general summary gcse

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Summary of the structure of the Periodic Table

1a. The basic structure of the Periodic Table

Only the top portion of the periodic table is shown above (full version)

See thenotes 1. to 4.in the full Periodic Table at the end of this page.

The idea of the Periodic Table is to arrange the elements in a way that enables chemists to understand patterns in the properties

of elements, but some reminders first.

An ATOM is the smallest particle of a substance which can have its own characteristic properties, BUTatomsare built up of even

more fundamental sub-atomic particles - the electron, proton and neutron and the structure of an atom ultimately determines its properties.
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An ELEMENT is a pure substance made up of only one type of atom , 92 of the elements in the Periodic Table (part of which is

shown above) naturally occur, from hydrogen H (element 1) to uranium U (element 92).

Note that each element has a unique symbol which is a single capital letter like H orU or a capital letter + small letter e.g. cobalt Co,

chlorine Cl or sodium Na.

The majority of elements are readily divided into two types with common characteristic physical and chemical properties.

o Most elements on the left areMETALS and their typical properties are described in section 2a.e.g. elements 3 to 4 (lithium to

beryllium), elements 11 to 13 (sodium to aluminium), elements 19 to 31 (potassium to gallium), elements 37 to 50 (rubidium to tin).

o The elements on the right areNON-METALS and their typical properties are described in section 2b.e.g. elements 1 to 2

(hydrogen to helium), elements 5 to 10 (boron to neon), elements 15 to 18 (phosphorus to argon), elements 35 to 36 (bromine to

krypton).

o BUT a few elements are referred to asSEMI-METALS which have mixed metal/non-metal character and not so easy to

classify, see section 2c.They occur in a diagonal band (down and L to R) e.g. silicon (14Si), germanium (32Ge), arsenic (33As)

and tellurium (52Te).

The elements are laid out in order of Atomic (proton) Number* (*see atomic structure page).

o Originally they were laid out in order of 'relative atomic mass' (the old term was 'atomic weight').

o This is not correct for some elements now that we know their detailedatomic structure(detailed GCSE notes) in terms of

protons, neutrons and electrons, and of course, their chemical and physical properties.

o For example: Argon (at. no. 18, electrons 2,8,8) has a relative atomic mass of 40. Potassium (at. no. 19, electrons 2,8,8,1) has a

relative atomic mass of 39. Argon, in terms of its physical, chemical and electronic properties is clearly a Noble Gas in Group 8

(0). Likewise, potassium is clearly an Alkali Metal in Group 1.

Many of the similarities and differences in the properties of elements can be explained by the electronic structure of the

atoms (electron configuration = electron arrangement in shells or energy levels, so watch out the varying phrases used!).

The idea of the Periodic Table is to arrange the elements in a way that enables chemists to understand patterns in the properties of

elements.

The main structural features of the periodic table are ...
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o to produce columns of similar elements called Groups.

They are usually similar chemically and physically BUT there are often important trends in physical properties and

chemical reactivity up/down a group.

They are similar elements because they have the same outer electron structure - same number of outer electrons.

o The resulting complete horizontal rows are called Periods and usually consist of a range of elements of different character.

There are important trends from left to right across a period e.g. the most important overall change is

from metallic ==> non-metallic element character.

Certain 'horizontal blocks' of elements within a period, which have specific chemical features in common, may be known

as a particularblock orseries e.g. from 21Sc to 30Zn are called the 1st Transition Metal Series within period 4.

o The ideas of Group and Period are totally connected with electron structure (see below).

1b. Electronic structure and the Periodic Table

Which electron arrangements are stable and which are not?

o The maximum electrons allowed in the shells or electronic energy levels up to atomic number 20 are:

1st shell 2, 2nd shell 8, 3rd shell 8, the 19th and 20th electrons go into shell 4 (this represents limit GCSE students

need to know about electron arrangements -details for GCSE/IGCSE on the Atomic Structure page).
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The chlorine cannot accept another electron to form a Cl2-

ion, because its electron structure would not be that of

a stable noble gas arrangement.

o The electron arrangements of the first 20 elements are shown below.

o NOTE: In the most modern periodic table notation Groups 3-7 and 0 are numbered Group 3 to 18.

Detailed Advanced Level Chemistry Notes on the electronic structure of atoms
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1c. Electronic structure and the arrangement of elements in the Periodic Table

All substances are made up of one or more of the different types of atoms we call elements and the elements identity is solely

determined by the atomic number of protons. Hydrogen, 1, H, the simplest element atom, does not readily fit into any group.

A Group is a vertical column of like elements e.g.Group 1 The Alkali Metals(for full GCSE notes on Li, Na, K etc.),Group 7/17 The

Halogens(for full GCSE notes on F, Cl, Br, I etc.) andGroup 0/8/18 The Noble Gases(for full GCSE notes on He, Ne, Ar etc.).

