|
Doc
Brown's Chemistry Introduction to the Periodic
Table
A summary survey of the Periodic
Table for revision KS4 Science IGCSE/O level/GCSE
Chemistry Information Study Notes for revising for AQA GCSE Science, Edexcel
360Science/IGCSE Chemistry & OCR 21stC Science, OCR Gateway Science
(revise courses equal to US grades 9-10)
Advanced
Level Periodic Table Notes - use index * EMAIL
query?comment
1.
Summary
of the structure of the Periodic Table
1a.
The basic structure of the Periodic Table * 1b.
Electronic structure and the Periodic Table
1c.
Electronic structure and the arrangement of elements in the
Periodic Table
1d.
More on patterns in the Periodic Table * 1e.
Valency and formula patterns in the
Periodic Table
2. Comparing Physical and
Chemical Properties of Elements
2a. Typical
Properties of Metallic
Elements * 2b. Typical
Properties of Non-metallic
Elements
2c.
The Properties of Semi-metals or
Metalloids
3.
Links to three
selected Data-Graphs of selected physical properties of elements
4. A
brief Summary of some Groups & Series of elements of the Periodic Table
5.
Snippets of the past and continuing history
of the Periodic Table
5a. Early
classification by Antoine Lavoisier of 1789 * 5b. The 1829 work of
Johann Döbereiner
5c.
The work of John Newlands 1864 * 5d. Dmitri Mendeleev's
Periodic Table of 1869
5e.
The
full modern
version of the Periodic Table
6. Where do we get the
elements from?
KEYWORDS-phrases for this page:
Electron arrangements and the Periodic Table * Gp
1 Alkali Metals
* Gp 7/17 Halogens
* history of Periodic Table * hydrogen
* Group
* Gp 0/18 Noble Gases * Period
* properties of metals * properties
of non-metals * metalloids/semi-metals * structure of Periodic Table
* Transition
Metals * Valency
 1.
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 the
notes 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, BUT
atoms
are built up of even more fundamental sub-atomic particles - the electron,
proton and neutron and the structure of an atom ultimately determines its
properties.
-
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 symbol which is a single
capital letter like H or U 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.
-
The elements are laid out in order of Atomic
(proton) Number* (*see atomic structure page).
-
Originally they were laid out in order of
'relative atomic mass'
(the old term was 'atomic
weight').
-
This is not correct for some elements now that
we know their
detailed atomic structure
(detailed GCSE notes) in terms of protons, neutrons and electrons, and of
course, their chemical and physical properties.
-
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
...
- 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.
- 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 particular block or series
e.g. from 21Sc to 30Zn are called the
1st Transition Metal Series within period 4.
- 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?
-
The maximum electrons
allowed in the shells or electronic energy levels up to atomic number 20
are:
-
When an atom has its outer
level full to the maximum number of electrons allowed, the atom is particularly
stable electronically and very unreactive. [2],[2,8]
and [2,8,8] etc. are known as the 'stable Noble Gas
arrangements',
-
BECAUSE this is
the situation with the Noble Gases: He is [2], neon is
[2,8] and argon is [2,8,8] etc.
-
These atoms are
the most reluctant to lose, share or gain electrons in any sort of
chemical interaction because they are so electronically stable.
-
Most
of the chemistry of an element is about what the outer electrons can do,
or can't, as the case maybe.
-
and the atoms of
other elements try to attain this sort of electron structure when
reacting to become more stable.
-
The most reactive metals
have just one outer electron.
-
These are the
Group
1 Alkali Metals, lithium [2,1], sodium [2,8,1],
potassium [2,8,8,1]
-
With one outer shell
electron, they have one more electron than a stable Noble Gas
electron structure.
-
So, they readily
lose the outer electron when they chemically react to try to
form (if possible) one of the stable Noble Gas electron arrangements
- which is why atoms react in the first place!
-
sodium atom ==>
sodium ion, Na ==> Na+ is [2.8.1] ==> [2.8] electronically
-
Group 1 metals
cannot lose two electrons to form an e.g. Na2+ ion because
too much energy is needed to get such a strongly held inner shell
electron involved in bonding, so you can't form NaCl2.
-
The most reactive
non-metals are just one electron short of a full outer shell.
-
These are the
Group
7/17 Halogens, namely fluorine [2,7], chlorine [2,8,7]
etc.
