Intramolecular bonding versus Intermolecular forces

Intramolecular bonding is the name given to the chemical bonds inside molecules. All intramolecular bonds are strong. The following are some examples:

• A pure covalent bond is a pair of electrons equally shared between atoms.
• A polar covalent bond is a pair of electrons unequally shared between atoms.
•  A dative covalent bond is a covalent bond where both electrons come from the same atom.

Intermolecular bonding is the name given to forces of attraction between molecules. There are three types, all of which are weak:

• Van der Waals forces
• Dipole-dipole
• Hydrogen bonding

1. Van der Waals forces

         Out of the three, this is the weakest intermolecular force. They form between temporary dipoles in one molecule and induced dipoles in another. They are constantly changing and the strength of attraction depends on the number of electrons in the molecule.

The shape of the molecule is also important:

A large surface area = more Van der Waals forces = strong attraction.

Below is showing Van der Waals forces occurring between 2 molecules of hydrogen:

The particle on the left has formed a temporary dipole maybe caused by one side of it being more dense than the other. For a brief moment one end of the particle has a positive charge and the other end with a negative charge, causing a nearby particle to also have a temporary dipole.

2. Dipole-dipole

         These forces are stronger than Van der Waals and occur between polar molecules meaning that they have permanent dipoles.

Remember! The greater the difference in electronegativity the bigger the dipole and the stronger the attraction is between the molecules (refer back to my previous blog 

The diagram below is showing the dipole-dipole attraction between two HCl molecules. The negative dipole on the chlorine atom is attracted to the positive dipole of another HCl molecule:

3. Hydrogen bonding

         This is the strongest intermolecular force but it only occurs between molecules in which hydrogen is bonded to nitrogen, oxygen or fluorine.

Hydrogen bonding is the attraction between an electron deficient atom on one molecule and a lone pair of electrons on a highly electronegative atom on a different molecule.

The diagram above shows hydrogen bonding in water. The positive dipole on the hydrogen atom is attracted to the lone pair of electrons on the oxygen atom in another water molecule.

Electronegativity

The definition of electronegativity is “the ability of an atom to attract a shared pair of electrons in a covalent bond”.

The first thing to learn is that electronegativity increases (strengthens) as you go along the periods of the periodic table and decreases (weakens) when you go down the groups, so the top right of the periodic table contains the elements with the higher electronegativities and the bottom left corner are the elements with the lower electronegativities.

Now, if we look at the periodic table, we can state that fluorine is the most electronegative element and francium is the least.

You are probably wondering why helium isn’t the most electronegative element and the answer is that noble gases are not included as they all have a full outer shell of electrons and therefore do not form a covalent bond with another element.

Electronegativity can also be used to find out the type of bond formed between two atoms by calculating the difference in value between the electronegativites of the two atoms:

• If the difference is large then the bonding is most likely to be ionic.
• If the difference is small then the bonding is most likely to be polar covalent.
• If there is no difference then the bonding will be pure covalent.

So the two elements that show the biggest difference in electronegativity are fluorine and francium and we would expect an ionic bond to form if these two elements came together.

Another important point to know is that metal elements tend to have low electronegativity values which indicates that they only have weak attractions for electrons.

And non-metal elements tend to have high electronegativity values indicating that they have strong attractions for electrons and they will form ionic bonds.

Finally, below shows some electronegativity values in relation to the periodic table and also some worked examples showing the electronegativity difference between atoms and the type of bonding formed:

Worked examples (using the table above)

• NH₃ – electronegativity difference of 3.0 – 2.1 = 0.9 (small). Type of bonding – polar covalent.
•  CaCl₂ – 3.0 – 1.0 = 2.0 (large). Type of bonding – ionic.
• BrCl – 3.0 – 2.8 = 0.2 (small). Type of bonding – polar covalent.
• NaCl – 3.0 – 0.9 = 2.1 (large). Type of bonding – ionic.

 http://en.wikipedia.org/wiki/File:Periodic_table.svg

Shapes of Molecules

Do you know what shape a H₂O molecule is? How about CH₄? You may have correctly guessed that CH₄ is tetrahedral (tetra meaning four).

This blog will hopefully show you how to correctly work out the shapes of molecules and ions, including those you may have never heard of!!

If you read my ‘Ionic and Covalent Bonding’ blog, you should know that the outer electrons of atoms are found in pairs in orbitals.

Remember that electrons have a negative charge so when an electron pair come together they will repel each other (like charges repel, unlike charges attract).

Because the electrons repel each other, they also want to be as far apart as possible from each other. This forms the basis of what is known as the electron pair repulsion theory which states that:

The following table shows how to work out the shapes of molecules that contain between 2 and 6 pairs of electrons in their outer shell. I also recommend drawing a dot and cross diagram of the molecule first, counting the number of electron pairs then referring to the table below:

Worked examples

Ø  Molecules with lone pairs of electrons

When trying to figure out the shape of a molecule that has a lone pair of electrons, an important thing to remember is that lone pairs repel more than bonding pairs.

See below:

It may not look like it but all the above molecules have 4 electron pairs around the central atom so they are all based upon the tetrahedral shape. However, the second molecule has only 3 bonded pairs and the 4^th pair of electrons is known as a lone pair which we can not see, so we say the shape is trigonal pyramidal and not tetrahedral.

A rough guide when working out the bond angle is that each lone pair pushes a bonded pair 2º closer together.

As I mentioned earlier, I recommend drawing a dot and cross diagram of the molecule first so you know for sure how many lone pairs you have.

Worked example

Formulae and Equations

It is always useful to know what elements make up a compound and just by learning a few simple rules it is a lot easier to name and recognise them:

• A compound made up of only two elements will end in –IDE, e.g. sodium chloride.
• When naming a compound, the metal is always written first.
• If oxygen is present the name will end in –ATE (or sometimes –ITE), e.g. magnesium sulphate contains magnesium, sulphur and oxygen.