Apart from hydrogen (doesn't really fit in any group), and helium (*), the Group number equals the number of electrons in the outer

shell and the number of electron shells used equals the Period number, e.g. chlorine's electron arrangement is 2.8.7, the second

element down in Group 7 on period 3. So after helium, elements in the same group have the same outer electron structure.

o * Although helium can't have 8 outer electrons like the rest of Group 0, its outer shell of 2 electrons is complete according to the

electron shell rules, just like neon and argon etc.

The elements in a group tend to have similar physical and chemical properties because of their similar outer shell electron

structure.

A Period is a horizontalrow of elements with a variety of properties, changing from very metallic elements on the left to non-metallic

elements on the right. A period starts when the next electron goes into the next available main energy level or shell (Group 1 alkali

Metals). The period ends when the main energy level is full (Group 0 or 8 Noble Gases).

All the elements on the same period use the same number of principal electron shells, and this equals the period number (e.g.

sodium's electron arrangement 2,8,1, the first element in Period 3).

The first element in a period is when the next electron goes into the next available electron shell or energy level (i.e. 1 electron in the outer

shell, after H it is the Group 1 Alkali Metals like sodium 2.8.1).
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The last element in a period is when the outer shell is full resulting in a very unreactive element, the Group 0 Noble Gases e.g.

argon 2.8.8. The next electron for the next element goes into the next highest level (shell) available, and so starts the next period with a

group 1 element again, periodicity - a very similar element every so often - but governed by the electron rules.

So in terms of electrons ....

o Period 1 is elements 1-2, H (1) to He (2)

o Period 2 is elements 3-10, Li (2,1) to Ne (2,8)

o Period 3 is elements 11-18, Na (2.8.1) to Ar (2.8.8)

o Period 4 is elements 19-36, starts with K (2,8,8,1) and Ca (2,8,8,2) and finishes with the Noble Gas Kr (2,8,18,8).

o Note that the number of shells containing electrons is equal to the period number.

The similarities (e.g. same Group) or differences (e.g. across a period) of the properties of the elements can be explained by the electronic

structure of the atoms.

From Period 4 onwards the length of a period significantly increases because it includes horizontal series of similar metals with their own

characteristic physical and chemical properties e.g.The 1st Transition Metals Series(detailed GCSE notes on Fe, Cr, Cu etc.)

Advanced Level Electron configurations, spectroscopy, hydrogen spectrum, ionisation energies

1d. More on patterns in the Periodic Table
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end is the very reactive non-metal liquid bromine, and the period finishes with the very unreactive gas krypton. Even more

complicated pattern!

o From left to right across a period the bonding in chlorides or oxides changes from ionic (with metals e.g. Na+Cl

-, Mg

2+O

2-to

covalent (with non-metals e.g. HCl, SO2).

o From left to right across a period the oxides change from alkaline/basic (with metals e.g. Na2O) to acidic (with non-metals e.g.

SO2).More on this in Group 6/16 Oxygen and oxides.

o Note on electron arrangements:

Except for boron, most non-metals have at least four electrons in the out shell.

Except for the noble gases, the more electrons in the outer shell the more non-metallic and the more reactive the element.

The most reactive non-metals only need to share/gain one or two electrons.

The most reactive metals only have 1 or 2 electrons in the outer shell which tend to be easily lost to form the metal ion in

reaction.

The most reactive metals have a low number of outer valency shell electrons (


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1e. Valency and formula patterns in the Periodic Table

The valency, or, combining power of an element is related to the elements position in the Periodic Table .

o For Groups 1 to 8, the group number gives the maximum valency possible and equals the number of outer

electrons (well, nearly always!).

o For many compounds, this rule works fine: e.g. for chlorine valency 1 and oxygen valency 2, you can deduce the following

formulae for valencies of 1 to 7 across the periods for Group 1 to 7 compounds (at least up to a point!) e.g. for period 3

chlorides: NaCl, MgCl2, AlCl3, SiCl4, PCl5, SCl6, then Cl itself and Ar can't combine with other elements.

oxides: Na2O, MgO, Al2O3, SiO2, P2O5, SO3, Cl2O7 and Ar can't combine with other elements.

