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CHAPTER

2The atomic theory of matter

the scientifi c method in the study of chemistry

the historical development of the model of the

atom, in particular the theories of Democritus and

John Dalton, and those scientists who contributed

to the discovery of the subatomic particles, the

nucleus and isotopes — Joseph Thomson, Ernest

Rutherford, Niels Bohr, James Chadwick, Henry

Moseley, Erwin Schrödinger, Marie Curie and

Frederick Soddy

••

the nature of radioactivity and its role in the

understanding of atomic structure

the use of the mass spectrometer in identifying isotopes

of elements

relative atomic masses of elements

the structure of the atom and the relationship of isotopes

to its subatomic particles

the electronic confi guration of the fi rst thirty elements of

the periodic table

•

•

••

•

The public’s fi rst awareness of X-rays and

their ability to penetrate fl esh to produce

skeletal images on fi lm came through

photographs like this, which was taken in

1896 by the American professor Michael

Pupin. The black spots are birdshot pellets

embedded in Pupin’s hand during a hunting

accident. Today X-rays are used extensively

for the medical diagnosis of fractures and

deformities. X-rays are also used extensively

in various branches of science. For example,

they are used to determine the arrangement

of atoms in crystal (crystal structure) because

their wavelengths are of a suitable size for

detecting the relative positions of atoms.

Studies of X-rays produced by atoms have

also provided important evidence about the

structure of atoms.

You will examine:

UNIT 1 The big ideas of chemistry28

Early ideas about matterLong before the process of scientific method was developed, humans have searched for answers to questions like ‘What is matter?’ and ‘What is the composition of matter?’

The ancient Greek philosophers, from around 600 BC, were the first people in recorded history to offer theories about the nature of matter. However, since their theories did not have an experimental basis, they could be neither supported nor rejected by experimental evidence. The philosophers relied on their powers of logical persuasion to convince others of their views.

In about 450 BC Empedocles (c. 490–c. 420 BC) suggested that all matter was made from combinations of four different ‘elements’: fire, air, water and earth.

Almost fifty years later, this view was opposed by another Greek philosopher, Democritus, (c. 470–c. 400 BC) who proposed what may be considered to be the first atomic theory of matter. He reasoned that if a sample of matter was cut into smaller and smaller pieces, a point would eventually be reached at which the pieces could not be further subdivided. At this point, the particles were indivisible. He named these very small particles atoms, from the Greek word atomos, meaning ‘unable to be divided’. He believed that atoms were in constant motion, that they combined with others in various ways and that they differed from each other only in shape and arrangement.

Aristotle (384–322 BC), a well-respected and influential Greek philosopher, opposed Democritus’ ideas and insisted that matter was continuous and infinitely divisible. He supported Empedocles’ four-element theory, proposing that it was possible to change one material into another (such as copper to gold) by varying the proportions of the four elements, namely fire, air, water and earth.

X-ray molecular diffraction of DNA

molecules. Images like this are used

to develop the understanding of the

structure of complex molecules.

X-ray molecular diffraction of DNA

molecules. Images like this are used

to develop the understanding of the

structure of complex molecules.

Democritus hypothesised the

existence of indestructible

particles that he called atoms,

from the Greek word atomos,

meaning ‘unable to be divided’.

Democritus hypothesised the

existence of indestructible

particles that he called atoms,

from the Greek word atomos,

meaning ‘unable to be divided’.

Aristotle adopted Empedocles’

theory that all matter is

continuous and is composed of

combinations of fire, air, water

and earth.

Aristotle adopted Empedocles’

theory that all matter is

continuous and is composed of

combinations of fire, air, water

and earth.

29CHAPTER 2 The atomic theory of matter

Alchemists

Although they relied heavily on experimentation, their activities were not guided by the development of scientific theories. Still, they developed a number of useful techniques such as distillation, evaporation, crystallisation and filtration. The idea and reality of the laboratory itself — and its equipment — began with the alchemists.

Aristotle rejected the idea of atoms, but accepted and refined the notion of four elements — water, air, fire andearth. Aristotle’s theory provided the philosophical basis of alchemy. During the period AD 500–1600 the practice of alchemy flourished. One of the major goals of an alchemist was to discover a substance that would turn base metals such as iron, copper and lead into gold. This substance was referred to as the ‘philosopher’s stone’. Among those who believed in alchemy were Isaac Newton (1643–1727) and Robert Boyle (1627–1691), who we see today as founders of modern science.

John Dalton (1766–1844) suggested that there were atoms that differed for each element and that they combined to make other substances. Robert Boyle (1600s) and Antoine Lavoisier (1700s) were the first scientists to begin testing

ideas by carrying out experiments and were credited with putting experimental systems and scientific method together in the study of chemistry. Often the advances made by one scientist depend on the earlier work of others or a collaborative effort by scientists from several research laboratories.

The discovery of electrolysis was very important in the history of the search for knowledge about atoms because it made possible the breaking down of substances that could not previously be split up. In the early 1800s, the English chemist Humphrey Davy, and later his assistant Michael Faraday, experimented with passing an electric current through substances. They found that they could break down some substances into their elements. For example, in 1807 Davy made the exciting discovery of sodium and potassium by passing electricity through molten salts. He isolated barium, calcium and magnesium

Alchemists tried to turn one

type of substance into another

using the ‘philosopher’s stone’.

They failed to turn lead into

gold.

Alchemists tried to turn one

type of substance into another

using the ‘philosopher’s stone’.

They failed to turn lead into

gold.

The alchemists devised chemical

apparatus and accumulated

knowledge about chemical

changes, but didn’t succeed in

turning lead into gold. This failure,

along with the growth of analytical

(rather than descriptive) knowledge

of chemical change, led to the

decline in belief in Aristotle’s ideas.

The alchemists devised chemical

apparatus and accumulated

knowledge about chemical

changes, but didn’t succeed in

turning lead into gold. This failure,

along with the growth of analytical

(rather than descriptive) knowledge

of chemical change, led to the

decline in belief in Aristotle’s ideas.

2.1 Making a spiral periodic table2.1 Making a spiral periodic table

Dalton’s theory stated that all

atoms of the same element are

identical, and atoms of different

elements are totally different.

Dalton’s theory stated that all

atoms of the same element are

identical, and atoms of different

elements are totally different.

This controlled burn in Kakadu

National Park combines the

elements of earth, air, fire and

water. Aristotle thought that these

elements of nature recombined

to form different substances, for

example wood burns to form ash,

smoke and water vapour.

This controlled burn in Kakadu

National Park combines the

elements of earth, air, fire and

water. Aristotle thought that these

elements of nature recombined

to form different substances, for

example wood burns to form ash,

smoke and water vapour.


Timeline of scientists’ contributions

to understanding of matter

ANCIENT GREEKS (500–400 BC)

reasoned and hypothesised

Democritus

Aristotle: Matter is continuous rather than atomistic

Dalton (1808)

experimented, speculated, theorised

empirical model: hard and

indivisible atoms

formulated the ‘Law of multiple

proportions’ to test this theory

Thomson (1903)

experimental model: solidsphere of positive chargewith negative electrons

experimented, calculated, theorised

Rutherford (1911)


experimental model: nuclearatom — positive nucleuscontaining protons,electrons outside

Bohr (1913)

reasoned and mathematically analysed

mathematical model: electrons in circularorbits and protonsin nucleus

OUR PRESENT CONCEPT OF THE ATOM

Electronic structure

Nuclear structure

quantum model — a complexmathematical model — electronlocation as probability cloud

many elementary particles inthe nucleus — more than200 elementary particles have been discovered so far

2200 y

ears

95 y

ears

8 y

ears

2 y

ears

10 y

ears

55 +

years

–

––

• ••

•

•• •

••

• •

• • • •

•

•

••

+–

––

––

–+

+ +

+

++

+

––

–

+––

Goldstein(1866)

positive raysprotons

Crookes(1879)

cathode rayselectrons

Soddy(1910)

discovery ofisotopes

Moseley(1914)

concept ofatomic number

Planck(1900)

quantumtheory

De Broglie(1923)

particle–wavedualism forthe electron

Chadwick(1932)

discovery ofneutron

Today

?