-
These atoms are one
electron short of a stable full outer shell and seek an 8th outer
electron to become electronically stable - yet again, this is
why atoms react!
-
They readily gain an
outer/share electrons to chemically react to form one of the
stable Noble Gas electron arrangements either by sharing electrons
(in a covalent bond) or by electron transfer forming a singly
charged negative ion (ionic bonding).
-
chlorine atom ==>
chloride ion, Cl ==> Cl- is [2.8.7] ==> [2.8.8]
electronically.
-
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.
-
The electron
arrangements of the first 20 elements are shown below.
-
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
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.) and
Group 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.
- *
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 horizontal
row 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).
- 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 ....
- Period 1
is elements 1-2, H (1) to He
(2)
- Period 2
is elements 3-10, Li (2,1) to Ne
(2,8)
- Period 3
is elements 11-18, Na (2.8.1) to Ar
(2.8.8)
- 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).
- 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
- More than three-quarters of the
109
known elements are metals (elements naturally
occur up to uranium 92, 93-109 are
'man-made' elements from the experiments of nuclear physicists.
- This work will continues as heavier and
heavier elements are likely to be made in nuclear reactions. They will
be all metals and radioactive. BUT one theory suggests that
'super-heavy' elements of about atomic number 150? may be in a nuclear
stability region and would prove most interesting to study. Chemists are
trying to predict their properties now!, so it may have started with
Mendeleev but it ain't finished yet!
- Only about 19 are
definitely are non-metal but about
7 more are semi-metals of mixed physical and chemical character.
- The metals in the periodic table are mainly found in the left hand columns (Groups 1 and 2) and in the central blocks of the transition
elements.
- There is a 'rough' diagonal division between the two principal types of element
zig-zagging from B-Al in group 3 to Te-Po in Group 6 (see
semi-metals section 2c.).
- The elements in this
'band' are sometimes referred to as 'semi-metals' or 'metalloids'
because of their 'mixture' of metallic and non-metallic physical or chemical character
e.g. the semi-conductor silicon in group 4.
- There tends to be gradual changes in
physical and chemical properties down a group e.g.
- Down Group 1 (Alkali Metals) and Group 2
the metals get more reactive.
- Down Group 7 (Halogens) the non-metals
get less reactive, their colour gets darker, their melting/boiling
points increase.
- Down Group 4 you start with a definite
non-metal carbon, and end up at the bottom with a the definite metal
lead, so there are quite significant changes in both physical and
chemical character.
- There tends to be major changes in
physical and chemical properties across a period e.g.
- Period 2 starts with a solid low melting
reactive metal lithium, in the middle there are the high melting and
rather unreactive non-metals boron and carbon, next to the end is the very
highly reactive non-metal gas fluorine, and the period finishes with the
very unreactive gas neon. Very complicated pattern!
- Period 4 starts with a solid low melting
very reactive metal rubidium, after calcium there are ten transition
metals with a wide variety of chemistry, followed by the metallic gallium,
semi-metal germanium and more non-metallic arsenic/selenium, next to the
end is the very reactive non-metal liquid bromine, and the period finishes
with the very unreactive gas krypton. Even more complicated pattern!
- From left to right across a period the
bonding in chlorides or oxides changes from ionic (with metals e.g. Na+Cl-,
Mg2+O2- to covalent (with non-metals e.g. HCl, SO2).
- 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.
- 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 (<= 3).
- The very reactive non-metals have 5 to 7
outer valency shell electrons.
- Elements in the 'middle' of the
Periodic Table e.g. Group 4 with 4 outer electrons, show non-metal e.g.
carbon or metal chemical character e.g. lead and
non are very reactive elements.
- The Noble Gas elements have full, very
stable, outer valency shells.
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.
- For Groups 1 to 8, the group number
gives the maximum valency possible and equals the number of outer
electrons (well, nearly always!).
- 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!
- 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.
- 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).
- e.g. in Group 4, C forms CO (nasty!)
but CO2 is more stable,
Pb forms PbCl2 which is much more stable than PbCl4.
- Xenon forms XeF8
and XO4 using its maximum valency of 8! and that got
somebody a Nobel Prize in Chemistry! (and in scrabble too?)
- 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.
- The valency of hydrogen is 1
(hydrides), oxygen 2 (most oxides) and chlorine is usually 1 (most chlorides).
- 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.
