit”in the

valence shell.

We use the electronegativity trend to decide which element of a molecule to write first. By convention,

the more electropositive element is named first (in ionic compounds, this is the metal). The more

electronegative element is written second, with its ending changed to “ide”. For example, the molecule

formed between carbon and oxygen is written “CO”, not “OC”, because carbon is more electropositive

than oxygen.

Unlike most ionic compounds, there can often be more than one combination between non-metal

atoms. For example CO and CO2 are not the same molecule and have to be named differently to tell

them apart. CO is called “carbon monoxide” and CO2 is called “carbon dioxide”. The prefixes mono, di,


tri, tetra, penta, or hexa are used to indicate 1, 2, 3, 4, 5 or 6 atoms of the more electronegative

element. The prefix mono is not used for the more electropositive element.

Some common compounds are named non-systematically: H2O, water; NH3, ammonia; and CH4

methane.

Composite ions – charged covalently bonded ions

Some non-metals bond together covalently to form stable structures, but these structures have an

overall charge. These are called composite or polyatomic ions. Some common ones are included in

the table below, together with an example of the type of compound formed.

Composite ion Name of ion Example of a salt formed

Compound name

NH4+ Ammonium NH4Cl Ammonium chloride

NH2– Amide NaNH2 Sodium amide

OH– Hydroxide KOH Potassium hydroxide

CN– Cyanide LiCN Lithium cyanide

O22– Peroxide BaO2 Barium peroxide

CO32– Carbonate K2CO3 Potassium carbonate

HCO3– Bicarbonate,

hydrogen

carbonate

NaHCO3 Sodium bicarbonate,

sodium hydrogen carbonate

(baking soda)

NO2– Nitrite NaNO2 Sodium nitrite

NO3– Nitrate NaNO3 Sodium nitrate

SO32– Sulfite MgSO3 Magnesium sulphite

SO42– Sulfate MgSO4 Magnesium sulphate

(Epsom salts)

Some polyatomic ions, for example bicarbonate, are important in the human body for helping to keep

blood and other fluids from becoming too acidic or basic.

When you write the formula for a compound formed with a composite ion, think of the ion as a single

item with a charge. It’s the charge that affects the formula for the salt overall.

Sometimes brackets are needed to show the correct numbers of atoms in the compound. For

example, ammonium oxide is written: (NH4)2O. The brackets make it clear that two ammonium ions

need to bond with one oxide to form an electrically neutral compound. This means that there are a

total of two nitrogen atoms, eight hydrogen and one oxygen atom in the compound. Another example

is magnesium hydroxide, written as Mg(OH)2.


Need to know check:

o How a covalent bond is formed to make a covalent or molecular compound; the properties

and general composition of a molecular compound; the formulas for common diatomic

molecules; what is meant by a composite or polyatomic ion; how to name and write the

correct formula of a molecular compound; the names of some common molecules with

non-systematic names; how to name and write the correct formula of a composite ion and

salts made from them.




o RQ18: Describe how a covalent bond is formed between fluorine atoms to make the

fluorine molecule. How many electrons are needed for each atom to complete its outer

shell? What type of covalent bond is formed?

o RQ19: How many electrons make up a double covalent bond? Name a molecule that

contains a double covalent bond.

o RQ20: How many electrons make up a triple covalent bond. Name a molecule that

contains a triple covalent bond.

o RQ21: Write the correct molecular formula for the following molecular compounds: (a)

water, (b) ammonia, (c) methane.

o RQ22: Write the correct formulas for the following molecular compounds: (a) sulphur

dioxide, (b) carbon disulfide, (c) diphosphorus pentoxide.

o RQ23: Write the correct formula for the following polyatomic ions: (a) hydroxide, (b)

ammonium, (c) hydrogen carbonate, (d) carbonate, (e) sulphate.

o RQ24: Write the correct names for the following compounds: (a) CaCO3, (b) Ca(HCO3)2,

(c) (NH4)2CO3, (d) K2SO4, (e) NH4OH


APPENDIX A:

Enabling Course Chemistry MODULE 2

ANSWERS to Review Questions

RQ1 & 3: .

Element number

Symbol Name Type of element

1 H Hydrogen Non-metal

2 He Helium Non-metal

3 Li Lithium Metal

4 Be Beryllium Metal

5 B Boron Metalloid

6 C Carbon Non-metal

7 N Nitrogen Non-metal

8 O Oxygen Non-metal

9 F Fluorine Non-metal

10 Ne Neon Non-metal

11 Na Sodium Metal

12 Mg Magnesium Metal

13 Al Aluminium Metal

14 Si Silicon Metalloid

15 P Phosphorus Non-metal

16 S Sulfur Non-metal

17 Cl Chlorine Non-metal

18 Ar Argon Non-metal

19 K Potassium Metal

20 Ca Calcium Metal

25 Mn Manganese Metal

26 Fe Iron Metal

27 Co Cobalt Metal

29 Cu Copper Metal

30 Zn Zinc Metal

34 Se Selenium Non-metal

35 Br Bromine Non-metal

47 Ag Silver Metal

50 Sn Tin Metal

53 I Iodine Non-metal

79 Au Gold Metal

80 Hg Mercury Metal

82 Pb Lead Metal

RQ2: Ans: Metals have high melting and boiling points which means that most are solids at room

temperature (except for mercury, a liquid); metals can be polished to shine, beaten into various shapes and drawn into wire – and are good conductors of heat and electricity. Non-metals have low melting and boiling points, are dull and brittle, and are poor conductors of heat and electricity.