1932

Pn

Becquerel,Pierre andMarie Curie

(1896)discovered

radioactivity

Rutherford,Marsden andGeiger (1911)

gold foilexperiment

Lavoisier(1789)‘Law ofmass’

philosophical model:

indestructible atomos

∞∞ ∞∞∞∞ ∞

∞∞ ∞

ANCIENT GREEKS (500–400 BC)

reasoned and hypothesised

Democritus

Aristotle: Matter is continuous rather than atomistic

Dalton (1808)

experimented, speculated, theorised

empirical model: hard and

indivisible atoms

formulated the ‘Law of multiple

proportions’ to test this theory

Thomson (1903)

experimental model: solidsphere of positive chargewith negative electrons


Rutherford (1911)


experimental model: nuclearatom — positive nucleuscontaining protons,electrons outside

Bohr (1913)

reasoned and mathematically analysed

mathematical model: electrons in circularorbits and protonsin nucleus

OUR PRESENT CONCEPT OF THE ATOM

Electronic structure

Nuclear structure

quantum model — a complexmathematical model — electronlocation as probability cloud

many elementary particles inthe nucleus — more than200 elementary particles have been discovered so far

2200 y

ears

95 y

ears

8 y

ears

2 y

ears

10 y

ears

55 +

years

–

––

• ••

•

•• •

••

• •

• • • •

•

•

••

+–

––

––

–+

+ +

+

++

+

––

–

+––

Goldstein(1866)

positive raysprotons

Crookes(1879)

cathode rayselectrons

Soddy(1910)

discovery ofisotopes

Moseley(1914)

concept ofatomic number

Planck(1900)

quantumtheory

De Broglie(1923)

particle–wavedualism forthe electron

Chadwick(1932)

discovery ofneutron

Today

?

1932

Pn

Becquerel,Pierre andMarie Curie

(1896)discovered

radioactivity

Rutherford,Marsden andGeiger (1911)

gold foilexperiment

Lavoisier(1789)‘Law ofmass’

philosophical model:

indestructible atomos

∞∞ ∞∞∞∞ ∞

∞∞ ∞

by 1808 and later established that boron, iodine and chlorine were also elements. Davy’s experiments suggested that the bonding between elements (atoms) was electrical. Much later, around the beginning of this century, scientists showed experimentally that matter itself was electrical. They also found that the atoms of some elements could change into atoms of other elements, a process called radioactivity. Much of modern physics is concerned with the particles that make up atoms, and scientists have built particle accelerators like the Australian Synchrotron to smash atoms so that they can study the subatomic particles that make up atoms.


Robert Boyle and Antoine Lavoisier — chemistry becomes a science

Robert BoyleIt was not until 1661, more than 2000 years after Democritus’ time, that the British scientist Robert Boyle emphasised the necessity of using experiments to test ideas which had been obtained by reason. Using the techniques of chemical analysis known to him, Boyle attempted to isolate Aristotle’s four ‘elements’.

Robert Boyle was the first to give the word ‘element’ its modern meaning. He defined an element as being a substance that could not be broken down into simpler substances. In Boyle’s time, fourteen elements were known, but this number was soon to increase dramatically. He was able to distinguish clearly between elements, compounds and mixtures.

Antoine LavoisierIn one of many experiments to investigate the nature of combustion, Lavoisier placed an accurately weighed amount of tin into a flask, sealed it, weighed it, heated it until no more of the powdery oxide was formed, and then reweighed the flask. No weight change was found. He then broke the seal on the flask, allowing air to enter, and weighed the flask again. In this case, he found a weight increase. Lavoisier attributed the weight increase to the weight of the air rushing in when the flask was opened and concluded that this weight had replaced the weight of the air originally in the closed apparatus, which had been used up in forming the oxide.

Lavoisier changed chemistry from a qualitative to a quantitative science. He showed that the mass of the products in a reaction is equal to the mass of the reactants, hence proving the law of conservation of mass. By proving that mass was neither created nor destroyed, he was able to disprove Aristotle’s ‘four-element’ theory that matter was composed of earth, water, air and fire. He showed that water was a compound, air was a mixture, earth contained numerous elements, and fire was a chemical process involving oxygen.

In 1789, Lavoisier and a small group of other scientists created the method of chemical nomenclature that classified the differences between elements. Lavoisier also clarified the distinction between elements and compounds.

Robert Boyle was an infant prodigy

and, in addition to his research

with acids and bases, he was

recognised for his investigations

of gases, dyes, glass, minerals,

phosphorus, qualitative analysis,

combustion, corrosion and the

nature of fire, heat and cold. In one

famous experiment, he placed a

burning candle in a jar and removed

the air from the jar to extinguish

the flame, demonstrating that

something in air is necessary for

combustion.

Robert Boyle was an infant prodigy

and, in addition to his research

with acids and bases, he was

recognised for his investigations

of gases, dyes, glass, minerals,

phosphorus, qualitative analysis,

combustion, corrosion and the

nature of fire, heat and cold. In one

famous experiment, he placed a

burning candle in a jar and removed

the air from the jar to extinguish

the flame, demonstrating that

something in air is necessary for

combustion.

Antoine Lavoisier is known as

the father of modern chemistry

and is famous for his theories

of burning, proving the law of

conservation of mass.

Antoine Lavoisier is known as

the father of modern chemistry

and is famous for his theories

of burning, proving the law of

conservation of mass.

To find out how atomic theory

developed, go to the website

for this book and click on the

Atomic theory weblink (see

Weblinks, page 531).

To find out how atomic theory

developed, go to the website

for this book and click on the

Atomic theory weblink (see


An element combined with another

element produces a compound.

This ratio is the same for any size

sample, large or small. Lavoisier

made the distinction between

elements and compounds.

An element combined with another

element produces a compound.

This ratio is the same for any size

sample, large or small. Lavoisier

made the distinction between

elements and compounds.


John Dalton — the fi rst modernatomic theoryIn 1808 the English headmaster and scientist John Dalton (1766–1844) revived and greatly expanded the atomic theory originally proposed by Democritus. Dalton’s atomic theory included the following ideas:

1. All matter is composed of extremely small particles called atoms that are held together by forces of attraction.

2. Atoms are indivisible and cannot be created or destroyed.

3. Atoms of the same element are identical and have the same mass.

4. Atoms of different elements have different masses.

5. Chemical reactions occur when atoms are separated, joined or rearranged. However, atoms of one element are not changed into atoms of another by a chemical reaction.

6. Atoms of different elements can combine with one another in simple whole number ratios to form compounds (for example, 1 : 1, 1 : 2, 2 : 3, 1 : 3, and so on).

Although some of Dalton’s ideas were later modifi ed, his fundamental idea that matter is made up of atoms is still valid. Despite its fl aws, Dalton’s theory inspired generations of scientists to conduct experiments that would clarify and refi ne our concept of the fundamental particles of matter.

