Chapter 13: the periodiC table
Contents
13.1-Classifying the elements
13.2-Trends in groups
13.3-Trends across a period
Learning objectives
Core
1. Describe the Periodic Table as an arrangement of elements in periods and groups
and in order of increasing proton number/ atomic number
2. Describe the change from metallic to non-metallic character across a period
3. Describe the relationship between group number and the charge of the ions formed
from elements in that group
4. Explain similarities in the chemical properties of elements in the same group of the
Periodic Table in terms of their electronic configuration
5. Explain how the position of an element in the Periodic Table can be used to predict its
properties
Supplement
6. Identify trends in groups, given information about the elements
The Periodic Table
Page 1 of 14
Notes by: Ms.Jarina
13.1 Classifying the elements
•
•
•
•
The Periodic Table is a way of classifying the elements. All modern versions of
the Periodic Table are based on the table proposed by Mendeleev in1869.
Elements are arranged on the Periodic Table in order of increasing atomic
number, where each element has one proton more than the element before it.
The table is arranged in vertical columns called Groups numbered I - VIII and in
rows called Periods.
Period: these are the horizontal rows that show the number of shells of electrons
an atom has. Eg elements in Period 2 have two electron shells, elements in
Period 3 have three electron shells
•
Group: The group number is the same as the number of outer-shell electrons in
the atoms, except for Group 0.
•
In Group I the atoms have one outer shell electron, in Group II they have two,
and so on.
•
The outer-shell electrons are also called valency electrons. And they are very
important: they dictate how an element behaves.
•
So, all the elements in a group have similar reactions, because they have the
same number of valency electrons.
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The atoms of the Group 0 elements have a very stable arrangement of electrons
in their outer shells. This makes them unreactive.
Metals and non-metals
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The main distinction in the Periodic table is between metals and non-metals.
•
The non-metals are grouped into the top right-hand region of the table, above the
thick stepped line.
•
The elements close to the line show properties that lie between the metals and
non- metals. These elements are now often referred to as metalloids (or semimetals). These elements have some of the properties of metals and others that
are more characteristic of non-metals. There are eight elements that are called
metalloids. They often look like metals but are brittle like non-metals. They are
neither conductors nor insulators but make excellent semiconductors. The prime
example of this type of element is silicon.
•
The elements present in Groups I to VIII / 0 of the table known as the maingroup elements. These vertical groups show most clearly how elements within
the same group have similar chemical and physical properties.
Page 2 of 14
Notes by: Ms.Jarina
•
Between Groups II and III of these main groups of elements is a block of metals
known as the transition elements (or transition metals). The first row of these
elements occurs in Period 4. This row includes such important metals as iron,
copper and zinc.
•
The noble gases, in Group VIII / 0 on the right-hand side of the table, are the
least reactive elements in the table.
•
Group VII, (halogens) and Group I (alkali metals), are the most reactive
elements.
•
The more unreactive elements, whether metals or non-metals, are in the centre
of the table.
The position of hydrogen in the Periodic Table
As the smallest atom of all, the properties of hydrogen are unique. It does not fit easily
into the trends shown in any one group.
Page 3 of 14
Notes by: Ms.Jarina
Electronic configuration and the Periodic Table
The outer electrons of an atom are mostly responsible for the chemical properties of any
element. Therefore, elements in the same group will have similar properties.
•
Elements in the same group have the same number of electrons in their outer
shell. For the main-group elements, the number of the group is the number of
electrons in the outer shell.
•
The periods also have numbers. This number shows us how many shells of
electrons the atom has.
•
Certain electron arrangements are found to be more stable than others. This
makes it more difficult to break up. The most stable arrangements are those of
the noble gases and they are so unreactive.
Why do atoms get smaller as you go across a period?
Within a period, all electrons are added to the same shell. At the same time, protons are
added to the nucleus, making it more positively charged. The effect of increasing proton
number is greater than that of the increasing electron number. This means that the
nucleus attracts the electrons more strongly, pulling the atom's shell closer to the
nucleus. As a result, the atomic radius decreases.
Page 4 of 14
Notes by: Ms.Jarina
The relationship between group number and ionic charge
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For the metal ions the positive charge is the same as the group number. These
atoms have lost their outer electrons to form the ion.
•
For non-metals, the negative charge on the ion is 8 minus the group number.
These atoms have gained electrons to form the negative ion.
