1

Atomic Structure &

the Periodic Table

Teacher Notes

Essential Revision for GCSE

Chemistry/Combined Science

These notes are intended for you to read and then summarise. Only through making

summaries (mind mapping, writing notes or creating flash cards) will learning become

more effective.

2

History of Atomic Models

In this historical section on models of the atom (pages 2-4), the bold writing is what is stated in the syllabus – i.e. the bold writing is what you actually need to learn.

New experimental evidence may lead to a scientific model being changed or replaced. Before the discovery of the electron, atoms were thought to be tiny spheres that could not be divided.

1808. Before the discovery of the electron, atoms were

thought to be tiny spheres that could not be divided.

John Dalton was a schoolteacher, a meteorologist, and

an expert on colour blindness, but he is best known for

his theory of atomism, the foundation of our modern

concept of the atom. Dalton held that there are as many

different kinds of atoms as there are chemical

elements.

It was the English physicist J. J. Thomson who first

discovered electrons in 1897; small negatively

charged particles in an atom.

The discovery of the electron led to the plum

pudding model of the atom. The plum pudding

model suggested that the atom is a ball of positive

charge with negative electrons embedded in it.

Ernest Rutherford

In 1911 Rutherford, Geiger and Marsden

conclude that that the mass of an atom was

concentrated at the centre (nucleus) and that

the nucleus was charged. This nuclear model

replaced the plum pudding model.

3

Rutherford’s Experiment in more detail

Investigating the structure of the atom In 1911 Geiger and Marsden were tasked by Rutherford to

probe the structure of matter by firing heavy (relative to

electrons) positively charged alpha particles at a thin gold sheet.

Results

Most of the alpha particles passed straight through without being deflected

A small number were deflected

An even smaller number were deflected back the way they came Explanation

Most of the alpha particles passed straight through the gold foil showing that the gaps between the nuclei are mostly empty space and that most of the mass of an atom is concentrated in a very small space in the middle of the atom called the nucleus

The ones that were deflected experienced an electric force of repulsion showed that the nucleus was charged (having the same charge as alpha particles)

and since only a small number of alpha particles were deflected this showed that the nucleus was small

In 1913 Bohr argued that electrons move around a nucleus in

separate energy levels called orbitals. Niels Bohr adapted

the nuclear model by suggesting that electrons orbit the

nucleus at specific distances. The theoretical calculations of

Bohr agreed with experimental observations.

4

Later experiments led to the idea that the positive charge of any nucleus could be subdivided into a whole number of smaller particles, each particle having the same amount of positive charge. The name proton was given to these particles.

History of the Periodic Table Before the discovery of protons, neutrons and electrons, scientists attempted to classify the elements by arranging them in order of their atomic weights (what we now call “relative atomic mass”. The early periodic tables were incomplete and some elements were placed in inappropriate groups if the strict order of atomic weights was followed.

Elements with properties predicted by Mendeleev were discovered and filled the gaps. Knowledge of isotopes made it possible to explain why the order based on atomic weights was not always correct. Today we understand that the elements in the periodic table are arranged in order of atomic (proton) number and so that elements with similar properties are in columns, known as groups. The table is called a periodic table because similar properties occur at regular intervals.

In 1932, James Chadwick discovered neutrons, and isotopes

were identified.

The experimental work of James Chadwick provided the evidence to show the existence of neutrons within the nucleus. This was about 20 years after the nucleus became an accepted scientific idea.

In 1869 Russian chemist Dimitri Mendeleev overcame some of the problems by leaving gaps for elements that he thought had not been discovered. Also, he wasn’t afraid of swapping the places of some elements (which were originally in order of atomic weight) so that elements with similar chemical properties were located under each other.

5

Elements in the same group in the periodic table have the same number of electrons in their outer shell (outer electrons) and it is this that gives them similar chemical properties

The Simplest Atom

The simplest atom is hydrogen. 1 electron moves around the nucleus.

The hydrogen atom

It is the outer electrons in atoms that control chemical behaviour. Chemists spend more time studying electrons than any other particle.

The Oxygen atom

8 electrons move around the nucleus. Two electrons are in the first shell (orbit). Six

electrons are in the second shell.

6

Why is the nucleus positively charged?

The nucleus is made of two types of particle:

protons and neutrons

Protons are positively charged. They make the whole nucleus positively charged.

Neutrons are particles which are similar in size to protons but they do not have a charge.

Particles found in the atom

Atoms are very small, having a radius of about 0.1 nm (1 x 10-10 m).

The radius of a nucleus is less than 1/10 000 of that of the atom (about 1 x 10-14 m).