John Dalton was the fi rst person to sketch the structures of many chemical compounds. Like many chemists before him, he devised his own set of symbols for the known elements.

Revision questions

1. John Dalton is regarded as the founder of modern atomic theory. In addition

to elaborating on Democritus’ view that all matter was built up from tiny,

indivisible particles called atoms, he prepared the fi rst table of atomic

weights. His description of an atom was often called the ‘billiard-ball’ model

of atomic structure.

(a) In what ways can Dalton’s atom be directly compared with a billiard

ball?

(b) In what ways is a billiard ball not comparable with an atom?

2. It has been suggested that a nation’s wealth would be indicated not by its

store of gold but by its ability to produce solid sodium from sea water. Can

you explain the reasoning behind this suggestion?

Subatomic particlesWith the help of increasingly powerful atom smashers, physicists have found numerous subatomic particles — bosons, leptons, hadrons and quarks are a few. So many particles have been found, in fact, that no single theory of atomic structure can account for them all. Within the scope of this course, however, we are concerned with the discovery of only three of these subatomic particles: electrons, protons and neutrons.

The fi rst subatomic particles were discovered as a result of the work of scientists investigating electricity. These scientists studied the fl ow of electric current through gases at low pressure in a discharge tube. This is a sealed glass tube containing the gas being investigated, with metal discs called electrodes

John Dalton


at each end. When connected to a high voltage, a bright green glow appears in the tube. One electrode, the anode, becomes positively charged while the other electrode, the cathode, becomes negatively charged. The glowing beam which travels from the cathode to the anode is called a cathode ray.

This cathode ray oscilloscope

shows a voltage versus time trace

sin wave. The green glow is caused

by excited electrons, directed by an

electric field, hitting the screen.

This cathode ray oscilloscope

shows a voltage versus time trace

sin wave. The green glow is caused

by excited electrons, directed by an

electric field, hitting the screen.

Since cathode rays travel towards the positively-charged anode, it was clear that they carry a negative charge. Scientists such as the English physicist J.J. Thomson (1856–1940) further discovered that cathode rays consist of tiny, negatively-charged particles, later called electrons.

TABLE 2.1 Properties of subatomic particles

Particle Symbol

Relative electrical

charge

Approximate relative

mass (amu)*

electron e 1– 1/1840

proton p 1+ 1

neutron n 0 1

* 1 amu = 1.66 × 10–24 g

Joseph John Thomson — discovery of the electronJoseph John (‘J.J.’) Thomson carried out quantitative experiments on the nature of cathode rays. In 1897, using a modified discharge tube such as the one shown below he discovered that the electrons that made up the cathode rays could be deflected by magnets or electrically charged plates.

Thomson found that in all cathode-ray experiments, regardless of the gas used in the cathode-ray tube or the type of metal used for the electrodes, the same charge-to-mass ratio was obtained. Thomson concluded that electrons must be a part of the atoms of all elements.

Thomson’s discovery of the electron, and the knowledge that atoms were neutral in charge overall, led him to propose a model for the internal structure of an atom. Since the negatively charged electrons must be balanced by positive

Joseph John Thomson, commonly

known as ‘J.J.’, discovered the

electron by using a specially built

discharge tube. He subsequently

proposed a ‘plum-pudding’ model of

atomic structure.


charges, he proposed that an atom was a sphere of positive charges throughout which tiny negative electrons were loosely embedded, like plums in a pudding. Every element had a different number of electrons in its atoms. This model became known as the ‘plum-pudding’ model of atomic structure.

Cathode

Positive

plate

Negative

plate

Magnet Zinc sulfide

detecting screen

Anode + –

+

S

N

Discovery of radioactivityIn 1896 the French chemist Antoine Becquerel (1852–1908) accidentally discovered that uranium ores emit invisible rays. While experimenting with minerals, he happened to store some uranium ores in a drawer with some

photographic plates. When he wanted to use the plates he found that they appeared to have been partially exposed. He investigated the cause of the fogging on the plates and discovered that uranium emitted rays that were able to expose a photographic plate even when the uranium and plate were separated by black paper of considerable thickness. Two of his graduate students, Marie Curie (1867–1934) and her husband Pierre Curie (1859–

1906), were interested in investigating this question further. They experimented with pitchblende, a uranium ore, and after two years of research isolated two new elements that gave off these rays: radium and polonium. They discovered that uranium and the new elements, radium and polonium, were disintegrating over time and emitting radiation that exposed Becquerel’s photographic plates. Marie gave the name radioactivity to the property.

Ernest Rutherford — discovery of the nucleusThe New Zealand physicist Ernest Rutherford (1871–1937) used radioactivity as a tool to investigate the structure of the atom. In 1911, Rutherford and his students Hans Geiger (1882–1945) and Ernest Marsden performed a classic experiment to test Thomson’s ‘plum-pudding’ atomic model. They bombarded a thin gold foil with positively charged alpha particles from a radioactive source. Almost all the particles passed through the gold foil (which was about 1000 atoms thick) but about one in every 10 000 bounced back, indicating repulsion by a highly concentrated positive charge.

The results of the experiment astounded Rutherford because, according to calculations based on the Thomson model, the alpha particles were expected

Thomson proposed that atoms

were spheres of positive charge

with small negative electrons

loosely embedded in them like

plums in a plum pudding.

Thomson proposed that atoms

were spheres of positive charge

with small negative electrons

loosely embedded in them like

plums in a plum pudding.

Thomson’s apparatusThomson’s apparatus

Marie and Pierre Curie discovered

that uranium, radium and polonium

were disintegrating over time

and emitting rays that exposed

photographic plates. Marie gave the

penetrating rays the name radiation.

Marie and Pierre Curie discovered

that uranium, radium and polonium

were disintegrating over time

and emitting rays that exposed

photographic plates. Marie gave the

penetrating rays the name radiation.

Ernest Rutherford discovered that

atoms are mostly empty space. He

thought that they have nuclei that,

in addition to containing the mass

and positive charge of the atom, are

orbited by the negatively charged

electrons. He also suggested the

existence of neutrons within the

nucleus.

Ernest Rutherford discovered that

atoms are mostly empty space. He

thought that they have nuclei that,

in addition to containing the mass

and positive charge of the atom, are

orbited by the negatively charged

electrons. He also suggested the

existence of neutrons within the

nucleus.


to pass directly through the gold foil. He commented on this experiment at his last public lecture as ‘it was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you’. (Note: 15 inches is approximately 38 cm.)

To explain his observations Rutherford suggested that the mass of an atom and the positive charge were concentrated in a small region in the centre of the atom which he called the nucleus. The majority of alpha particles missed the tiny gold nuclei and passed through the empty space without interference from the relatively small electrons.

Source of

alpha

particles

Lead shield Atoms of

gold foil

Nucleus

Alpha particles

Beam of

alpha

particles

Gold

foil

Following this experiment, Rutherford discounted Thomson’s solid ball model and proposed a new model of the atom known as the nuclear model of atomic structure. According to his model, most of the atom consists of empty space. The electrons orbit around a central nucleus that contains virtually all of the mass and positive charge of the atom.

Although Rutherford’s model solved the problem of the observed alpha particles’ deflections, it created another problem. If the positive charge was located in the centre of the atom and the negatively charged electrons were orbiting the nucleus, why were the electrons not immediately attracted to the nucleus, since opposite charges attract?

In 1920, Rutherford suggested that some nuclei also contained another particle with a neutral charge, which he named the neutron. However, it was not until 1932 that the neutron was actually discovered.