BUT things are not always so simple!

o Na to Si no problem! great! In fact, apart from N, O, F (which have valency restrictions NOT for GCSE though!) you can usually

make a reasonable prediction of the maximum valency compound for all of the elements in Groups 1 to 7.

o However there are lots of other compounds where the element's valency is less than its group number and there is even a

pattern of decreasing valency from Group 4 to Group 7 (as well as the pattern of increasing valency mentioned above, see

table below and the decreasing pattern for hydrides which is important for GCSE level).

o e.g. in Group 4, C forms CO (nasty!) but CO2 is more stable, Pb forms PbCl2 which is much more stable than PbCl4.

o Xenon forms XeF8 and XO4 using its maximum valency of 8! and that got somebody a Nobel Prize in Chemistry! (and in scrabble

too?)

o Tabulated below are some formulae you are likely to come across in your GCSE or equivalent course in bold, but others you are

unlikely to come across are included because they fit in with general formula patterns.

o The valency of hydrogen is 1 (hydrides), oxygen 2 (most oxides) and chlorine is usually 1 (most chlorides).

o The expected-theoretical formulae for the hydride, chloride and oxide for element X of valency 1 to 5 are given below and

examples of all these formulae can be found in the Period 2-3 table further down.

o


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LiCl

Li2O

BeCl2

BeO

BCl3

B2O3

CCl4

CO2

NCl3

several

Cl2O

O2

ClF

F2O

-

-

Period 3

Na

-

NaCl

Na2O

Mg

-

MgCl2

MgO

B

BH3

AlCl3

Al2O3

Si

SiH4

SiCl4

SiO2

P

PH3

PCl3, PCl5

P2O3, P2O5

S

H2S

-

SO2, SO3

Cl

HCl

Cl2

Cl2O

Ar

-

-

-

Advanced Level Chemistry Notes on Period 2 survey Li to Ne

Advanced Level Chemistry Notes on Period 3 survey Na to Ar

2. Comparing Physical and Chemical Properties of Elements
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2a. Typical Properties of Metallic Elements

Physical properties of metals

Usually high melting points and boiling points so all solid bar one (exceptions like mercury the only liquid metal at room temperature

and theAlkali Metals[GCSE notes]have untypical low melting points).

Often very good conductors of heat and electricity. This is due to the mobility of the free moving electrons in thestructure of a metal.

Most have a high density (exceptions like the Alkali Metals have untypical low densities, the first three Li, Na and K float on water before

the 'fizzing'!).

Their appearance is always 'shiny' (usually silvery, except for copper and gold)

Usually quite strong materials (exceptions like the Alkali Metals which are atypically very soft, and metals like lead are relatively soft

too) They are easily beaten into shape (malleable) or drawn into wire (ductile) of varying strength, from very weak sodium to very strong

iron).

Solids sonorous, they ring or resonate to produce a note when struck.

Chemical Properties of metals

They tend to form basic oxides that react with acids to form salts (if the oxide is soluble in water it forms an alkali of pH > 7,

universal indicator blue or violet). Most metals react with acids to form a salt and hydrogen. (see metal reactions:reactivityand metal-acid

reactions/equations[1]and[2]with answers).

Metals readily form positive ions in compounds by losing electrons e.g.

o sodium Na - e-==> Na

+, magnesium Mg - 2e

-==> Mg

2+or aluminium Al - 3e

-=> Al

3+

Theiroxides and chlorides are usually ionic* in terms of chemical bonding. e.g.
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o magnesium oxide, MgO orMg2+

O2-

, sodium oxide Na2O or(Na+)2O

2-,

and aluminium oxide Al2O3 or(Al3+

)2(O2-

)3

*At least at GCSE level, but there are some chloride exceptions at Advanced level such as FeCl3 and AlCl3.

Reactivity of Metals Notesand Metal Extraction Notes

2b. Typical Properties of Non-metallic Elements

Physical properties of non-metals

They usually have low melting points and boiling points and so can be gases, liquids or solids(exceptions like silicon, and

carbon as diamond or graphite, see GCSE notes).

Usually poor conductors of heat and electricity (exceptions like carbon in the form of graphite).

N0n-metals generally have a low density.

The appearance can be quite varied but tend to be dull if solid.

Often weak materials e.g. soft or brittle solids (exceptions like silicon, and carbon as diamond, which are very hard and strong)

When solid they are not easily beaten into shape or drawn into wire, the solids tend to be too brittle.
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Solid non-metals are not usually sonorous, e.g. they do not usually resonate or ring with sound, like when a piece of metal is

struck.

Chemical properties of non-metals

They form acidic oxides when burned in air or oxygen, these react with alkalis to form salts, if soluble in water they form acid

solutions of pH chloride: Cl2 + 2e-==> 2Cl

-(more simply Cl + e

-==> Cl

-typical of Group 7 Halogens)

o oxygen ==> oxide: O2 + 4e-==> 2O

2-(more simply O + 2e

-==> O

2-typical of Group 6 elements)

The oxides and chlorides, when combined with other non-metals are always covalent in terms of chemical bonding.o e.g. waterH2O(l), methane CH4(g), sulphur dioxide SO2(g) and hydrogen chloride HCl(g)

The oxides and chlorides, when combined with metals tend to be ionic in terms of chemical bonding e.g.

o sodium chloride, NaCl orNa+Cl

-, magnesium chloride MgCl2 orMg

2+(Cl

-)2 ,

and magnesium oxide, MgO orMg2+

O2-

2c. The Properties of Semi-metals or Metalloids
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A very tricky topic, only the basic idea should be dealt with at KS3/GCSE level.