All atoms are electrically neutral.

We believe this because we don’t normally get an electric shock when

we touch things

Therefore, every atom must contain as many protons as electrons

1/1836 (negligible)

7

Arrangement of electrons in atoms

Electrons move around the nucleus in orbits or shells

The electron shell nearest the nucleus can hold a maximum of 2 electrons

The second shell can hold a maximum of 8 electrons

The third shell can hold as many 18 electrons but you’ll never come across these atoms at

GCSE. For the first 20 atoms in the periodic table, the maximum number of electrons the

third shell can hold never exceeds 8

How to find out how many protons and neutrons are in the nucleus

Each element in the Periodic Table is given two numbers. The top number is called the

mass number. The bottom number is called the atomic number.

Example: Sodium

ATOMIC NUMBER: This is the number of protons in the nucleus of the atom.

Since, all atoms are neutral, the atomic number also tells you

how many electrons are in the atom.

MASS NUMBER: This is the number of protons + neutrons in the

nucleus of the atom.

8

• For example, sodium has an atomic number of 11. This means a sodium atom has 11

protons (and, therefore, 11 electrons).

• Sodium has a mass number of 23. This means it has 23 protons and neutrons in its

nucleus.

• Therefore, sodium must have 11 protons and (23-11) 12 neutrons in its nucleus.

Number of neutrons = mass number – atomic number

The Sodium Atom

The sodium atom has 3 electron shells. That's why it's in Period 3 of the Periodic Table.

The sodium atom has 1 electron in its outer shell. That's why it's in Group 1 of the

periodic table.

The group number tells you the number of electrons in the outer shell of the atom.

The period number (row) = number of electron shells in the atom.

9

Isotopes An element may have two or more different isotopes.

Isotopes are atoms of the same element, with the same number of protons but

different numbers of neutrons.

EXAMPLE: Chlorine has two isotopes:

Since each isotope has the same number of protons, they also have the same number of electrons

(Remember, atoms are neutral).

On Earth, about 75% of all chlorine is chlorine-35 and 25% is chlorine-37. This means that, the "average"

mass number (i.e. the relative atomic mass) of chlorine is:

(75 x 35) + (25 x 37)

------------------------------- = 35.5 (relative atomic mass is basically the average of all the mass numbers)

100

Important: Isotopes have identical chemical properties because they have identical electronic structures. In other words, having different numbers of neutrons has no effect on the chemical properties of isotopes

10

Electronic Structure

Instead of drawing the electron shells of an atom, it is quicker to write the electronic

structure of the atom.

The electronic structure of sodium (atomic number = 11) is:

Example Question

Write down the electronic structure of aluminium:

The atomic number of aluminium is 13. This means an aluminium atom has 13 protons

and, therefore, 13 electrons.

Aluminium is in Period 3 so there are 3 electron shells containing electrons

Aluminium is in Group 3 so there are 3 electrons in the outer shell

Therefore, the electronic structure of aluminium is:

Al (2, 8, 3)

A good way of checking is to see if the numbers add up to 13.

11

The Modern Periodic Table

The Periodic Table (first 30 elements shown only)

An element is a substance made of only one type of atom.

The Periodic Table shows all the elements, arranged in order of ascending atomic number. There are over 100 elements. More than ¾ are metals The columns in the Periodic Table are called GROUPS. The rows in the Periodic Table are called PERIODS. Some Groups in the Periodic table are given "family names": Group Number Family Name ___________________________________________

1 The Alkali Metals 7 The Halogens 8 The Noble Gases ___________________________________________ There is one family of elements that are not in any Group. They are called the Transition Metals. They are located between Group 2 and Group 3.

12

Elements in the same group have similar chemical properties. This means they behave in a similar way, when they react with other substances. The reason why elements in the same group have similar chemical properties is because they all have the same number of electrons in their outer shells

Metals and Non-metals Metals are found on the left of the Periodic Table. Non-metals are found on the right of the Periodic Table.

Useful Observation Non-metal oxides (for example sulphur dioxide) are acidic.

Metal oxides (for example copper oxide) are basic (a base is any chemical that neutralises an acid to make salt + water).

Oxidation

Most elements react with oxygen. You are expected to be able to write balanced equations for the oxidation of familiar elements. Examples: sulphur + oxygen sulphur dioxide S + O2 SO2 This is a non-metal oxide so it’s acidic

magnesium + oxygen magnesium oxide 2Mg + O2 2MgO This is a metal oxide so it’s a base

Example Question: Which of the following substances are acidic?

(a) iron oxide (b) sodium oxide (c) sulphur dioxide (d) carbon dioxide (e) chlorine oxide (f) calcium oxide ANSWER: (c), (d) and (e) are acidic because they are non-metal oxides.