-
| valency of X |
1 |
2 |
3 |
4 |
5 |
6 |
| formula of hydride |
XH |
XH2 |
XH3 |
XH4 |
- |
- |
| formula of chloride |
XCl |
XCl2 |
XCl3 |
XCl4 |
XCl5 |
- |
| formula of oxide |
X2O |
XO |
X2O3 |
XO2 |
X2O5 |
XO3 |
- Examples of how to work out formula
from valencies or ionic charge is on the
Elements, Compounds and
Mixtures page and the structure,
bonding and properties of many of these substances is discussed
elsewhere.
- In the table below you can see how the formula change
horizontally from left to right with change in valency AND the vertical connection where
elements in the same group form compounds of the same formula.
|
element, hydride, chloride and oxide |
Group 1
Alkali Metals |
Group 2 |
Group 3 |
Group 4 |
Group 5 |
Group 6 |
Group 7
The Halogens |
Group 8
Noble Gases |
|
valency |
1 |
2 |
3 |
4 |
3, 5 |
2, 4,
6 |
often
1, can be 3,4,5,7 |
- |
| Period 2 |
Li
LiH
LiCl
Li2O |
Be
BeH2
BeCl2
BeO |
B
BH3
BCl3
B2O3 |
C
CH4
CCl4
CO2 |
N
NH3
NCl3
several |
O
H2O
Cl2O
O2 |
F
HF
ClF
F2O |
Ne
-
-
- |
| 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
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 the
Alkali
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 the
structure 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: reactivity
and metal-acid reactions/equations [1]
and [2]
with answers).
Metals readily form positive ions
in
compounds by losing electrons e.g.
- sodium Na - e- ==> Na+,
magnesium Mg - 2e- ==> Mg2+ or aluminium
Al - 3e- => Al3+
Their oxides and chlorides are
usually ionic*
in terms of chemical bonding. e.g.
- magnesium oxide, MgO or Mg2+O2-,
sodium oxide Na2O or (Na+)2O2-
,
- 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 Notes and 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.
- 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 <7, universal indicator yellow-orange-red
- Non-metals do not usually react with acids
e.g.
to produce a salt and hydrogen like most metals do.
- Non-metals readily form negative ions in
compounds by losing electrons e.g.
- chlorine ==> chloride: Cl2
+ 2e- ==> 2Cl- (more simply Cl +
e- ==> Cl- typical of Group 7 Halogens)
- oxygen ==> oxide: O2
+ 4e- ==> 2O2-
(more simply O + 2e-
==> O2-
typical of Group 6 elements)
- The oxides and chlorides, when combined
with other
non-metals are always covalent in terms of chemical
bonding.
- e.g. water H2O(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.
- sodium chloride, NaCl or Na+Cl-
, magnesium chloride MgCl2 or Mg2+(Cl-)2
,
- and magnesium oxide, MgO or Mg2+O2-
2c.
The Properties of Semi-metals or
Metalloids
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
6/16 |
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 character e.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-conductor and 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-conductor with metallic and non-metallic
properties and is used in photo-electric cells (solar cells) and xerography
(photocopying). (*Allotropes
are 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/hydroxides that 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-conductor with 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 0/8 Noble Gases
and the 1st series of Transition Metals.
Elements
1-20
covering Periods 1-3 and start of Period 4
Elements
1-38
covering Periods 1-4 and start of Period 5
Elements
1-96
covering Periods 1-6 and start of Period 7
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 reactive Group 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 or Na+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.
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
or CaO.
Detailed advanced
Level Chemistry Notes on Group 1 and Group 2 Metals
Group 3 (Group 13 in modern notation)
Group
4 (Group 14 in modern notation)
Group
5 (Group 15 in modern notation)
Group
6 (Group 16 in modern notation)
- 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
- 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+)2S2-.
- They form covalent small
molecule compounds with other non-metallic elements e.g. H2O
or CS2.
- The top element in the group is oxygen, a
most important element.
- Made by green plants in photosynthesis.
- Consumed in the reverse process of
respiration.
- Pure oxygen is obtained from the fractional
distillation of liquified air, though for many industrial process, the
21% in air is quite adequate to use directly (fractional
distillation is explained on another page, oxygen has a higher
boiling point than nitrogen).