Gp

3/13

Gp

4/14

Gp

5/15

Gp

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BASIC IDEA: A narrow diagonal band of elements can show both metallic and non-metallic physical or

chemical properties and are referred to as 'semi-metals' or 'metalloids'. Although most tend to be nearer

being a metal or a non-metal, they do represent the point elements change from metal to non-metal as

you move from left to right across the Periodic Table BUT please read the notes below carefully!

B C N O

To me boron, B, is clearly a non-metal, showing no real metallic character and I'm not sure why it is sometimes

shown as a semi-metal on some periodic tables? and is very different in character to metallic aluminium below it

in the same group. Boron's oxide is acidic only, and the solid element consists of a non-conducting giant covalent

structure, both classic non-metallic properties. Carbon, C, is also clearly a non-metal, its oxide is acidic and in

the form of diamond, it is a non-electrical conducting 3D giant covalent structure. However, in the form of

graphite, it has a layered 2D giant covalent structure that does allow electricity to conduct through the layers.

(more details)

Al Si P S

Physically and chemically aluminium, Al, is very much a metal, but the oxide/hydroxide reacts with both acids

(metallic) and alkalis (acidic) to form salts showing dual character. Silicon is mainly non-metallic

charactere.g. the oxide is acidic but, although the solid element has a giant covalent structure, it shows slight

electrical conducting properties (semi-conductor), especially when doped with other elements and used in

computer chip technology. To me, neither are true semi-metals.

Ga Ge As Se

Germanium, Ge, is considered as a true semi-metal (metalloid). Like silicon, germanium is a semi-conductorand used in electronic technology. Its oxide/hydroxide react with both acids/alkalis showing dual metal/non-metal

character. Arsenic, As, is also a true metalloid with oxides/hydroxides that react both with acids/ and alkalis to

form salts and the element exists in two allotropic* crystalline forms. One form is less dense, non-conducting and

covalent in structure (non-metal) and the other is more dense and weakly electrical conducting (metallic) and

used in transistors. Selenium, Se, is also a semi-conductorwith metallic and non-metallic properties and is
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used in photo-electric cells (solar cells) and xerography (photocopying). (*Allotropesare different physical forms

of the same element in the same physical state.)

In Sn Sb Te

Arsenic, As, (like antimony in the same group), is also a true semi-metal (metalloid) with oxides/hydroxidesthat react both with acids/ and alkalis to form salts and the element exists in two allotropic* crystalline forms (non-

metallic and metallic). Tellurium, Te, is also a semi-conductorwith metallic and non-metallic properties. Both

As and Te are used in electronic devices.

3. Links to three selected Data-Graphs of selected physical properties of elements

Links to the first 'experimental' editions of these new web pages are below. They are of more use to Advanced Level students studying

'Periodicity', but they are a source of useful data. There are also summaries of data for Group 1 Alkali Metals, Group 2 Alkaline Earth Metals,

Group 7 Halogens, Group 0 Noble Gases and the 1st series of Transition Metals.

Elements 1-20 covering Periods 1-3 and start of Period 4


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4. A brief Summary of some Groups & Series of elements of the Periodic Table

with links to more detailed GCSE notes where necessary

Group 1 Alkali Metals

The very reactiveGroup 1 The Alkali Metals [GCSE notes] have low density (some float on water).

They readily react with non-metals to form ionic compounds e.g. NaCl orNa+

Cl-

, Li2O or(Li+

)2O2-

. These are colourless crystals or white solids, soluble in water to give colourless solutions (usually pH 7 for their

salts, pH 13-14 for the oxides because MOH alkali formed).

The metals react rapidly, maybe violently, with water to form alkaline hydroxides and hydrogen gas.

Alkali metal atoms have one outer electron, which is readily lost to form a stable single positive ion M+.

Down the group, the metals get more reactive, and the melting points and boiling points decrease.
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Detailed advanced Level Chemistry Notes on Group 1 and Group 2 Metals

Group 2 Alkaline Earth Metals

Group 2 are the 2nd group of metals (sometimes called "Alkaline Earth Metals").

They are not quite so reactive as the Alkali Metals for the same period.

They have two outer electrons and readily lose them to form the M2+

ion.