13

Group Chemistry

Each group in the periodic table is distinct and contains elements with similar chemical properties. In this topic you focused on three Groups in the periodic table:

Group 0 (The Noble Gases) are a group of unreactive gases

Group 1 (The Alkali Metals) are a group of reactive metals

Group 7 (The Halogens) are a group of reactive non-metals

Group 0 - "The Noble Gases"

The Noble Gases are a group of unreactive gaseous elements. They are placed in Group 0 because the number of electrons these atoms need to gain or lose to achieve a full outer shell is zero. The first 4 elements in Group 0 are:

Helium, He Neon, Ne Argon, Ar

Krypton, Kr They are unreactive gases that exist as single atoms. This is because they have a complete outer electron shell. They do not need to lose or gain any electrons because they already have a stable electronic structure. For some reason, full outer shells are very stable.

It seems that all the other elements in the periodic table are reactive because they

desire a Noble Gas electronic structure, i.e. there is a tendency for atoms to acquire a

full outer shell of electrons.

Any element not in Group 0 can gain a full outer shell by undergoing a chemical

reaction.

14

Physical Properties Although the noble gases have no interesting chemical properties (they’re unreactive) they do have some notable physical properties. For instance, as you go down Group 0, the boiling point increases:

Uses of the Noble Gases Noble Gas use ___________________________________________________________

helium balloons

neon, argon and krypton electric lighting ___________________________________________________________

The reason why the boiling points

increase is all to do with the size of

the atoms.

There seems to be a correlation

between how many electrons

molecules have and the strength of

attraction between any two

molecules.

In the case of the Noble gases, as you

go down the group, the atoms get

bigger (they have more electrons) and

the strength of the attraction between

the atoms increases. Therefore the

boiling points increase.

15

Group 1 - "The Alkali Metals"

The Alkali Metals themselves are not alkali. If you dabbed a piece of universal indicator paper on the metal, nothing would happen. The reason why they are called the “Alkali Metals” is because they all react vigorously with water forming an alkaline solution. The first 3 elements in Group 1 are:

Lithium, Li Sodium, Na

Potassium, K They are very reactive metals. They react rapidly with oxygen to form the metal oxide:

4Li + O2 2Li2O 4Na + O2 2Na2O 4K + O2 2K2O The alkali metals react vigorously with water to form a hydroxide (this is an alkaline solution) and hydrogen gas. 2Li + 2H2O 2LiOH(aq) + H2 lithium + water lithium hydroxide + hydrogen 2Na + 2H2O 2NaOH(aq) + H2 sodium + water sodium hydroxide + hydrogen 2K + 2H2O 2KOH(aq) + H2 potassium + water potassium hydroxide + hydrogen alkaline solutions always contain OH-(aq) ions

The Alkali Metals are stored in oil. This is because oil protects the metals from oxidation.

Therefore, all these metal oxides are

bases – they neutralise acids to

produce salt + water

16

Comparing the reactivity of Group 1 Metals with water Alkali Metal o b s e r v a t i o n ___________________________________________________________ lithium reacts quickly & dissolves, fizzes (gas given off) sodium reacts and fizzes more vigorously & dissolves, potassium reacts violently, hydrogen gas catches fire. Dissolves ___________________________________________________________ Reaction with chlorine gas They burn quickly in chlorine with a bright flame. Potassium burns fastest and sodium next. Lithium is the least reactive. White solids are left behind. lithium + chlorine lithium chloride {slowest} 2Li + Cl2 2LiCl

sodium + chlorine sodium chloride 2Na + Cl2 2NaCl

potassium + chlorine potassium chloride {fastest} 2K + Cl2 2KCl

white solids

Why are Group 1 metals reactive?

They have only 1 electron in their outer shell. During a chemical reaction a Group 1

metal loses its outer electron easily to form a positively charged ion which has a stable full outer shell of electrons.

Some Symbols

Li = lithium atom (reactive) Li+ = lithium ion (stable)

Na = sodium atom (v. reactive) Na+ = sodium ion (stable)

K = potassium atom (extremely reactive) K+ = potassium ion (stable)

17

Why do Group 1 metals get more reactive as you go down the Group?

2, 1 2, 8, 1 2, 8, 8, 1

As you down Group 1, the alkali metals become more reactive This is because the outer shell electron gets further away from the nucleus. Therefore the attraction from the nucleus is weaker and the atom loses its electron more easily.

Also, as you down Group 1 the atoms have more shells of electrons. These inner

electron shells shield the outer electron from the electrostatic attraction of the nucleus and again make it easier for the outer electron to be lost

18

Bromine: poisonous

reddish-brown liquid.

