Periodic Table: Trends Across Period 2 Chemistry Tutorial

Key Concepts

• The horizontal rows of the periodic table are called periods.
• Period 2, or the second period, refers to the second row from the top of the periodic table.
• The elements in period 2 of the periodic table are:

  Name of Element	lithium	beryllium	boron	carbon	nitrogen	oxygen	fluorine	neon
  Chemical Symbol	Li	Be	B	C	N	O	F	Ne

• The following general trends are observed as you go across period 2 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 neon)

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

  (i) melting points change from high to low

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

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

  (iv) halides change from ionic on the left to covalent on the right

  (v) oxides change from ionic and basic on the left to covalent and acidic on the right

• The general trends in the properties of period 2 elements is summarized in the table below:

  Trends in Properties of Period 2 Elements
  Chemical Symbol	Li	Be	B	C	N	O	F	Ne
  Metallic Character	metals	→	semi-metal	→	non-metals
  Bond Type	metallic			→	covalent
  Melting Point	high			→	low
  Conductivity	high			→	low
  Lustre	metallic			→	dull
  Atomic Radius	high			→	low
  1st Ionisation Energy	low			→	high
  Electronegativity	low			→	high
  Atomic Number	low			→	high
  No. electrons	low			→	high
  Ions of Elements	cations			→	anions
  Halides of Elements	ionic	→	covalent
  Oxides of Elements	basic		→ amphoteric →			acidic

Please do not block ads on this website.
No ads = no money for us = no free stuff for you!

Properties of Elements in Period 2 of the Periodic Table

The table below gives a number of different properties of the elements of period 2 of the periodic table.

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

Physical Properties of Period 2 Elements

Elements to the left of period 2 are metals, those to the right are non-metals.
Metals usually have higher melting and boiling points than non-metals, so let's examine the melting points of period 2 elements:

The elements on the right, nitrogen, oxygen, fluorine and neon all have low melting points and are all non-metals.
In fact, apart from neon which exists as a monatomic gas (Ne_(g)) at room temperature and pressure, the others are all diatomic gases, nitrogen gas (N_2(g)), oxygen gas (O_2(g)) and fluorine gas (F_2(g)).
Only weak intermolecular forces (London forces or dispersion forces) act between these covalent molecules, so little energy is required to disrupt the attraction and melt the solid, or indeed, boil the liquid to produce a gas.

The elements on the left, lithium and beryllium have high melting points and are metals.
Strong metallic bonds hold the "atoms" in a 3-dimensional array and it requires a lot of energy to disrupt these attractive forces so the melting points are high.

The melting point of boron is very high, so is it a metal? ,
Boron has other properties such as being a semiconductor, not a metallic conductor which place it in the borderline region between being a metal and a nonmetal so we shall classify it as a semi-metal (or metalloid).

What about the high melting point of carbon, is carbon a metal?
Carbon has other properties that make it definitely a non-metal, such as the formation of covalent bonds between carbon atoms which produce large (infinite) 3-dimensional lattices that make graphite and diamond (see allotropes).
It requires a lot of energy to disrupt these strong covalent bonds holding the carbon atoms in place in the lattice so the melting point of carbon (either as graphite or diamond) is high.

So, we see a trend from high to low melting points across period 2 from left to right, a trend from solid to gas, as well as a trend from metal to semi-metal to non-metal, and a trend from metallic bonds to covalent network solid to diatomic molecules to monatomic gas:

We can also look at the conductivity of period 2 elements to confirm the same trend: metallic to semi-metal (metalloid) to non-metal:

The delocalised electrons in the structure of a metal allows metals to conduct electricity well. Lithium and beryllium are metals.
At room temperature, boron is a poor conductor of electricity, but its conductivity increases when it is heated. Boron is a semi-metal (or metalloid).
The delocalised electrons between the carbon layers of graphite allow electricity to be conducted in this way, but the electrons are "locked up" in covalents bonds in the diamond structure so diamonds do not conduct electricity, diamond is an insulator.
Similarly, there are no mobile electrons to conduct electricity in the structure of the other non-metals.

The trend from conductor to poor conductor to insulator across period 2 from left to right confirms the trend from metal to semi-metal (metalloid) to non-metal.

Atomic Structure of Period 2 Elements

Consider the atomic number of each period 2 element :

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 2.
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 arround the nucleus of each atom?
Each atom of each element in period 2 has a filled K shell (or first energy level), that is, each of atoms has 2 electrons in the K shell.
Each subsequent electron is positioned in the L shell, or second energy level.
With the group 18 element neon, Ne, this second energy level (L shell) is complete.

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 element (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 2 increases as you go from left to right across the period. ⁽¹⁾
This suggests that elements on the left hand side of the period 2 are more likely to form positive ions, that is lose an electron, than elements on the right hand side of period 2.

If we consider the ability of an atom of each period 2 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 2 from left to right is shown in the table below:

This suggests that the elements on the right hand side of period 2 (exclusing neon) are more likely to gain electrons and form negative ions (anions) than those on the left hand side of period 2.

Chemical Properties of Period 2 Elements

With the exception of neon, period 2 elements react in order to achieve a stable "octet" of electrons in their valence shells.
Neon is an exception because, being a group 18 (Noble Gas) element, it already has a stable electronic configuration.
The other period 2 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.
Otherwise the valence electrons will be shared and a covalent bond formed between the atoms.

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

Let's consider chemical reactions that produce the halide of each element in period 2:

Lithium chloride, LiCl, is an ionic compound composed of lithium cations (Li⁺) and chloride anions (Cl^-).
The difference in electronegativity between lithium and chlorine is:

electronegativity chlorine - electronegativity lithium = 3.16 - 0.98 = 2.18
2.18 > 1.7 therefore bond is ionic

Across period 2 from left to right the bonding in the halides changes from ionic to covalent.

Let's consider an oxidation reaction for each period 2 element to produce an oxide of each element:

Now we can consider the difference in electronegativity between each period 2 element and oxygen in order to decide if each oxide is ionic or covalent:

Once again we see a trend across period 2 from left to right, the oxides change from being ionic to being covalent.

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:

Problem Solving