- Oxygen is used in:
- oxy-acetylene burners to
produce a much hotter and intense flame for 'cutting' and welding
metal,
- oxygen 'tents' in hospitals for
respiratory problems,
- oxidant gas for burning rocket fuel.
- Oxygen combines
with most other elements to form oxides of varying physical chemical
character.
- On the left and middle of the Periodic
Table are the basic metal oxides
which react with acids to form salts e.g. Na2O, MgO, CuO etc.
These metal oxides tend to be ionic
in bonding character with high melting points. The Group 1 Alkali
Metals, and to a less extent, Group 2 oxides, dissolve in water to form
alkali solutions. All of them react with , and neutralise acids to form
salts.
- As you move left to right the oxides
become less basic and more acidic.
- So on the right you have the acidic
oxides of the non-metals CO2, P2O5,
SO2, SO3 etc. These tend to be covalent in bonding
character with low melting/boiling points. Those of sulphur and
phosphorus are very soluble in water to give acidic solutions which can
be neutralised by alkalis to form salts.
- These oxides are another example of
the change from metallic element to non-metallic element chemical
behaviour from left to right across the Periodic Table.
- BUT life is never that simple in
chemistry!:
- Some oxides react with both acids and
alkalis and are called amphoteric oxides. They are usually
relatively insoluble and have little effect on indicators. An
example is aluminium oxide dissolves in acids to form
'normal' aluminium salts like the chloride, sulphate and nitrate.
However, it also dissolves in strong alkali's like sodium hydroxide
solution to form 'aluminate' salts. This could be considered as
'intermediate' basic-acidic character in the Periodic Table.
- Some oxides are neutral, tend
to be of low solubility in water and have no effect on litmus, and
do not react with acids or alkalis. e.g. CO carbon monoxide
(note that CO2 carbon dioxide is weakly acidic) and NO
nitrogen monoxide (note that NO2 nitrogen dioxide is
strongly acidic in water). There is no way of simply predicting this
kind of behaviour from periodic table patterns!
- Sulphur is an important element used
in the GCSE notes on the
manufacture of sulphuric acid.
- 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)
-
The
Group 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. KBr or MgCl2, 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.
-
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)
The
1st Transition Metal Series (Scandium to zinc)
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 at room 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 also
Metal
Extraction (detailed GCSE notes) and more on metal uses on
the
Extra
Industrial Chemistry - detailed GCSE notes - use index of sub-pages.
Detailed Advanced
Level Chemistry Notes on the 3d-block of elements and Transition Metals
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) |
- 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 Döbereiner
-
Johann
Döbereiner noted that certain elements seemed to occur as
'triads' of similar elements e.g.
-
(i)
lithium, sodium and potassium
-
(ii)
calcium, strontium and barium
-
(iii)
chlorine, bromine and iodine
|
-
Döbereiner
was 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 |
|
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 similar 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 the
modern
Periodic Table e.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
-
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 the modern Periodic
Table'.
5e.
The
full modern
version of the Periodic Table
|
Pd |
metal 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 |
|
6 |
55Cs |
56Ba |
*57-71 |
72Hf |
73Ta |
74W |
75Re |
76Os |
77Ir |
78Pt |
79Au |
80Hg |
81Tl |
82Pb |
83Bi |
84Po |
85At |
86Rn |
|
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 |
-
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.
-
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).
-
This means that
21Sc to 30Zn can be now considered as the top
elements in the vertical Groups 3 to 12.
-
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.
-
Detailed Advanced Level Chemistry Notes on the Periodic Table
|
-
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 the
electronic
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 of electron structure.
-
We now know the electronic structure of elements and can
understand sub-levels and the 'rules' in electron structure (see
atomic
structure 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 (see radioactivity-nuclear
reactions page). 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 the structure 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?
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.
-
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 of chemical process to separate or extract an element
from a compound e.g.
-
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 (see
GCSE notes on Metal Extraction)
-
Non-metals are obtained by a variety of
means e.g. chlorine is obtained by electrolysis of sodium chloride
solution (see Group 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.
-
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.
-
Gold/platinum is are the least
reactive metals and are usually found 'native' as the yellow/silver
elemental metal.
-
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.
-
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.
-
-
 Website
content copyright
© Dr W P Brown 2000-2010 All rights reserved on revision notes, puzzles,
quizzes, worksheets, x-words etc. * Copying of website material is not
permitted * I do not personally endorse the adverts - but they do pay
for the site! |