Forms a brown vapour

Group 7 - "The Halogens" The halogens are reactive non-metals. Each atom needs to gain 1 electron to achieve a full outer shell.

You only need to learn about chlorine, bromine and iodine.

Halogen Molecular formula State at room temperature

Colour

chlorine Cl2 gas Greenish/yellow

bromine Br2 liquid brown

iodine I2 solid Silvery grey

Common properties of Halogens

They are non-metals

Low melting and boiling points

Poor conductors of heat and electricity

Have coloured vapours

Their molecules each contain two atoms – i.e. they are diatomic (e.g. Cl2) They are reactive because they only need to gain 1 more electron to complete their outer electron shells.

They react vigorously with Group 1 metals: sodium + chlorine sodium chloride {fastest reaction} sodium + bromine sodium bromide sodium + iodine sodium iodide {slowest reaction}

As you go down Group 7, the halogens become less reactive

Chlorine: poisonous

greenish-yellow gas Iodine: harmful silvery-grey

solid – readily turns into a

purple vapour when heated

Therefore, they have simple

covalent structures

19

Displacement Reactions of Halogens

These are simple test tube reactions which can be used to decide the order of reactivity of the halogens.

Method 1. Put some potassium chloride solution in a test tube, potassium bromide solution in

another test tube and potassium iodide solution in a third test tube (By the way, it doesn’t have to be a potassium salt – it could be any metal salt (e.g. a sodium salt), as long as the halide is soluble

2. Then, using a dropping pipette, squirt some chlorine water (chlorine dissolved in water) to each of the test tubes and see whether there is a chemical reaction. You’ll know whether there is a reaction because the solution will darken. Record your observations

3. Prepare three more test tubes of potassium chloride solution, potassium bromide solution and potassium iodide solution.

4. Using a dropping pipette, squirt some bromine water (bromine dissolved in water) to each of the test tubes. Once again, make a note of any chemical reactions by seeing whether or not the solution darkens.

5. Prepare three more test tubes of potassium chloride solution, potassium bromide solution and potassium iodide solution.

6. Using a dropping pipette, squirt some iodine solution (iodine dissolved in water) to each of the test tubes. Once again, make a note of any chemical reactions by seeing whether or not the solution darkens.

Please note: in lessons you may have used white cavity tiles instead of test tubes

20

Results

Potassium chloride Potassium bromide Potassium iodide Chlorine water

Bromine water No reaction

Iodine water No reaction No reaction

Conclusions

Chlorine is more reactive than bromine because there was a reaction between chlorine

and potassium bromide – the solution went darker. So, chlorine displaced bromine from

potassium bromide:

chlorine + potassium bromide → potassium chloride + bromine

Cl2(aq) + 2KBr(aq) → 2KCl(aq) + Br2(aq)

Chlorine is also more reactive than iodine because chlorine reacted with potassium

iodide – the solution darkens. So, chlorine displaced iodine from potassium iodide:

chlorine + potassium iodide → potassium chloride + iodine

Cl2(aq) + 2KI(aq) → 2KCl(aq) + I2(aq)

Bromine is more reactive than iodine because there was a reaction between bromine and

potassium bromide – the solution darkens. Therefore, bromine displaced iodine from

potassium iodide:

bromine + potassium iodide → potassium bromide + iodine

Br2(aq) + 2KI(aq) → 2KBr(aq) + I2(aq)

In summary

Iodine didn’t displace anything so it must be the least reactive halogen that was tested. Bromine only managed to displace iodine so is the 2nd most reactive halogen tested. Chlorine displaced both bromine and iodine so is the most reactive halogen tested.

21

Ionic Equations

Take a look at this displacement reaction:

chlorine + potassium bromide → potassium chloride + bromine

Cl2(aq) + 2KBr(aq) → 2KCl(aq) + Br2(aq)

Potassium bromide is a salt. It’s made of ions rather than atoms. That’s because when

potassium reacted with bromine to make potassium bromide, the potassium atom

(which is in Group 1) lost its outer electron and in the process gained a full outer shell.

However, when potassium lost its outer electron, it became positively charged (as there

are now less electrons than protons).

Similarly, when bromine (which is in Group 7) reacted with potassium, it gained an

electron (from potassium) ending up with stable full outer shell. However, when

bromine gained an electron it ended up with an overall negative charge.

Therefore, if we are going to be accurate, the formula of potassium bromide should really

be written as:

K+Br-

In this displacement reaction, the potassium bromide is dissolved. This means the

particles are separated and all mixed up with water molecules.

