Periodic Table: Trends Across Period 3 Chemistry Tutorial

Key Concepts

• The horizontal rows of the periodic table are called periods.
• Period 3, or the third period, refers to the third row from the top of the periodic table.
• The elements in period 3 of the periodic table are:

  Name of Element	sodium	magnesium	aluminium (1)	silicon	phosphorus	sulfur (2)	chlorine	argon
  Chemical Symbol	Na	Mg	Al	Si	P	S	Cl	Ar

• The following general trends are observed as you go across period 3 from left to right:

  (a) atomic number, and therefore charge on the nucleus (nuclear or core charge) increases

  (b) number of valence electrons increases

  (c) atomic radius decreases

  (d) first ionisation energy increases

  (f) electronegativity increases (excluding argon)

  (g) elements on the left are metals, elements on the right are non-metals:

  (i) melting points change from high to low

  (ii) boiling points change from high to low

  (iii) electrical and heat conductors on the left to insulators on the right

  (iv) metallic lustre on the left to dull on the right

  (v) colour changes from grey on the left to coloured on the right
              (except argon which is colourless)

  (vi) Properties of the chlorides:

  (I) bond type: ionic chlorides on the left to covalent chlorides on the right

  (II) aqueous solutions: neutral on the left to acidic on the right

  (vii) Properties of the oxides:

  (I) bond type: ionic on the left to covalent on the right

  (II) aqueous solutions: basic on the left to amphoteric to acidic on the right

• The general trends in the properties of period 3 elements are summarised in the table below:

  Period 3 Elements	Na	Mg	Al	Si	P	S	Cl	Ar
  Trend: Atomic Number (Z)	low	→	→	→	→	→	→	high
  Trend: No. Electrons	low	→	→	→	→	→	→	high
  Trend: Atomic Radius	largest	→	→	→	→	→	→	smallest
  general Trend: 1st Ionisation Energy	lowest	→	→	→	→	→	→	highest
  Trend: Electronegativity	lowest	→	→	→	→	→	highest	none
  General Melting Point	higher			highest	lower
  Physical Description	shiny-grey "metallic" solid			dull grey solid	white solid	yellow solid	greenish-yellow gas	colourless gas
  Metallic Character	metals			semi-metal (metalloid)	non-metals
  Bonding in Chlorides	ionic		covalent					none
  Acidity of Chlorides	neutral		acidic					none
  Bonding in Oxides	ionic			covalent (3-D network)	covalent molecular			none
  Acidity of Oxides	basic		amphoteric	acidic				none

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Properties of Period 3 Elements

The table below gives a number of different properties of elements in period 3 of the periodic table.

Can you see any patterns, or trends, in the data?

Trend in Physical Properties of Period 3 Elements

First, let's compare the physical properties of metals and non-metals.

In general metals are:

• hard (EXCEPT Group 1, group IA or alkali metals, metals which are quite soft)
• shiny, they have metallic lustre
• solids at room temperature and pressure (Except mercury which is a liquid metal)
  That is, metals have high melting point and boiling points (Except for mercury and the Group 1 (group IA or alkali metals) metals which have low melting and boiling points compared to other metals)
• good electrical conductors

In general, non-metals are:

• dull
• brittle
• solids, liquids or gases at room temperature and pressure
  Non-metals usually have low melting and boiling points.
• electrical insulators (non-conductors of electricity)

First, let's determine the state of each element at room temperature (25°C) and pressure (100 kPa):

• Solid: melting point (M.P.) is greater than 25°C

  M.P. > 25°C (or, 25°C < M.P.)

• Liquid: melting point (M.P.) less than 25°C but boiling point (B.P.) is greater than 25°C

  M.P. < 25°C < B.P.

• Gas: boiling point (B.P.) is less than 25°C

  B.P. < 25°C

Use the data to determine the state of each of element as shown in the table below:

There are only 2 gaseous elements in period 3; chlorine and argon. These are definitely non-metals since metals DO NOT exist as gases at room temperature and pressure.⁽³⁾
We can not classify any of the other elements as metals or non-metals based on their state at room temperature and pressure because both metals and non-metals can exist as solids.
But do any of these period 3 elements look like metals?

Let's consider the physical description of each element at room temperature and pressure:

"Shiny grey solid" is a good general description of metal (some might say "silvery"), so sodium, magnesium and aluminium all look like metals.
A solid that has dull lustre, like silicon, or is coloured, like phosphorus and sulfur, look more like non-metals.
If we look at the trend in melting points and boiling points, however, we may be able to confirm that these elements are all non-metals ...