James Chadwick — discovery of the neutronIn 1932, the British scientist James Chadwick (1891–1974) bombarded a sample of beryllium with alpha particles and found that a ray was given off by the beryllium. The particles that made up the ray were not deflected by a magnetic or an electrical field, indicating that they carried no charges. Chadwick suggested that the alpha particles striking the beryllium foil displaced uncharged particles called neutrons from the nuclei of beryllium atoms. Further experiments showed that neutrons had almost the same mass as protons.

Rutherford’s alpha particle

experiment: (a) Apparatus

(b) Alpha-particle behaviour

near nuclei

Rutherford’s alpha particle

experiment: (a) Apparatus

(b) Alpha-particle behaviour

near nuclei

Gold foil was used for the alpha

particle experiment because gold is

so malleable that it can be beaten

out into thin foils without breaking.

Gold foil can be made so thin that a

strong light held behind it will shine

through.

Gold foil was used for the alpha

particle experiment because gold is

so malleable that it can be beaten

out into thin foils without breaking.

Gold foil can be made so thin that a

strong light held behind it will shine

through.

(a) (b)(a) (b)

In 1932 James Chadwick

discovered the neutron.

In 1932 James Chadwick

discovered the neutron.

James Chadwick was awarded the

Nobel Prize in physics in 1935 for

his discovery of the neutron three

years earlier.


3.

Revision questions

Rutherford’s nuclear model of the atom is often compared with the solar system, with the nucleus being analogous to the sun and the electrons revolving around it like planets.

(a) In what ways does Rutherford’s model resemble the arrangement of the solar system?

(b) In what ways is Rutherford’s model different from the arrangement of the solar system?

4. What is radioactivity? Why are some atoms radioactive?

5. How was radioactivity discovered?

6. How did the discovery of radioactivity upset ideas about atoms at the time?

7. What was the principal method used for the early exploration of the atom? Name a scientist who used this method and describe the nature of the work.

8. On what basis did Rutherford conclude that the nucleus was (a) small (b) positively charged?

9. Describe how J.J. Thomson arrived at his ‘plum-pudding’ model of the atom.

Frederick Soddy — isotopesIn 1910 the English physicist Frederick Soddy (1877–1956) made observations of the radioactive decay of thorium and found what he described as ‘two kinds of thorium’. Both forms of the thorium element contained atoms having 90 electrons and 90 protons, but one form of the element had 140 neutrons in the nucleus, and the other form of the element had 142 neutrons. Soddy referred to these ‘different forms of the same element’ as isotopes.

Soddy’s observation of isotopes,

elements with the same atomic

number but different mass

numbers, lead to the discovery

of the neutron.

Soddy’s observation of isotopes,

elements with the same atomic

number but different mass

numbers, lead to the discovery

of the neutron.

A mass spectrometer can be used

to analyse the isotopes of carbon,

nitrogen and strontium found in

African elephant tusks. This helps

to control the illegal ivory poaching

trade by identifying where the tusks

have come from.

A mass spectrometer can be used

to analyse the isotopes of carbon,

nitrogen and strontium found in

African elephant tusks. This helps

to control the illegal ivory poaching

trade by identifying where the tusks

have come from.


Identifying isotopes of elements with a mass spectrometerIn 1913 Thomson discovered that some elements could have atoms of different masses. Thomson’s equipment was developed by Francis Aston into a mass spectrometer, which could make comparisons of the relative masses of atoms. The figure below shows the main features of a mass spectrometer. A sample of the element to be analysed is injected as a gas into the ionisation chamber, where the atoms are ionised by bombardment with electrons produced by the hot filament. The positive ions formed are accelerated through an electric field and deflected in a magnetic field that forces the ions to travel in different paths.The curved paths of deflection depend on the mass-to-charge ratio of the ions. For a given charge (for example, singly charged ions), the heavier ions are harder to deflect and so travel in a wider curve. Ions corresponding to a fixed mass-to-charge ratio will be picked up by the ion collector and the ion current will be amplified and displayed.

The mass spectrometer provides us with information about:• the number of isotopes in a given sample of an element• the relative isotopic mass of each isotope• the percentage abundance of the isotopes.

The relative isotopic mass is the mass of a single isotope and is determined by comparing the mass of ions of the isotope to the value of a standard, 6

12C (which has been assigned a mass of 12 exactly). The figure below shows the mass spectrum for neon. The two peaks in the trace represent the two isotopes of neon. Both isotopes of neon have an atomic number of 10 (that is, the number of protons is 10), but one isotope has a relative isotopic mass of 20 and the other has a relative isotopic mass of 22. The atoms of neon-20 will have 10 protons and 10 neutrons, and atoms of neon-22 have 10 protons and 12 neutrons. The number of protons added to the number of neutrons in an atom is called its mass number.

Find out more about mass

spectrometry by going to

the website for this book

and clicking on the Mass

spectrometry weblink (see


Find out more about mass

spectrometry by going to

the website for this book

and clicking on the Mass

spectrometry weblink (see


A mass spectrometer. Different

types of atoms of the same element

can be separated on the basis of

their mass:charge ratio.

A mass spectrometer. Different

types of atoms of the same element

can be separated on the basis of

their mass:charge ratio.

Charged plates to

generate electric

field to accelerate

positive ions

Magnetic field causes positive

ions to deflect. The curve of the

deflection depends on the

mass (for a given charge) of

the atom.

To

pump

Sample

Hot filament

produces electron

beam

Ionisation

chamber

Positive ions

Ion

collector

Ion current detected

and displayed

Charged plates to

generate electric

field to accelerate

positive ions

Magnetic field causes positive

ions to deflect. The curve of the

deflection depends on the

mass (for a given charge) of

the atom.

To

pump

Sample

Hot filament

produces electron

beam

Ionisation

chamber

Positive ions

Ion

collector

Ion current detected

and displayed

A mass spectrometer will deflect

the path of the lightest element

the most. Similarly, it will deflect

the most highly charged particle

the farthest.

A mass spectrometer will deflect

the path of the lightest element

the most. Similarly, it will deflect

the most highly charged particle

the farthest.

A mass spectrum shows

information about the relative

isotopic mass and percentage

abundance of each isotope.

A mass spectrum shows

information about the relative

isotopic mass and percentage

abundance of each isotope.

Rela

tive

ab

un

dan

ce (

%)

Relative isotopic mass

10

18 19 20 21 22 23

50

90

100

Rela

tive

ab

un

dan

ce (

%)

Relative isotopic mass

10

18 19 20 21 22 23

50

90

100

A mass spectrum for

the element neon. The

element neon is made

up of two isotopes of

different proportions.

90% of the element neon

is made up of the isotope

with relative mass of

20, while the remaining

10% is the isotope with

relative mass of 22.

A mass spectrum for

the element neon. The

element neon is made

up of two isotopes of

different proportions.

90% of the element neon

is made up of the isotope

with relative mass of

20, while the remaining

10% is the isotope with

relative mass of 22.


Distinguishing the atoms: atomic numbers, mass numbers, relative isotopic masses and relative atomic massesThe structure of an atom is identified using the convention

Z

A E where E is the

symbol of the element, A is the mass number (number of protons + neutrons)

and Z is the atomic number (number of protons). For example, the two isotopes

of neon are represented as 1020Ne (atomic number, Z = 10; mass number, A = 20)

and 1022Ne (Z = 10; A = 22).