This ion occurs in the ionic compounds they readily form with non-metals like the Group 7 Halogens or oxygen and sulphur from Group 6

e.g. MgCl2 orCaO.

Detailed advanced Level Chemistry Notes on Group 1 and Group 2 Metals
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Group 3 (Group 13 in modern notation)

Group 3/13 contains the metal Aluminium (seeGCSE metal extraction notesand uses of metals:GCSE comparison with transition

metalsortitanium/steel/alloy uses GCSE notes.

Advanced Level Chemistry Notes on Group 3/13 Introduction - Boron & Aluminium


Group 4/14 contains the non-metal carbon - which forms lots of compounds with hydrogen formed in oil (seeGCSE Oil Products notes).

The structure of the allotropes of carbon and the structure and properties of silicon dioxide-silica-SiO2 (GCSE notes ondiamond,

graphite and silica) are important.

Advanced Level Chemistry Notes on Group 4/14 Introduction - Carbon & Silicon - semi-metals e.g. Ge
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Group 5/15 contains the non-metal nitrogen - important element in natural and manmade artificial fertilisers (seeGCSE notes on

ammonia and nitric acid). Nitrogen forms 79% (4/5th's) of air.

Advanced Level Chemistry Notes on Group 5/15 Introduction - Nitrogen & Phosphorus


Group 6/16 are a Group of non-metallic elements, the first 2 are O oxygen and below it S sulphur.

They have 6 outer electrons and readily gain 2 electrons to form an X2-

ion in ionic compounds

o e.g. they form ionic compounds with metallic elements e.g. magnesium oxide MgO and sodium sulphide Na2S,

or written ionically: Mg2+

O2-

and (Na+)2S

2-.

They form covalent small molecule compounds with other non-metallic elements e.g. H2O orCS2.

The top element in the group is oxygen, a most important element.

o Made by green plants in photosynthesis.
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NO2 nitrogen dioxide is strongly acidic in water). There is no way of simply predicting this kind of behaviour from periodic

table patterns!

Sulphuris an important element used in theGCSE notes on the manufacture of sulphuric acid.

o Sulphur or its compounds in oil burn to form the acidic polluting gas sulphur dioxide, one of the causes of acid rain (see Oil

Product Notes).

Advanced Level Chemistry Notes on Group 6/16 Introduction - Oxygen & Sulfur

Group 7 The Halogens (Group 17 in modern notation)

TheGroup 7 Halogens[GCSE notes] are coloured non-metals with low melting points and

boiling points.

They are brittle when solid e.g. iodine and poor conductors of heat and electricity when liquid or solid.

Halogens exist as molecules of pairs of atoms, X2(diatomic molecules), form ionic salts with metals e.g. KBrorMgCl2, but

form covalent molecular compounds with other non-metallic elements e.g. HCl, CBr4.

The halide ions, X-, are formed by halogen atoms, with 7 outer electrons, gaining 1 electron to form a stable noble gas electron structure.

Down the group the melting points and boiling points increase and the reactivity decreases.

Sodium chloride is a very important raw material from which hydrogen, chlorine and sodium hydroxide can be manufactured by

electrolysis.
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These products have many uses and are important in the manufacture of other useful compounds ranging from bleaches, hydrochloric

acid and plastics etc.

Advanced Level Chemistry Notes on Group 7/17 The Halogens

Group 0 The Noble Gases (Group 18 in modern notation)

TheGroup 0/18 Noble Gases[GCSE notes] are colourless non-metals with very low melting and boiling

points (they are all gases at room temperature).

They exist as individual atoms (NOT diatomic molecules) and are very unreactive chemically due to their

very stable full outer shell electron arrangements.

Helium has a very low density and so is used in balloons and airships.

Their lack of chemical reactivity makes them useful to provide an 'inert' atmosphere to prevent oxidation e.g. argon in filament bulbs and in

arc welding.

Advanced Level Chemistry Notes on Group 0/18 The Noble Gases
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The 1st Transition Metal Series (Scandium to zinc)

The ten horizontal elements Sc to Zn are called the 1st series ofTransition Metal Elements[GCSE notes] e.g.

iron and copper.

These elements in the central blocks of the periodic table are typical metals - good conductors of heat and

electricity and can be bent or hammered into shape (malleable) and they can be drawn into wire (ductile).

However, compared to the group 1 Alkali Metals, they have higher melting points (except mercury - a liquid atroom temperature); they are harder, tougher and stronger; they are much less reactive and so do not react (corrode) as quickly with

oxygen or water.

These properties make them useful structural materials (e.g. steel) and were things need to be good conductors e.g. copper electrical

wiring or steel radiators.

Most transition metals form coloured compounds (e.g. blue copper salt solutions) and are used in pottery glazes, stained glass and

weathered copper roofs turn green!