Therefore, the best way of writing the formula of potassium bromide solution is:

K+(aq) + Br-(aq)

In this displacement reaction another salt is produced – potassium chloride. Just like

potassium bromide, this is also made of dissolved ions.

Therefore, the best way of writing the formula of potassium chloride solution is:

K+(aq) + Cl-(aq)

OK, let’s write the whole equation out again, but this time instead of writing “KBr(aq)” or

“KCl(aq)” we’ll use the more accurate formula:

22

Cl2(aq) + 2K+(aq) + 2Br-

(aq) → 2K+(aq) + 2Cl-(aq) + Br2(aq)

Although it looks really weird, writing it out like this makes it easier to notice something

strange. Have you noticed that the potassium ions (K+) haven’t reacted? They were

there at the start of the reaction and they’re still there at the end of the reaction! But a

chemical equation is meant to show just the chemical changes taking place. There really

is no point writing the formulae of particles that are not involved in a chemical reaction!

There is a special name for ions like the potassium ions – they’re called spectator ions. A

spectator ion is an ion that isn’t involved in the main reaction.

An ionic equation is a chemical equation with the spectator ions removed.

Here’s the same equation, with the spectator ions (the potassium ions) left out:

Cl2(aq) + 2Br-(aq) → 2Cl-(aq) + Br2(aq)

An ionic equation is really a simplified equation. You’re just focusing on what reacts. In

this case, a chlorine molecule reacts with 2 bromide ions and you produce 2 chloride ions

and a bromine molecule – chlorine displaces the bromine because chlorine is more

reactive. The reaction had nothing to do with potassium. We could just as easily have

used sodium chloride, sodium bromide and sodium iodide instead – the same reaction

would have happened but this time, sodium ions would be the spectator ions left out of

the ionic equation.

By the way, the other two displacement reactions (written on page 19) can also be

written as ionic equations:

Cl2(aq) + 2I-(aq) → 2Cl-(aq) + I2(aq) chlorine displaces iodine

Br2(aq) + 2I-(aq) → 2Br-

(aq) + I2(aq) bromine displaces iodine

23

Why do Group 7 Elements get less reactive as you go down the group?

This is the complete opposite trend to Group 1 – which get more reactive. Why is this? It’s because the Group 7 elements achieve stability by doing the complete opposite to Group 1. Whereas Group 1 atoms lose their outer electron to achieve a full (and stable) outer shell, Group 7 atoms need to gain 1 electron to complete their outer shell. As you go down Group 7, it becomes steadily more difficult to gain an electron because:

(i) As the atoms get bigger, the nucleus gets further away from the electron it’s trying to gain so the attraction for this electron is weaker

(ii) As the number of shells increase there is greater shielding (or repulsion) of the electron it’s trying to gain. Therefore, the attraction from the nucleus is weaker

Therefore, the reactivity of the halogens decreases as you go down the group.

Fluorine is more reactive than chlorine Chlorine is more reactive than bromine Bromine is more reactive than iodine

24

Trends in Boiling Points Just like the Noble gases, the halogen molecules get bigger as you down Group 7.

Therefore, the molecules have more electrons and this causes the intermolecular forces

to become stronger. Consequently the boiling points increase.

Uses of Group 7 elements and compounds

Name of substance Use ___________________________________________________________ chlorine kills bacteria in water, raw material for PVC plastic

iodine solution antiseptic widely used in operating theatres

Two Important points about the chemistry of metals and non-metals

Since metals are on the left hand side of the periodic table (i.e. Groups 1, 2 and 3), they only have a few outer electrons (just 1, 2 or 3!) Therefore, during chemical reactions it’s always easier for metal atoms to lose electrons (and become positively charged ions) to achieve a full outer shell. And since non-metals are found on the right hand side of the periodic table (e.g. Groups 5, 6 and 7), they already have quite a lot of electrons in their outer shells. Therefore, during chemical reactions non-metal atoms usually gain electrons (and become negatively charged ions) to complete their outer shells.

25

The Transition Metals (Separate Sciences only)

The Transition Elements are located between Group 2 and Group 3 in the Periodic Table.

Iron is the cheapest and most useful transition metal because:

(a) Iron ore (haematite) is plentiful (b) Iron is relatively cheap to extract using well established blast furnace technologies (c) Iron is useful because it is a hard and strong

Properties of Transition Metals and their compounds • The metals are hard, strong and have high densities • The metals have high melting points • The metals are shiny • They are good catalysts (e.g. iron is a catalyst used in the manufacture of ammonia) • Their compounds are brightly coloured {for example, copper sulfate is blue}