Most elements consist of a mixture of isotopes. The relative atomic mass

(Ar) of an element represents the average mass of one atom, taking into

consideration the number of isotopes of the element, their relative isotopic

mass (RIM) and their relative abundance. Using data from the mass spectrum,

one could calculate the relative atomic mass for an element by using the

following method of calculation:

For example, the two isotopes of neon, 2010 Ne and 22

10 Ne, have relative

isotopic masses of 20 and 22 respectively. Their relative abundances are 90.0%

and 10.0% respectively (see the graph on the previous page).

Ar(Ne)=

(20 90)+(22 10)

100

= 20.2

s s

Sample problem: 2.1 Calculating isotopic mass

Use the data in the table below to find the relative isotopic mass of 1737Cl. The

relative atomic mass, Ar , of chlorine is 35.45.

Isotope RIM % Abundance

35Cl17 34.97 75.80

3717 Cl unknown 24.20

Solution:

RIMof first isotope %abundance + RIMofs� second isotope %abundance

100=

s�

(34.97 75.80)+ RIM( Cl) 24.20

100= 35.45

265

1737s s

00.65+ 24.20 RIM( Cl)= 3545

24.20 RIM(

1737

17377

1737

Cl)= 3545 – 2650.65

RIM( Cl)=894.35

24.220

= 35.96

Z

AE A

Z

, = mass number

= atomic number

Z

AE A

Z

, = mass number

= atomic number

r

RIM abundan e

100�

s¤

cA

r

RIM abundan e

100�

s¤

cA

Ar =RIMof first isotope abundance +s� RRIM of second isotope abundance +.s� ...

100Ar =

RIMof first isotope abundance +s� RRIM of second isotope abundance +.s� ...

100

Ar (C1)Ar (C1)


Revision questions

10. Gallium has two isotopes. One isotope is 6931Ga and has a relative abundance

of 60.50%. The relative atomic mass of gallium is 69.70. Find the relative

isotopic mass of the other isotope.

11. Three isotopes of magnesium and their relative abundances are2412 Mg

(78.8%), 2512 Mg (10.2%), and 26

12 Mg (11.0%).

(a) Sketch on a graph the mass spectrum for magnesium.

(b) Calculate the relative atomic mass of magnesium.

12. What are the three kinds of particles in an atom? Where are they found in the

atom? List some of their physical properties.

13. (a) What evidence led to the prediction of the neutron?

(b) What are the atomic numbers of tellurium (Te) and iodine (I)?

(c) What are the relative atomic masses of Te and I?

(d) What are the numbers of protons, neutrons and electrons in the isotopes

of tellurium ( Te)12852

and iodine ( I)?12753

(e) Why does Te come before I in the periodic table?

14. Distinguish between the terms:

(a) neutron and proton

(b) mass number, atomic mass and atomic number.

15. What are isotopes? How were they discovered?

16. Why do isotopes of the one element show the same chemical behaviour even

though they differ in composition?

17. Work out how many neutrons, protons and electrons are present in an

atom of:

(a) 1123Na

(b) 1428Si

(c) oxygen-16

(d) 11 +H

(e) 1224 2+Mg

18. Silicon-containing ores have three isotopes: 92% silicon-28, 5% silicon-29

and 3% silicon-30.

(a) Draw the mass spectrum for silicon.

(b) What is the atomic number of silicon?

(c) What are the relative isotopic masses of the three silicon isotopes?

(d) What is the relative atomic mass of silicon?

Electrons in an atom

The continuing search to understand the structure of the atom is the concern

of both chemists and physicists. We have examined experimental evidence

that led to the idea of atoms possessing the fundamental subatomic particles:

electrons, protons and neutrons. Investigations also showed that protons and

neutrons make up the nucleus of an atom with electrons orbiting around it.

Increasing experimental evidence led to the present-day model of the atom.

Modern atomic theory is concerned primarily with electron arrangement

around the nucleus and has helped the understanding of the chemical


properties and behaviour of atoms. In chemical reactions, when atoms react, bonds are broken and atoms rearrange themselves to form new bonds. The formation of new bonds involves the redistribution of electrons.

Electron

(a) Thomson

model

(b) Rutherford

model

(c) Bohr model

Nucleus Orbit

(d) Quantum-

mechanical

model

Nucleus Electron

90% probability of

finding the electron

inside this line

Niels Bohr and the hydrogen spectrumIn 1913 Niels Bohr (1885–1962) proposed a new model of atomic structure based on a series of experiments involving the spectra produced by hydrogen

atoms.

The spectrum of atomic hydrogen

can be observed by passing an

electric current through hydrogen

gas at low pressure.

The visible hydrogen spectrum

consists of four lines: red, green, blue

and violet. There are also two nearby

lines in the ultraviolet region.

To explain the emission spectra

produced by hydrogen atoms, Bohr

hypothesised that if atoms emit only

discrete wavelengths, they must have

only discrete energies. He proposed

that electrons moved around the

nucleus of an atom in certain fixed

paths, called orbits. Each orbit

corresponded to a different energy

level. An electron could change its

energy by going from one energy

level to another, but could not

remain between energy levels. Themovement of an electron from one energy level to another is called a quantum jump. A specific quantity of energy (a photon) is associated with each quantum jump made by an electron.

According to Bohr, an electron could move from a lower energy level to a

higher energy level by absorbing energy, for example from a flame. When the

electron falls back to a lower level, energy is released as a consequence of

the quantum jump. The energy given out is the difference in energy between

the two energy levels and will be associated with a specific wavelength.

Each of these wavelengths corresponds to a coloured line in the emission

spectrum; some of these lines are in the visible part of the electromagnetic spectrum.

Models of the atom have changed

as chemists and physicists have

learned more about atomic

structure. (a) The ‘plum pudding’

arrangement of electrons

(b) Electrons orbit around the

nucleus like planets circle around

the sun (c) Electrons orbit around

the nucleus in different energy

levels (d) Electrons as clouds of

negative charge inside orbitals

Models of the atom have changed

as chemists and physicists have

learned more about atomic

structure. (a) The ‘plum pudding’

arrangement of electrons

(b) Electrons orbit around the

nucleus like planets circle around

the sun (c) Electrons orbit around

the nucleus in different energy

levels (d) Electrons as clouds of

negative charge inside orbitals

The Bohr model of the atom

suggested that electrons orbited

the nucleus in shells. Atoms

emit a quantum of energy in the

form of a photon of light if they

move from the ground state to

an excited state.

The Bohr model of the atom

suggested that electrons orbited

the nucleus in shells. Atoms

emit a quantum of energy in the

form of a photon of light if they

move from the ground state to

an excited state.

Niels Bohr proposed that the

electrons of an atom were restricted

to certain energy levels. Bohr won

the Nobel prize in physics in 1922

for his work on atomic spectra.

Niels Bohr proposed that the

electrons of an atom were restricted

to certain energy levels. Bohr won

the Nobel prize in physics in 1922

for his work on atomic spectra.

The Bohr model of an atom

Electrons would not be found in these ‘non-orbit’ areas

1st orbit

Nucleus

2nd orbit

3rd orbit

°

Electrons would not be found in these ‘non-orbit’ areas

1st orbit

Nucleus

2nd orbit

3rd orbit

°


Violet Blue Green Black Red

When the electrons in an atom are in the lowest possible energy levels, the atom is said to be in its ground state. If an electron moves to a higher energy level by absorbing energy, the atom is said to be in an excited state. An atom in an excited state will eventually return to its ground state and in doing so will

release energy.

Revision questions

19. (a) What does the emission spectrum of hydrogen look like?

(b) In 1913, Niels Bohr proposed a model for the hydrogen atom to explainthe emission spectrum of hydrogen. Explain how he accounted for thepattern of this emission spectrum.