Many transition metals e.g. iron and platinum are used as catalysts. Cast iron is hard and used as man-hole covers. Steel is an

alloy* based on iron and used for car bodies. *alloy means a metal mixed with at least one other element.

see alsoMetal Extraction(detailed GCSE notes) and more on metal uses on theExtra Industrial Chemistry- detailed GCSE notes -

use index of sub-pages.

Detailed Advanced Level Chemistry Notes on the 3d-block of elements and Transition Metals
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5. Snippets of the past and continuing history of the Periodic Table

5a. The early classification of Antoine Lavoisier of 1789

Antoine Lavoisier's 1789 classification of substances into four 'element' groups

acid-making elements gas-like elements metallic elements earthy elements

sulphur light cobalt, mercury, tin lime (calcium oxide)

phosphorus caloric (heat) copper, nickel, iron, magnesia (magnesium oxide)

charcoal (carbon) oxygen gold, lead, silver, zinc barytes (barium sulphate)

azote (nitrogen) manganese, tungsten argilla (aluminium oxide)

hydrogen platina (platinum) silex (silicon dioxide)


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The understanding that an element as a unique atomic 'building block' which could not be split into simpler substances and a compound is

a chemical combination of two or more elements were not at all understood at the time of Lavoisier.

However, Lavoisier was the first to define an element in the correct 'chemical sense' as a substance that could not be divided into simpler

substances.

'light' and 'caloric' (heat), were considered 'substances' and the last 'scientific' vestige of the elements of 'earth, fire, air and water' which

had there conceptual origin in the Greek civilisation of 2300-2800 years ago.

However, Lavoisier was correct on a few things e.g. the elements sulphur, phosphorus and carbon and correctly described their oxides as

acidic e.g. dissolved in water turned litmus turns red.

Many metallic elements, were correctly identified though I doubt if they were pure though!

What he described as the 'earthy elements' are of course compounds, a chemical combination of a metal plus oxygen or sulfur (both in

case of barium).

He didn't have very high temperature smelting technology, or a reactive metal from electrolysis (came in about 1806 onwards)' to

'separate' the elements in some way e.g. he couldn't extract a reactive metal! In other words, at this time, the wrong 'classification' was

due to a lack of chemical technology as much as lack of knowledge.

5b. The 1829 work of Johann Dbereiner

Johann Dbereiner noted that certain elements seemed to occur as 'triads' of similar elements e.g.

o (i) lithium, sodium and potassium


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o (ii) calcium, strontium and barium

o (iii) chlorine, bromine and iodine

Dbereinerwas amongst the first scientists to recognise the 'group' idea of chemically very similar elements.

Three groups he 'recognised' were (i) Group 1 Alkali Metals, (ii) Group 2 Alkaline Earth Metals, (iii) Group 7 Halogens.

5c. The work of John Newlands 1864

Newlands' Octaves (his 'Periodic Table' of 1864)

H Li Ga B C N O

F Na Mg Al Si P S

Cl K Ca Cr Ti Mn Fe

Co, Ni Cu Zn Y In As Se

Br Rb Sr Ce, La Zr Di, Mo Ro, Ru

Pd Ag Cd U Sn Sb Te

I Cs Ba, V Ta W Nb Au


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Pt, Ir Tl Pb Th Hg Bi Cs

Newlands recognised that every 7 elements, the 8th seemed to be very similar to the 1st of the previous 7 when laid out in a 'periodic'

manner and he was one of the first scientist to derive a 'Periodic Table' from the available knowledge.

e.g. his 'table' consists of almost completely genuine elements (Di was a mix of two elements), classified roughly into groups of simi lar

elements and a real recognition of 'periodicity'

He also recognised that the 'groups' had more than 3 elements (not just 'triads'), and was correct to mix up metals and non-metals in same

group e.g. in 5th column there is carbon, silicon, tin (Sn) from what we know call Group 4. However, indium is in group 3 but Ti, Zr have a

valency of 4, like Group 4 elements and do form part of vertical column in what we know call the Transition Metal series

Other correct 'patterns' if not precise are recognisable in terms of themodern Periodic Tablee.g. half of column 2 is Group 1, half of

column 3 is Group 2, half of column 5 is Group 4, half of column 6 is Group 5, half of column 7 is Group 6. If we put his column 1 as

column 7, it would seem a better match of today!

Although none of his vertical column groups match completely but the basic pattern of the modern periodic table was emerging. However

column's 1 and 7 do seem particularly mixed up compared to the modern periodic table!

5d. Dmitri Mendeleev's Periodic Table of 1869
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Mendeleev laid out all the known elements in order of 'atomic weight' (what we know call relative atomic mass) except for several

examples like tellurium (Te) and iodine (I) whose order he reversed because chemically they seemed to be in the wrong vertical column!