13.2 Trends in groups
1. The metals in Group I are called alkali metals.
2. They are soft solids with relatively low melting points and low densities.
3. They are highly reactive and are stored in oil to prevent them reacting with the
oxygen and water vapour in the air.
4. When freshly cut with a knife, all these metals have a light-grey, silvery surface, which
quickly tarnishes (becomes dull).
5. There are several different trends in this group of elements:
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the melting points and boiling points decrease down the group,
•
the metals get softer as we go down the group,
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the densities increase as we go down the group.
Page 5 of 14
Notes by: Ms.Jarina
* the densities of rubidium and caesium are 1.53 and 1.88, respectively
By observing the trends in these physical properties, we can also predict values for
other alkali metals
The reactions of the alkali metals with water.
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The alkali metals (Group I) are the most reactive metals that occur. They are
known as alkali metals because they react vigorously with water to produce
hydrogen and an alkaline solution. The equation for the reaction between
sodium and water is as follows:
sodium + water → sodium hydroxide + hydrogen
2Na + 2H2O → 2NaOH + H2
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The reactions range from vigorous in the case of lithium to explosive in the case
of caesium
•
When a few drops of universal indicator is added to the water, the solution will
turn purple showing that an alkali is present.
Page 6 of 14
Notes by: Ms.Jarina
The reactions of the alkali metals with water.
Group VII – the halogens
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The most reactive non-metals are the halogens in Group VII of the table. In
contrast with Group I, reactivity decreases down the group.
•
For example, fluorine is a dangerously reactive, pale yellow gas at room
temperature.
•
There is a steady increase in melting points and boiling points as we go down the
group, and the elements change from gases to solids as the atomic number
increases.
•
The lowest element in this group is a highly radioactive and rare element,
astatine.
Page 7 of 14
Notes by: Ms.Jarina
The general properties of some of the halogens (Group VII).
Page 8 of 14
Notes by: Ms.Jarina
The common properties of halogens are:
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They are all poisonous and have a strong smell.
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They are all non-metals.
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They all form diatomic molecules (for example Cl2, Br2, I2).
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They all have a valency (combining power) of 1 and form compounds with
similar formulae, for example hydrogen chloride (HCl), hydrogen bromide (HBr),
hydrogen iodide (HI).
•
Their compounds with hydrogen are strong acids when dissolved in water, for
example hydrochloric acid (HCl), hydrobromic acid (HBr), hydriodic acid (HI).
•
They produce a series of compounds with other elements: chlorides, bromides
and iodides. Together these are known as halides.
•
The react directly with metals to form metal halides (or salts).
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They all form negative ions carrying a single charge, for example chloride ions
(Cl-), bromide ions (Br -), iodide ions (I-).
The chemical reactivity of the halogens
1. When chlorine is passed over heated aluminium, the metal glows white and
forms aluminium chloride:
heat
2Al + 3Cl2 ⎯⎯→
2AlCl3
Aluminium also reacts strongly with bromine and iodine.
2. Hydrogen will burn in chlorine to form hydrogen chloride.
H2 + Cl2 → 2HCl
3. Chlorine dissolves in water to give an acidic solution. This mixture is called
chlorine water and contains two acids. Chlorine water acts as an oxidising
agent
Cl2 + H2O
→
HCl
hydrochloric acid
+
hypochlorous acid
Page 9 of 14
Notes by: Ms.Jarina
HClO
Test for chlorine
Damp litmus or Universal Indicator paper is bleached when held in the gas.
The halogen displacement reactions
The halogens become steadily less reactive as you go down the group.
For example, if you add chlorine to a solution of potassium bromide, the chlorine
displaces bromine. Chlorine is more reactive than bromine, so it replaces it, and
potassium chloride is formed. Potassium bromide solution is colourless. It turns orange
when chlorine is bubbled through it:
Cl2 + 2KBr
→
2KCl + Br2
colourless
orange
Chlorine will also displace iodine from potassium iodide:
Cl2 + 2KI → 2KCl +
colourless
I2
yellow–brown
Group VIII / 0 – the noble gases
All the noble gases are present in the Earth’s atmosphere. Together they make up
about 1% of the total, though argon is the most common. These gases are unreactive.
They were sometimes referred to as inert gases, meaning they did not react at all.
However, since the 1960s, some compounds of xenon and krypton have been made
and their name was changed to noble gases.
The uses of noble gases depend on their unreactivity.
Helium - Used in airships and balloons because it is both light and unreactive.