20. Explain the relationship between the terms ‘ground state’ and ‘excited state’.

Erwin Schrödinger — the quantum-mechanical model of an atomAlthough Bohr’s model of an atom could account for the lines seen in the emission spectra of hydrogen, it did not explain and could not mathematically predict the lines seen in the emission spectra of the more complex atoms. Furthermore, it did not explain why electrons moving around the nucleus and emitting electromagnetic radiation did not fall into the nucleus of the atom, causing it to collapse.

In 1923 the French scientist Louis de Broglie (1892–1987) proposed, using Albert Einstein’s and Max Planck’s quantum theory, that electrons show both particle and wave behaviour. In 1926, the Austrian physicist Erwin Schrödinger (1887–1961) made use of the new quantum theory to refi ne Bohr’s model of the atom. He wrote and developed a wave equation describing the location and

Scientists have been able to

determine the elements present

in stars from an analysis of their

spectra. During a solar eclipse

in 1868, the light from the outer

layers of the sun was analysed and

spectral lines not corresponding to

any known element at the time were

observed. The lines were attributed

to a new element that was given

the name helium, derived from the

Greek word helios, meaning sun.

Helium was not identifi ed on

Earth until 1895, almost 30 years

later!

Scientists have been able to

determine the elements present

in stars from an analysis of their

spectra. During a solar eclipse

in 1868, the light from the outer

layers of the sun was analysed and

spectral lines not corresponding to

any known element at the time were

observed. The lines were attributed

to a new element that was given

the name helium, derived from the

Greek word helios, meaning sun.

Helium was not identifi ed on

Earth until 1895, almost 30 years

later!

The visible region of the emission

spectrum of hydrogen consists of

four lines. Each line corresponds to

a specifi c wavelength. This series of

lines is known as the Balmer series

for hydrogen.

The visible region of the emission

spectrum of hydrogen consists of

four lines. Each line corresponds to

a specifi c wavelength. This series of

lines is known as the Balmer series

for hydrogen.

Schrödinger developed the

quantum mechanical model

of the atom, which features

electrons arranged in shells,

subshells and orbitals within

an atom.

Schrödinger developed the

quantum mechanical model

of the atom, which features

electrons arranged in shells,

subshells and orbitals within

an atom.


energy of an electron in a hydrogen atom. From this, he developed the currently accepted quantum-mechanical model of the atom—a complex mathematical model based on particles such as electrons showing wave-like behaviour.

According to quantum mechanics, the electron is not considered as moving along a definite path. Instead, the electron is found in a region of space around a nucleus called an orbital. An orbital may be visualised as a blurry cloud of negative charge: the cloud is most dense where the probability of finding the electron is large and less dense where the probability of finding the electron is small.

1s1s

2s

In the quantum-mechanical model an electron around a nucleus may be visualised as a

cloud of negative charge. As in the Bohr model, the electron is attracted to the nucleus

by electrostatic forces and moves in such a way that its total energy has a specific

value. In (a) the charge cloud for the 1s electron in hydrogen is shown, and in (b) charge

clouds for the electrons in the 1s and 2s subshells are shown.

In the quantum-mechanical model of the atom:

1. The energy levels of electrons are designated by principal quantum numbers, n, and are assigned specific values: n = 1, 2, 3, 4, 5 and so forth. These principal quantum numbers may be referred to as shells and are also called the K, L, M and N shells respectively.

2. Within each shell, several different energy levels called subshells may be found. The number of subshells equals the shell number; for example, if the shell number is 2, there will be 2 subshells at that energy level. Each subshell corresponds to a different electron cloud shape. Subshells are represented by the letters s, p, d, f, and so on.

TABLE 2.2 Energy levels within shells of an atom

Shell

number

(n) Shell symbol

Number of

subshells

Subshell

symbol

Maximum

number of

electrons in

subshell

1 K 1 s 2

2 L 2 s 2

p 6

3 M 3 s 2

p 6

d 10

4 N 4 s 2

p 6

d 10

f 14

Erwin Schrödinger proposed a

charge cloud model of the atom.

Erwin Schrödinger proposed a

charge cloud model of the atom.


Electron configurationThe way in which electrons are arranged around the nucleus of an atom is

called the electron configuration of the atom.

Generally, the order of subshell filling is from lowest energy first:

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p . . .

1s

2s

3s

4s

5s

6s

7s7p

6p

5p

4p

3p

2p

Note: represents an orbital

Incre

asin

g e

nerg

y

6d

5d

4d

3d

5f

4f

Notice that the 4s subshell is filled before the 3d subshell, which is of a higher

energy than the 4s subshell. Likewise, the 4d subshell is higher in energy than

the 5s subshell, and so on.

Excited states

When an atom moves to a higher energy level than the ground state by

absorbing energy, its electron configuration changes. The outermost electron

moves to a higher energy level subshell. For example, neon has 2 electrons

in the first shell and 8 in the second shell; when the outermost electron

gains energy, 1s22s22p6 becomes 1s22s22p53s1. Note that once the order of

filling subshells has been determined, the subshells are written in increasing

numerical order, not the order of increasing energy. So 1s22s22p63s23p64s23d2

becomes 1s22s22p63s23p63d24s2.

The electron configuration for the elements hydrogen, sodium and scandium,

for example, can be written as:

H (Z = 1) 1s1

Na (Z = 11) 1s22s22p63s1

Sc (Z = 21) 1s22s22p63s23p64s23d1, or alternatively

1s22s22p63s23p63d14s2 to show that the last occupied shell is the fourth one;

nevertheless, the fourth shell starts to fill before the third shell is complete.

The group and period in which an element is found is easily read from the electron configuration, as shown in the example below:

group 2 (is the number of valence electrons in the last shell being filled)

Be (Z = 4) 1s22s2

period 2 (the highest shell number being filled)

group 18 (2 + 6)

Ar (Z = 18) 1s22s22p63s23p6

period 3

The order of filling subshells is

1s2 2s2 2p6 3s2 3p6 4s2 3d10.

The order of filling subshells is

1s2 2s2 2p6 3s2 3p6 4s2 3d10.

The energy levels of atomic orbitalsThe energy levels of atomic orbitals

The order of filling of subshells may

be found by following the arrows

from tail to head, starting with the

top arrow.

The order of filling of subshells may

be found by following the arrows

from tail to head, starting with the

top arrow.

#6d

1s

2s

3s

4s

5s

6s

7s

2p

3p

4p

5p

6p

7p

3d

4d

*5d

7d

*4f#5f

6f

7f

* and # In both these cases it is

found that the two subshells

fill at the same time.

#6d

1s

2s

3s

4s

5s

6s

7s

2p

3p

4p

5p

6p

7p

3d

4d

*5d

7d

*4f#5f

6f

7f

* and # In both these cases it is

found that the two subshells

fill at the same time.

To use an interactive periodic

table, go to the website for this

book and click on the Interactive

periodic table weblink (see


To use an interactive periodic

table, go to the website for this

book and click on the Interactive

periodic table weblink (see



The group number is found by using the total number of electrons in the last shell to count across the periodic table. The period is the number of the last shell occupied by electrons.

Sample problem: 2.2 Ground state electron confi guration for fl uorine

Find the ground state electron confi guration for a fl uorine atom.

Solution:

Fluorine has 9 electrons. According to the order of subshell fi lling, the 1s subshell in a fl uorine atom will fi ll fi rst, and will contain 2 of fl uorine’s 9 electrons. The next energy level is the 2s subshell — this will hold another two electrons. The 2p subshell can hold six electrons, but since only 5 electrons remain to be placed, one of the 2p subshell orbitals will be incomplete.