The table below gives the melting point and boiling point of period 3 elements. Let's see if we can classify each of the elements of period 3 as either a metal or a non-metal:

Elements on the left of the period 3 have higher melting and boiling points than those on the right.
Argon, a group 18 or Noble Gas has the lowest melting point and boiling point. Atoms of argon do not form molecules, so argon exists as monatomic atoms of argon, Ar_(g). One atom of argon does not experience a significant force of attraction to another atom of argon so the attraction between them is easily weakened by heat resulting in a low melting point and a low boiling point.
Chlorine also has a low melting point and low boiling point. Chlorine exists as a diatomic molecule; 2 atoms of chlorine joined together by a single covalent bond, Cl_2(g). The forces of attraction between these diatomic molecules is quite low so they are easily weakened by heat resulting in a low melting point and low boiling point.
Sulfur and phosphorus also exist as discrete molecules⁽⁴⁾ with weak forces of attraction acting between the molecules which are weakened on heating resulting in low melting points for the solids.
These low melting and boiling points are typical of non-metals. Argon, chlorine, sulfur and phosphorus are non-metals.
On the left hand side of period 3 we find the elements sodium, magnesium and aluminium. The high melting points suggest that the force of attraction between the "atoms" making up the solids are strong. We call these strong forces of attraction metallic bonds. The "atoms" are held together in a 3-dimensional array with delocalised electrons creating the metallic bonds between them. Sodium, magnesium and aluminium are all metals.
We can also see that the melting point of these metals is increasing from left to right. The more electrons a metallic element can contribute to the metallic bond the more mobile electrons there will be and the stronger the metallic bond will be!

In general, we can say that the elements become less metallic in nature (more non-metallic) across period 3 from left to right.

But what about silicon? It has the highest melting point and boiling point. Is silicon a metal?
The high melting point and boiling point reflects the strength of the strong covalent bonds holding atoms of silicon together in a 3-dimensional covalent network, similar in structure to diamond.
In this respect, silicon would appear to be a non-metal...

But let's look at another physical property of the period 3 elements, their electrical conductivity at room temperature and pressure, in order to decide if silicon is a metal or a non-metal:

Note that the metals, sodium, magnesium and aluminium are excellent electrical conductors.
The metal with the greatest number of electrons making up the metallic bond, aluminium, has the greatest value for electrical conductivity, while the metal with the least electrons contributed to the metallic bond has the lowest electrical conductivity of these metals.
That is, there is a trend in the conductivity of the period 3 metals: from left to right the conductivity of METALS increases.
The non-metals sulfur, chlorine and argon are very poor electrical conductors.
And the electrical conductivity of silicon lies somewhere in between that typical of metals and of non-metals. Silicon is best described as a semi-metal (or metalloid).

We can now see that as you go from left to right across period 3 of the periodic table the elements change from metals to a semi-metal (metalloid) to non-metals:

Trends in Atomic Structure of Period 3 Elements

Consider first the atomic number (Z) of each element in period 3 of the periodic table:

The modern periodic table is arranged in order of increasing atomic number (Z) from left to right across a period.
This means that a proton is added to the nucleus of each atom of each successive member of period 3.
Since each proton carries a charge of 1+, we say that the charge on the nucleus is increasing across the period, or that the nuclear charge is increasing, or that the core charge is increasing across the period.

Remember that an atom has no overall charge, so the total positive charge on the nucleus due to the protons must equal the total negative charge resulting from the electrons.
For an atom of each element the number of protons = the number of electrons.

But how are those electrons arranged around the nucleus of each atom?
Each atom of each element in period 3 has a filled:

• K shell (or first energy level)
  (that is, each atom has 2 electrons in the K shell)

  AND,

• L shell (second energy level)
  (that is, each atom has 8 electrons in the L shell)

The electrons in the highest energy level (last shell) are referred to as valence electrons. These are the electrons used in chemical reactions.
Consider the number of valence electrons in an atom of each element:

As you go across period 3 from left to right, the number of valence electrons increases.

As electrons are added to the same energy level (shell) across the period, the increasingly large positive charge on the nucleus pulls these electrons closer so that the size of the atoms decrease across the period from left to right.
We measure the size of an atom by its atomic radius:

Since the valence electrons (those electrons in the highest energy level) are more strongly attracted to the nucleus, they will be more difficult to remove.
We have evidence to support this idea from an inspection of the values for the first ionisation energy of the period 3 elements (the energy required to remove an electron from the gaseous atom to form a gaseous cation);

In general, the energy required to remove an electron from each atom in period 3 increases as you go from left to right across the period. ⁽⁶⁾
This suggests that elements on the left hand side of period 3 are more likely to form positive ions, that is lose an electron, than elements on the right hand side of period 3.

If we consider the ability of an atom of each period 3 element to attract elctrons towards itself (its electronegativity) we would expect the elements on the right which have a smaller atomic radius and greater nuclear charge to be better at this than those atoms on the left hand side of the period.
This trend of increasing electronegativity across period 3 from left to right is shown in the table below:

This suggests that the elements on the right hand side of period 3 (exclusing argon) are more likely to gain electrons and form negative ions (anions) than those on the left hand side of period 3.