RQ4: Draw a model of an atom that shows where each particle can be found and the charge of each particle.

Ans: The proton has a positive charge and is found in the nucleus; the neutron has no charge and is found in the nucleus; the electron has a negative charge and is found at a distance from the nucleus. A diagram such as the one below could be drawn.

From Atkin et al. 2009, p. 26. OR Parker 2004. Structure of the atom.

RQ5: Describe the relative size of particles in an atom. Ans: The proton and neutron are about the

same size and the electron is 1/1836th or about one-two thousandth (1/2000) the size of a proton or neutron.

RQ6: Ans: Column 1, element number. RQ7: Ans: 2, 4, 6, 8, 10 – the number of electrons is the same as the number of protons, the atomic

or element number. RQ8: Relative atomic masses

Element Relative atomic mass

Lithium 6.94

Beryllium 9.01

Boron 10.81

Carbon 12.01

Nitrogen 14.01

Oxygen 16.00

Fluorine 19.00

Neon 20.18

Neutron Neutron


RQ9: Ans: The relative atomic mass increases along a period from metals to non-metals (left to right). The relative atomic mass of each atom increases as the atomic number or number of protons increases.

RQ10: Ans: The relative atomic mass of 16.00 means that this is the average weighted atomic mass

of the isotopes of oxygen. It means that on average, over all the isotopes of oxygen, an oxygen atom has 8 protons and 8 neutrons.

RQ11: Look at the periodic table to find how many electrons there are in the outermost shell of (the

group number): (a) lithium, 1 (b) krypton 8 (c) tin 4 (d) phosphorus 5 (e) potassium 1 (f) barium 2 (g) calcium 2 (h) caesium 1 (i) silicon 4 (j) carbon 4 (k) radium 2 (l) neon 8

RQ12: Ans: Arsenic is in group 5 of the periodic table. It has 5 electrons in the valence shell. RQ13: Ans: Look at the electronic configuration diagrams above. Write out the electronic

configuration of: (a) sodium, Na: 2, 8, 1 (b) aluminium, Al: 2, 8, 3 (c) helium, He: 2 (d) beryllium, Be: 2, 2 (e) chlorine, Cl: 2, 8, 7 (f) fluorine, F: 2, 7 (g) boron, B: 2, 3 (h) oxygen: O: 2, 6 (i) potassium. K: 2, 8, 8, 1

RQ14: Ans: Bromine atomic number 35 has an electron configuration in shells of: 2, 8, 18, 7. The

outermost shell has 7 electrons. The electron dot diagram shows the electrons in the valence shell only.

RQ15: Work out the names for the negative ions formed from the following atoms: Ans:

(a) sulfur sulfide (b) bromine, bromide (c) chlorine, chloride (d) fluorine fluoride (e) oxygen. Oxide

Br


RQ16: What is the formula of the compounds formed between the following sets of atoms: Ans:

(a) sodium and bromine NaBr (b) potassium and sulphur K2S (c) magnesium and nitrogen Mg3N2

RQ17: Write the formulas for the following compounds: Ans:

(a) sodium oxide Na2O (b) lithium bromide LiBr (c) calcium chloride CaCl2 (d) aluminium oxide Al2O3 (e) sodium chloride NaCl (f) magnesium chloride MgCl2 (g) potassium iodide KI (h) barium sulphide BaS (i) aluminium sulfide Al2S3 (j) sodium fluoride NaF.

RQ18: Ans: Describe covalent bonding in a fluorine molecule. Each fluorine atom has seven

electrons in the valence shell. It needs 8 to form a stable octet. Therefore each atom needs one electron.The fluorine atoms can share ONE PAIR of electons to form a stable octet. This forms a SINGLE covalent bond.

RQ19: Ans: Two PAIRS of electrons, or four electrons, make up a double covalent bond. An oxygen

molecule (O2) contains a double covalent bond between two oxygen atoms. RQ20: Ans: Three PAIRS of electrons make up a triple covalent bond. A nitrogen molecule (N2)

contains a triple covalent bond between two nitrogen atoms. RQ21: Write the correct molecular formula for the following molecular compounds: Ans:

(a) water H2O (b) ammonia NH3 (c) methane CH4.

RQ22: Write the correct formulas for the following molecular compounds: Ans:

(a) sulfur dioxide SO2 (b) carbon disulfide CS2 (c) diphosphorus pentoxide P2O5.

RQ23: Write the correct formula for the following polyatomic ions: Ans:

(a) hydroxide OH– (b) ammonium NH4

+ (c) hydrogen carbonate HCO3

– (d) carbonate CO3

2– (e) sulphate SO4

2–.


RQ24: Write the correct names for the following compounds: Ans:

(a) CaCO3, calcium carbonate (b) Ca(HCO3)2, calcium hydrogen carbonate, calcium bicarbonate (c) (NH4)2CO3 ammonium carbonate (d) K2SO4 potassium sulfate (e) NH4OH ammonium hydroxide

chemistry module 2 enabling course · this will be discussed in section 2.2 of this module. 1.3...

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