Fluorine’s electron confi guration may be written as follows:

1s22s22p5

Sample problem: 2.3 Ground state electron confi guration for potassium

Find the ground state electron confi guration for a potassium atom.

Solution:

Potassium has 19 electrons. Two electrons fi ll the 1s subshell, 2 electrons fi ll the 2s subshell, 6 electrons fi ll the 2p subshell, 2 electrons fi ll the 3s subshell and 6 electrons fi ll the 3p subshell. This means that 1 electron remains to be placed. The next subshell, 3d, is of a higher energy level than the 4s subshell. The remaining electron will be placed in the 4s subshell. The electron confi guration of potassium is:

1s22s22p63s23p64s1

Revision questions

21. Write the full electron confi guration of the following elements: lithium,

potassium, carbon, chlorine, argon and nitrogen.

22. The isotope 1532P is used in the treatment of leukaemia.

(a) Write the full electron confi guration of this isotope.

(b) How does it differ from 1531P?

23. An element X has confi guration 1s22s22p63s23p4.

(a) What group is it in?

(b) What period is it in?

(c) Give its name and symbol.

24. Name the elements with the following confi gurations, and state the group

number and period number in each case.

(a) 1s22s1

(b) 1s22s22p6

(c) 1s22s22p63s1

(d) 1s22s22p63s23p1

(e) 1s22s22p63s23p6

(f) 1s22s22p63s23p64s2


25. For each of the elements in the following two sets of atomic numbers, write the electron confi guration, and state which elements belong to the same group.(a) Z = 20, 12, 4, 9(b) Z = 5, 6, 8, 16.

26. Write the electron confi guration and element symbols for:(a) the fi rst three elements in group 18(b) the element in group 3, period 3(c) the element in group 5, period 2(d) the element in group 17, period 3.

Sample problem: 2.4 Ground state electron confi guration of an aluminium ion

Find the ground state electron confi guration of an aluminium ion, Al3+.

Solution:

An aluminium ion, Al3+, is an aluminium atom that has lost three electrons. Therefore, its electron confi guration is as follows:

1s22s

22p6

Sample problem: 2.5 Comment on the electron confi guration below

1s22s

22p63p

1

Solution:

The species contains 11 electrons. It represents an atom or ion in an excited state, because the fi nal electron would, in the ground state, be placed in a 3s

orbital.

Revision questions

27. Distinguish between the terms ‘shell’ and ‘subshell’.

28. Write the ground state electron confi guration for the fi rst 30 elements of the periodic table.

29. Which of the following electron confi gurations are ground state con-fi gurations and which are excited state confi gurations?(a) 1s22s22p63s1

(b) 1s22s22p63s23p63d1

(c) 1s22s22p63s23p63d54s2

30. A neutral magnesium atom has an electron confi guration of 1s22s22p63s16s1.(a) How can you tell that the atom is in an excited state?(b) Describe what would happen if the atom changed its electron

confi guration to the ground state.(c) Write the ground state electron confi guration for the magnesium ion

Mg2+.

31. Identify the following elements with ground state electron con fi gurations of:(a) 1s22s22p63s23p63d54s2

(b) 1s22s22p63s23p64s2


Chromium and copper — atypical electron configurationsA few elements have electron configurations that do not follow the usual pattern. Chromium (atomic number 23) and copper (atomic number 29), for example, may be expected to be written as follows:

chromium 1s22s22p63s23p63d44s2

copper 1s22s22p63s23p63d94s2.

The correct electron configurations are shown below:

chromium 1s22s22p63s23p63d54s1

copper 1s22s22p63s23p63d104s1.

These arrangements give chromium a half-filled d subshell and copper a filled d subshell. Filled subshells are more stable than half-filled subshells. However, half-filled subshells are more stable than other partly filled subshells.

Chromium has a half-filled d

subshell. Copper has a filled

d subshell. These unusual

electron configurations make

the elements more stable.

Chromium has a half-filled d

subshell. Copper has a filled

d subshell. These unusual

electron configurations make

the elements more stable.

The Australian Synchrotron is a

particle accelerator in Victoria that

is used to discover more about the

structure of the atom, including

bosons, quarks and hadrons.

The Australian Synchrotron is a

particle accelerator in Victoria that

is used to discover more about the

structure of the atom, including

bosons, quarks and hadrons.


SummaryThe changing models of the atom have been due largely to the application of the scientifi c method of investigating (formulating a hypothesis and testing it) and the development of technology capable of testing theories put forward by scientists.

Around 400 BC, Democritus proposed that matter is made up of atoms. He suggested that they are small, indivisible particles and that the differences between atoms resulted from differences in their size, shape and weight.In 1808, John Dalton proposed the fi rst modern atomic theory in which he stated that:1. All matter is composed of extremely small

particles called atoms.2. Atoms are indivisible and cannot be created or

destroyed.3. Atoms of the same element are identical and

have the same mass.4. Atoms are rearranged in chemical reactions.5. Atoms of different elements can combine with

one another in simple whole number ratios to form compounds.

In 1896, Antoine Becquerel, Pierre Curie and Marie Curie discovered radioactivity. This led to the notion that atoms are not indivisible, disputing a previous theory of Dalton’s.In 1897, J. J. Thomson discovered electrons by studying the defl ection of cathode rays towards the positively charged plates of his specially designed cathode-ray tube. He proposed the ‘plum-pudding’ theory of the atom which suggested that the atom is a sphere of positive charges throughout which electrons are embedded.In 1910, Frederick Soddy discovered different forms of the same element which he called isotopes.In 1911, Ernest Rutherford discovered the nucleus by bombarding very thin gold foil with alpha particles. He found that most of the positively charged particles went straight through the gold atoms while a small number bounced back. He suggested that the atoms had a nucleus containing positively charged particles (protons) with electrons orbiting around it.In 1932, James Chadwick discovered the neutron.

A mass spectrometer is an instrument that can separate isotopes of an element based on their mass-to-charge ratio. A mass spectrum of an element shows the number of isotopes present in an element, their relative isotopic mass and their proportion

•

−

−

−

−

−

−

−

•

(% abundance). The relative atomic mass of an element can be calculated using the method:

Ar RIM of first isotope abundan� s1

100[( % cce

RIM of second isotope abundance

)

( %

�

s ))� . . .]

where Ar is the relative atomic mass and RIM is the relative isotopic mass.Atoms are identifi ed by the general conventionA

Z, where Z is the atomic number (the number of

protons) and A is the mass number (the number of protons and neutrons).Niels Bohr re-emphasised Rutherford’s idea of electrons moving around the nucleus. His model, however, suggested that electrons orbited around the nucleus in defi nite paths of energy called quantum levels or shells. A quantum of energy or photon emitted is the result of an electron gaining enough energy to move to a higher energy level and, when falling back to its ground state position, emitting the photon. Bohr determined the different energy levels by using mathematical formulas which measured the wavelengths of the different energy levels.Erwin Schrödinger’s most important contribution to modern atomic theory was his development of the mathematical description of the paths electrons would most likely follow in their orbits around the nucleus. The formulas he developed formed the basis of the quantum mechanical model of the atom. In it he proposed that, instead of Bohr’s idea of the electrons following predetermined paths, they would be moving around in regions of space called orbitals. The quantum mechanical model has the following features:

Orbitals are of various shapes and are found within subshells which in turn are found in shells. Each shell has a different energy level, with the shell furthest away from the nucleus having the highest energy level and the one closest to the nucleus having the lowest energy level.Each energy shell has a principal quantum number, n. The fi rst four shells have n = 1, 2, 3 and 4; they are also called the K, L, M and N shells respectively.Subshells are energy levels found within shells. There are four different types of subshells — the s, p, d and f subshells.The maximum number of electrons the different subshells can hold are:

s subshell: two electrons p subshell: six electrons d subshell: ten electrons

f subshell: fourteen electrons.