Trends in the Chemical Properties of Period 3 Elements

With the exception of argon, period 3 elements generally react in order to achieve a stable "octet" of electrons in their valence shells.
Argon is an exception because, being a group 18 (Noble Gas) element, it already has a stable electronic configuration.
The other period 3 elements can achieve a stable octet of electrons by

• a transfer of electrons from one species to another resulting in the formation of ions and ionic bonds
• sharing of valence electrons resulting in covalent bonds

Electrons will only be transferred from one atom to another if there is a large difference in electronegativities of the two atoms (we will define a large difference as greater than 1.7)
Otherwise the valence electrons will be shared and a covalent bond formed between the atoms.

• Ionic bond: difference in electronegativity > 1.7
• Covalent bond: difference in electronegativity < 1.7

Consider how each element could form an ion with a stable electronic configuration:

Let's consider chemical reactions that produce the halide, in this case the chloride, of each element in period 3:

Sodium chloride, NaCl, is an ionic compound composed of sodium cations (Na⁺) and chloride anions (Cl^-).
The difference in electronegativity between sodium and chlorine is:

electronegativity chlorine - electronegativity sodium = 3.16 - 0.93 = 2.23
2.23 > 1.7 therefore bond is ionic

Across period 3 from left to right the bonding in the chlorides changes from ionic to covalent.

We can confirm this with reference to the melting points of the chlorides:

• High melting point indicates an solid ionic chloride.
• Low melting point indicates a covalent molecular chloride.

If you were to dissolve each of these chlorides in water, you would find another interesting trend:

Aqueous solutions of period 3 chlorides change from neutral to acidic as you go across the period from left to right.
Sodium chloride dissolves in water to produce sodium ions (Na⁺_(aq)) and chloride ions (Cl^-_(aq)) but neither ion reacts with water (hydrolyses) so the aqueous solution is neutral (because water is neutral).
This is also true for magnesium chloride.
Other chlorides do react with water to produce acidic solutions as shown in the balanced chemical equations below:

AlCl₃ + 3H₂O_(l) → 3HCl_(aq) + Al(OH)_3(s)
SiCl_4(l) + 2H₂O_(l) → 4HCl_(aq) + SiO_2(s) (silica, weakly acidic)
PCl_3(l) + 3H₂O_(l) → 3HCl_(aq) + H₃PO_3(aq) (phosphorus acid)
PCl_5(s) + 3H₂O_(l) → 5HCl_(aq) + H₃PO_4(aq) (phosphoric acid)
Cl_2(aq) + H₂O_(l) → HCl_(aq) + HOCl_(aq) (hypochlorous acid)

Let's consider an oxidation reaction for each period 3 element to produce an oxide of each element:

Now we can consider the difference in electronegativity between each period 3 element and oxygen in order to decide if each oxide is ionic or covalent:

As we go from left to right across period 3 the bonding within the oxides of the elements change from ionic to covalent.

Once again, we can confirm the ionic or covalent bonding within the oxides by inspecting the values for the melting points of the oxides as shown in the table below:

Note the high melting point of the ionic oxides (the metal oxides) and the lower melting points of the covalent molecular oxides (non-metal oxides).
The oxide of silicon, SiO₂, is a giant 3-dimensional covalent network solid similar in structure to diamond. Melting solid SiO₂ requires enough energy to weaken these strong covalent bonds so the melting point of silica (SiO₂) is very high.

If we were to dissolve each oxide in water, we would find another interesting trend.
Aqueous solutions of the oxides of elements on the left are basic, then become amphoteric, while the oxides of the non-metals on the right hand side are usually acidic:

Sodium oxide reacts with water to produce a basic solution of sodium hydroxide (NaOH_(aq)).
Magnesium oxide reacts with water to produce a basic solution of magnesium hydroxide (Mg(OH)_2(aq)).
You won't dissolve aluminium oxide in water, the bonds holding all the ions together are just too strong, but, aluminium oxide will react with both acids and bases so it is amphoteric.
Likewise silica, SiO₂, has very low solubility in water because its atoms are held together in a 3-dimensional covalent lattice, but will react with bases so it is an acid.
The oxides of phosphorus dissolve in water to produce phosporus acid (H₃PO₃) and phosphoric acid (H₃PO₄).
The oxides of sulfur dissolve in water to produce sulfurous acid (H₂SO₃) and sulfuric acid (H₂SO₄).
And the oxides of chlorine dissolve in water to produce hypochlorous acid (HOCl) and perchloric acid (HClO₄).

Problem Solving using the StoPGoPS approach