Smart thinking!

With an increased number of known elements, groups becoming more clearly defined, and he used a double column approach which is

NOT incorrect, i.e. a sort of group xA and xB classification. This is due to the 'insert' of transition metals, some of whom show chemical

similarities to the vertical 'groups', but this is needed to be understood for GCSE or A level!

However, his 'presentation' was sufficiently accurate, and Mendeleev was sufficiently confident to predict missing elements and their

properties * e.g. germanium (which he called eka-silicon, below Si and above Sn in Group IV and Mendeleev is deservedly called the

'father of themodern Periodic Table'.
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5e. The full modern version of the Periodic Table

Pdmetal groups

horizontal blocks of Transition Metal Series (Periods 4 to 7)

metal ==> non-metal groups

Gp1 Gp2 Gp3 Gp4 Gp5 Gp6 Gp7 Gp0

1 1H Note: H does not readily fit into any group which are the vertical columns 2He

2 3Li 4Be The full Modern Periodic Table of Elements

ZSymbol, z = atomic or proton number

5B 6C 7N 8O 9F 10Ne

3 11Na 12Mg 13Al 14Si 15P 16S 17Cl 18Ar

4 19K 20Ca 21Sc 22Ti 23V 24Cr 25Mn 26Fe 27Co 28Ni 29Cu 30Zn 31Ga 32Ge 33As 34Se 35Br 36Kr

5 37Rb 38Sr 39Y 40Zr 41Nb 42Mo 43Tc 44Ru 45Rh 46Pd 47Ag 48Cd 49In 50Sn 51Sb 52Te 53I 54Xe

655

Cs56

Ba *57-7172

Hf73

Ta74

W75

Re76

Os77

Ir78

Pt79

Au80

Hg81

Tl82

Pb83

Bi84

Po85

At86

Rn

7 87Fr 88Ra *89-103 104Rf 105Db 106Sg 107Bh 108Hs 109Mt 110Ds 111Rg 112Cp 113? 114? 115? 116? 117? 118?

Gp 1 Alkali Metals

Gp 2 Alkaline Earth Metals

Gp 7/17 Halogens

Gp 0/18 Noble Gases

Take note of the four points on the

right

*57La 58Ce 59Pr 60Nd 61Pm 62Sm 63Eu 64Gd 65Tb 66Dy 67Ho 68Er 69Tm 70Yb 71Lu

*89Ac 90Th 91Pa 92U 93Np 94Pu 95Am 96Cm 97Bk 98Cf 99Es 100Fm 101Md 102No 103Lr

1. Using 0 to denote the Group number of the Noble Gases is historic i.e. when its valency was

considered zero since no compounds were known. However, from 1961 stable compounds of

xenon have been synthesised exhibiting up to the maximum possible valency of 8 e.g. in XeO4.

2. Because of the horizontal series of elements e.g. like the Sc to Zn block (10

elements), Groups 3 to 7 & 0 can also be numbered as Groups 13 to 18 to fit in with the

maximum number of vertical columns of elements in periods 4 and 5 (18 elements per period).

3. This means that 21Sc to 30Zn can be now considered as the top elements in the vertical

Groups 3 to 12.


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4. I'm afraid this can make things confusing, but there it is, classification is still in progress! and

GCSE students can, as far as I can judge, ignore points 2 and 3.

5. Advanced Level Periodic Table Notes

With are knowledge of atomic structure, the modern Periodic Table is now laid out in order of atomic (proton) number and is directly linked

to theelectronic structure of elements.

Due to isotopic masses, the relative atomic mass does go 'up/down' occasionally (there is no obvious 'nuclear' rule that accounts for this,

at least at GCSE/GCE level!). BUT chemically Te is like S and Se etc. and I is like Cl and Br etc. and this is now backed up by modern

knowledge ofelectron structure.

We now know the electronic structure of elements and can understand sub-levels and the 'rules' in electron structure (seeatomicstructure page) e.g. 2 in shell 1 (period 1, 2 elements H to He), 8 in shell 2 (period 2, 8 elements Li to Ne), there is a sub-level which

allows an extra 10 elements (the transition metals) in period 4 (18 elements, K to Kr). this also explains the sorting out of Mendeleev's A

and B double columns in a group (but that's for much more advanced chemistry!). The periods are complete now that we know about

Noble Gases.

The use and function of the Periodic Table will never cease! Newly 'man-made' elements are being synthesised. In the 1940's the

research team developing the materials required to produce the first atomic bombs dropped on Hiroshima and Nagasaki realised that

'trans-uranium' elements were being formed in nuclear reactions (seeradioactivity-nuclear reactionspage). From element 93 to 111 are

now known, but sometimes just a few atoms from a cyclotron experiment and all are highly t radioactive due to unstable nuclei but thestructure of the bottom part of the periodic table will continue to grow and grow! Physicists are hoping to eventually make some 'nuclear

stable' super-heavy metallic elements around atomic number 150?
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6. Where do we get the elements from?