•

•

•

−

−

−

−

Chapter review


Electrons fill shells and subshells of lowest energy first. The order of filling is:

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d

The ground state electron configuration of an atom refers to electrons in their lowest energy level. Any other configurations represent the atom in an excited state and in a higher energy level.

Multiple choice questions 1. Which of the following is incorrect?

A Ernest Rutherford was famous for his ‘gold-foil’ experiment.

B The cathode-ray tube was the basis of much experimental evidence for the initial development of modern atomic theory.

C Pierre and Marie Curie were the first scientists to discover radioactivity.

D The idea of atoms was proposed as far back as 400 BC by Democritus.

2. Which of the following ideas of John Dalton’s atomic theory is no longer true?A All matter is composed of tiny indestructible

particles called atoms.B Atoms of the same element are alike in every

way.C Atoms of different elements are different.D Atoms can combine together in small numbers

to form molecules. 3. J. J. Thomson’s basis for the ‘plum-pudding’ model

for atomic structure was that:A he discovered electrons by investigating

cathode rays in a specially designed cathode-ray tube

B he found that cathode rays travelled towards the positively charged anode, so the particles in the rays must be negatively charged

C he measured the amount of deflection of cathode rays attracted by a positively charged plate

D he postulated that an atom consists of a core of positive charges surrounded by negatively charged electrons.

4. Ernest Rutherford discovered a nucleus in atoms. What experimental evidence did he have?A Alpha particles were bombarded at the atoms

of a thin gold foil.B Most of the alpha particles passed directly

through the foil because gold atoms were mainly made of empty space.

C A small number of alpha particles bounced back because they hit the centre of the atom which was made up of a core of negatively charged particles.

−

•

D He proposed the nuclear model of the atom, which suggested that the atom has a small central core of protons surrounded by electrons orbiting around it.

5. Which of the following particle pairs has approximately the same mass?A a proton and an electronB a proton and a neutronC a neutron and an electronD an electron and a hydrogen atom

6. Three atoms, I, II and III, each have an atomic number of 12. Atom I has 12 neutrons, atom II has 13 neutrons and atom III has 14 neutrons. Which of the following sentences is correct?A Atoms I, II and III are allotropes of each other.B Atoms I, II and III are isotopes of the same

element.C Atom I is a neutral atom while atoms II and III

are cations of atom I.D Atom I is a neutral atom while atoms II and III

are anions of atom I. 7. A neutral atom of the isotope 6

13C would consist of:A 6 protons, 13 neutrons, 13 electronsB 0 protons, 13 neutrons, 13 electronsC 13 protons, 7 neutrons, 13 electronsD 6 protons, 7 neutrons, 6 electrons.

8. Which of the following features is not present in a mass spectrometer?A The gas sample is bombarded by alpha

particles.B The gas sample is ionised into positively

charged particles.C The positive ions are accelerated by an electric

field.D A magnetic field forces the particles to separate

in curved paths according to their mass-to-charge ratio.

9. Which of the following statements regarding subatomic particles is correct?A Protons are positively charged particles and

neutrons are negatively charged.B The relative masses of an electron, a proton and

a neutron are all about 1 unit.C In a neutral atom, the number of neutrons is

equal to the number of protons.D Isotopes of an element have the same number of

protons but a different number of neutrons.10. Bohr’s theory of the atom proposed that:

A electrons orbit the nucleus like planets move around the sun

B no more than two electrons are allowed in any energy level

C electrons move around the nucleus in fixed orbits, each of which has a different energy level

D energy shells have subshells which contain regions of space called orbitals.


11. The maximum number of electrons that can be placed in the shell n = 3 is:A 18 C 2B 8 D 32.

12. Which of these electron configurations represents an atom in an excited energy state?A 1s22s22p43s2

B 1s22s22p63s23p4

C 1s22s22p6

D 1s22s22p63s23p63d34s2

13. The electron configuration of 17

35 C1– is:A 1s22s22p63s23p53d1

B 1s22s22p63s23p44s2

C 1s22s22p63s23p6

D 1s22s22p63s23p43d14s1.14. The electron configuration 1s22s22p63s23p6 rep-

resents which of the following ions?A O2–

C Al3+

B S2– D Na+

15. The electron configuration of an atom X is 1s22s22p63s23p1. Which of the following formulae is most likely to be a compound formed with X?A XF2

B CaXC XCl3

D MgX2

16. The ground state electron configuration of a neutral atom with an atomic number of 19 is:A 1s22s22p63s23p63d1

B 1s22s22p63s23p7

C 1s22s22p63s23p64s1

D 1s22s22p63s23p53d14s1.

Review questions 1. (a) What did the alchemists contribute to the

growth of chemistry?(b) Would you consider the alchemists to be

scientists? Justify your answer. 2. Examine the points made by Dalton in his atomic

theory. Consider each statement, and record:(a) whether or not each suggestion still holds(b) which other scientists were able to contribute

to refining or changing each proposal, and in what way

(c) our current understanding of each statement. 3. From the results of Rutherford’s ‘gold-foil’

experiment, he suggested that the protons of the atoms in the metal must be concentrated in the

centre, or nucleus, of the atom, while the electrons are outside the nucleus. If this experiment were to be repeated, what results would have been

obtained if:(a) Thomson’s ‘plum-pudding’ model was correct(b) the electrons were concentrated in the nucleus

and the protons orbited the nucleus?

4. Complete the following table of the structural properties of atoms.

Element

Atomic

number

(Z )

Atomic

mass

(A)

Number of:

Protons Neutrons Electrons

1 0

nickel 59 28

gold 197 79

6 8

osmium 76 190

19 20

silver 47 107

silicon 16 14

mercury 80 120

5. Oxygen consists of three isotopes: 16O, 17O and 18O. Show how many different masses an oxygen molecule, O2, can have.

6. Naturally-occurring chromium consists of the following four isotopes:4.31% 50Cr (relative isotopic mass 49.946)83.76% 52Cr (relative isotopic mass 51.941)9.55% 53Cr (relative isotopic mass 52.941)2.38% 54Cr (relative isotopic mass 53.939).(a) Calculate the relative atomic mass of

chromium.(b) Explain the difference between ‘relative atomic

mass’ and ‘mass number’, selecting appropriate data from the above list.

7. A sample of carbon is found to contain two isotopes with relative isotopic masses of 12.00 and 13.00. If the relative atomic mass of carbon is 12.01, calculate the relative abundances of each of the isotopes.

8. Suppose that you are a research chemist and you have just discovered a new element. How would you identify its different isotopes and determine its relative atomic mass?

9. Explain the difference between:(a) a ‘shell’ and a ‘subshell’(b) an atomic ‘orbit’ and an ‘orbital’.

10. What is the maximum number of electrons that may be found in:(a) a 3p subshell(b) a 2s subshell(c) a 4d subshell(d) the third shell?

11. Write the electron configuration of each of the following in their ground states:(a) sodium atom(b) nitrogen atom(c) argon atom(d) iron atom(e) copper atom(f) calcium ion(g) chloride ion.

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