Neon – Used in brightly coloured advertising lights
Argon - Used to fill light bulbs because it will not react with the filament even at high
temperatures.
Page 10 of 14
Notes by: Ms.Jarina
The transition elements
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The general features of the transition elements make them the most useful
metals available to us. They are much less reactive than the metals in group I to
II.
•
The transition metals are one of the major contributors of colours in our lives.
Trace amounts of these metals produce the colour of gemstones such as
sapphire and ruby.
General features of transition metals (or transition elements):
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They are hard and strong.
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They have high density.
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They have high melting and boiling points.
Two of their distinctive properties:
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Many of their compounds are coloured.
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The transition metals, or their compounds, are often used as catalysts Iron, for
example is used as a catalyst in the Haber process for making ammonia.
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They often show more than one valency (variable oxidation state) – they form more
than one type of ion. For example, iron can form compounds containing iron(II) ions
(Fe2+) or iron(III) ions (Fe3+).
Page 11 of 14
Notes by: Ms.Jarina
13.3 Trends across a period
The first period of the table contains just two elements, hydrogen and helium.
Each of the remaining periods of elements starts with a reactive alkali metal and
finishes with an unreactive, non-metallic, noble gas.
In periods 2 and 3 there is a shift from metallic lattice to giant covalent structure and
then to simple molecular structure as we move across the table.
In Period 3, from sodium to argon, there appears to be a gradual change in physical
properties across the period. The change in properties seems to centre around silicon.
The elements before this are metals and those after this are non-metals
As you move across a period there are also differences in some chemical properties
linked to the metal to non-metal transition. Metal oxides (magnesium oxide) are usually
basic, while those of non-metals are acidic (sulfur dioxide). Therefore, there is a change
from basic oxides to acidic oxides as we move across a period.
If we look at Period 4 in the Periodic Table, we see that there is a whole ‘block’ of
elements in the centre of the table. This block of elements falls outside the main groups
of elements. They are best considered not as a vertical group of elements but as a row
or block. They are usually referred to as the transition elements (or transition metals).
They are some of the most useful metallic elements available to us.
Why does the size of atom decrease as we go across a period?
Within a period, protons are added to the nucleus as electrons are being added to the
same energy level.(number of shells remain the same) These electrons are gradually
pulled closer to the nucleus because of its increased positive charge. Since the force of
attraction between nuclei and electrons increases, the size of the atoms decreases.
Page 12 of 14
Notes by: Ms.Jarina
The changes in properties of the elements in Period 3 and in Group IV of the Periodic Table.
1. The first period of the Periodic table contains just two elements. They are
_______________ and _______________.
2. The first element in period 3 is ________________ and the last element is
________________.
3. The change from metal to non-metal across period 3 also shows changes in
bonding type from _________________ bonding to _______________bonding.
4. Metal oxides are usually _________________ (acidic, basic), while those of nonmetals are ________________ (acidic, basic).
5. In the above diagram, as we go down group _______________, the change is
from non-metal to metal.
6. The metalloids in this group are _______________ and _______________.
Page 13 of 14
Notes by: Ms.Jarina
Checklist
Group I properties (Core / supplement)
Describe the Group I alkali metals, lithium, sodium and potassium, as relatively soft metals with
general trends down the group, limited to:
(a) decreasing melting point
(b) increasing density
(c) increasing reactivity 2
Predict the properties of other elements in Group I, given information about the elements.
Group VII properties (Core/supplement)
Describe the Group VII halogens, chlorine, bromine and iodine, as diatomic non-metals with
general trends down the group, limited to:
(a) increasing density
(b) decreasing reactivity 2 State the appearance of the halogens at r.t.p. as:
(a) chlorine, a pale yellow-green gas
(b) bromine, a red-brown liquid
(c) iodine, a grey-black solid
3 Describe and explain the displacement reactions of halogens with other halide ions
4 Predict the properties of other elements in Group VII, given information about the elements
Transition elements - Core/ Supplement
1 Describe the transition elements as metals that:
(a) have high densities
(b) have high melting points
(c) form-colored compounds
(d) often act as catalysts as elements and in compounds
Supplement
2 Describe transition elements as having ions with variable oxidation numbers, including iron(II)
and iron(III)
Noble gases - Core / supplement
Describe the Group VIII noble gases as unreactive, monatomic gases and explain this in terms of
electronic configuration
Page 14 of 14
Notes by: Ms.Jarina