The ultimate origin of all elements is the nuclear reactions that go on when stars are formed from inter-stellar dust and gas forming a huge

combined mass due to gravity, and then 'chunks' of a star cool down to form planets. All the elements from atomic numbers 1-92 (H-U)

naturally occur on Earth, though some are very unstable-radioactive and decay to form more nuclear stable elements.

Everything around you is made up of the elements of the periodic table, BUT most are chemically combined with other elements in the

form of many naturally occurring compounds e.g.

o hydrogen and oxygen in water, sodium and chlorine in sodium chloride ('common salt'), iron, oxygen and carbon as iron

carbonate, carbon and oxygen as carbon dioxide etc. etc.!

Therefore, most elements can only be obtained by some kind ofchemical process to separate or extract an element from a compound

e.g.

o Less reactive metals are obtained by reduction of their oxides with carbon and more reactive metals are extracted by electrolysis

of their chlorides or oxides (seeGCSE notes on Metal Extraction)

o Non-metals are obtained by a variety of means e.g. chlorine is obtained by electrolysis of sodium chloride solution (seeGroup 7

The Halogens GCSE Notes).

However some elements never occur as compounds or they occur in their elemental form as well as in compounds e.g.

o The Group 0 Noble Gases are so unreactive they are only present in the atmosphere as individual atoms. Since air is a mixture,

these gases are separated from air by a physical method of separation by distillation of liquified air. The elements oxygen and

nitrogen are obtained from air at the same time, which is far more convenient than trying to get them from compounds like oxides

and nitrates etc.
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o Gold/platinum is are the least reactive metals and are usually found 'native' as the yellow/silver elemental metal.

o Relatively unreactive metals like copper and silver, can also be found in their elemental form in mineral deposits as well as in

metal ores containing compounds like copper carbonate, copper sulphide and silver sulphide.

o The non-metal sulphur is found combined with oxygen and a metal in compounds known as sulphates, but it can occur as

relatively pure sulphur in yellow mineral beds of the element.

-

WHAT ARE ALLOTROPES?

Oxygen atoms usually form 'stable' O2 oxygen molecules (also called dioxygen), BUT they can form an unstable molecule O3 ozone (also

called trioxygen). The mass of the oxygen atoms in each of the molecules is mainly 16 (99.8%), and about 0.2% of two other stable isotopes of

masses 17 and 18. Whatever isotope or isotopes make up the molecule, it doesn't affect the molecular structure or the respective

chemistry of the O2 or O3 molecules.

However, what sometimes confuses the issue is the fact that oxygen O2 and ozone O3 are examples of allotropes.

Allotropes are defined as different forms of the same element in the same physical state.

The different physical allotropic forms arise from different arrangements of the atoms and molecules of the element and in the case of solids,

different crystalline allotropes.

They are usually chemically similar but always physically different in some way e.g.


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O2 (oxygen, dioxygen) and O3 (ozone, trioxygen) are both gases but have different densities, boiling points etc.

Graphite, diamond and buckminsterfullerene are all solid allotropes of the element carbon and have significantly different physical and in

some ways chemical properties! (details on bonding page)

Rhombic and monoclinic sulphur have different geometrical crystal structures, that is different ways of packing the sulphur atoms (which

are actually both made up of different packing arrangements ofS8 ring molecules). They have different solubilities and melting points. There is

also a 3rd unstable allotrope of sulfur called plastic sulphurmade by pouring boiling molten sulphur into cold water which forms a black plastic

material consisting of chains of sulphur atoms -S-S-S-S-S- etc..

It doesn't matter which isotopes make up the structure of any of an element's allotropes described above, so to summarise by one example ...

oxygen-16, 17 or 18 are isotopes of oxygen with different nuclear structures due to different numbers of neutrons,

and O2 and O3 are different molecular structures of the same element in the same physical state and are called allotropes irrespective of

the isotopes that make up the molecules.
http://www.docbrown.info/page04/4_72bond4.htm#Large%20Covalent%20Molecules%20and%20their%20Propertieshttp://www.docbrown.info/page04/4_72bond4.htm#Large%20Covalent%20Molecules%20and%20their%20Propertieshttp://www.docbrown.info/page04/4_72bond4.htm#Large%20Covalent%20Molecules%20and%20their%20Propertieshttp://www.docbrown.info/page04/4_72bond4.htm#Large%20Covalent%20Molecules%20and%20their%20Properties

general summary gcse

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