Below are some examples of compounds with the elements stated that they contain:

• copper oxide – copper and oxygen
• hydrogen sulphide – hydrogen and sulphur
• sodium phosphate – sodium phosphorus and oxygen
• caesium chloride – caesium and chlorine
• sodium hydrogencarbonate – sodium, hydrogen, carbon and oxygen

 It can also be done the other way round:

• sodium and oxygen – sodium oxide
• zinc and sulphur – zinc sulphide
• oxygen and potassium – potassium oxide
•  nickel, phosphorous – nickel phosphate
• chlorine and copper – copper chloride

Ø  Do you know your ions?

Remember! An ion is a charged particle that has either lost or gained electrons.

The following is a table of some common ions that you might come across. Most of them you can work out from their position in the periodic table, but others you will just have to learn (who said Chemistry was fun?!)

Ø  The Formula!

When working out the formula for an ionic compound, follow these three simple steps:

• Look at the name of the compound
• Decide which ions are in it
• Balance the charges

Look at the following examples:

Ø  Compounds containing complex ions

Ions such as carbonate, CO₃^2- and sulphate, SO₄^2- are known as complex ions as they contain more than one atom.

When balancing a formula that contains a complex ion, you need to put it in brackets, e.g. if you need two nitrates to balance a 2+ charge then you write it as (NO₃)_{2 ….}not NO₃₂! Below are a couple more examples:

 

Ø  Formulas of molecules

Molecules are made up of atoms not ions so when working out the formula we can’t use the balancing the charges method as shown before. Instead we have to use its position in the periodic table to work out how many bonds it will form:

• Carbon is in group 4 of the periodic table; it has 4 electrons in its outer shell so it can form 4 bonds to get a complete outer shell.
• Chlorine is in group 7 of the periodic table; it has 7 electrons in its outer shell so it forms 1 bond to get a full outer shell.
• So 1 carbon atom can bond with 4 chlorine atoms so the formula is CCl₄.

Here are some more examples:

• Hydrogen and bromine - hydrogen is in group 1 of the PT, it has 1 electron in its outer shell so forms 1 bond to get a full outer shell. Bromine is in group 7 of the PT, it has 7 electrons in its outer shell so forms 1 bond to get a full outer shell. So 1 hydrogen atom can bond with 1 bromine atom so the formula is HBr.
• Nitrogen and hydrogen – nitrogen is in group 5 of the PT, it has 5 electrons in its outer shell so forms 3 bonds to get a full outer shell. As said above, hydrogen can form 1 bond to get a full outer shell. So 3 hydrogen atoms can bond with 1 nitrogen atom so the formula is NH₃.
• Phosphorous and oxygen – phosphorous is in group 5 of the PT, it has 5 outer shell electrons so can form 3 bonds. Oxygen is in group 6, it has 6 outer shell electrons so can form 2 bonds. 2 phosphorous atoms can bond with 3 oxygen atoms so the formula is P₂O₃.

Ø  Just one last bit…. balancing equations

A chemical equation tells us the reactants and products for that reaction.

When they are balanced they can also tell us the amount of each reactant and product.

Follow these few steps and you won’t go far wrong (fingers crossed):

• Write down the word equation for the reaction.
•  Now write down the (correct) formulas for each compound in the equation.
• Now make sure that the number of atoms of each element are the same on both sides i.e. balanced.

Ionic and Covalent Bonding

Group 8 or 0 of the periodic table contain a group of atoms called noble gases; Helium, Neon, Argon, Krypton, Xenon and Radon.

They all have a full outer shell of 8 electrons, which make them very stable and unreactive.

All the other atoms in the periodic table want to be just like the noble gases and do so by bonding with other atoms to form a complete outer shell.

Atoms can do this by either transferring their electrons to another atom forming an ionic bond, or by sharing to form a covalent bond.

Ionic bonding

• Usually occurs between a metal and a non-metal. 
• When metals react they usually lose electrons. As they now have fewer electrons than protons they form positive ions.
• When non-metals react they usually gain electrons. As they now have more electrons than protons they form negative ions.
• So when a metal atom bonds with a non-metal atom, the metal donates its electrons to the non-metal to form a positive metal ion and a negative non-metal ion.
• The ions are attracted to each other by an ionic bond.

Below is a dot and cross diagram to show what happens when a sodium atom bonds with a chlorine atom:

 The sodium ion that forms has a +1 charge and the chloride ion has a -1 charge. The formula of the compound formed is NaCl. 


Covalent bonding

• Usually occurs between two non-metals.
• As explained before, non-metals usually gain electrons but as they both can’t gain electrons they share their electrons between them. This is known as a covalent bond, with a zero overall charge.

Below is a dot and cross diagram showing what happens when two chlorine atoms bond together to form a chlorine molecule (again only the outer shell electrons are shown):

Each chlorine atom has 7 electrons in their outer shell so when they come together a single covalent bond is formed from one shared pair of electrons. The formula of the chlorine molecule that forms is Cl₂.

When two oxygen atoms come together, it is slightly different because they each have 3 pairs of electrons in their outer shell, so when the two atoms attract a double covalent bond is formed consisting of two pairs of electrons.

Electron Structure

Atoms consist of electrons circling the nucleus of the atom in paths called shells or orbitals.

Electron shells increase in energy, as they get further away from the nucleus.
There are four types of sub shells; s, p, d and f. Each sub-shell contains orbitals:

• The s sub-shell contains 1 x s orbital
• The p sub-shell contains 3 x p orbitals
•  The d sub-shell contains 5 x d orbitals
•  The f sub-shell contains 7 x f orbitals

An orbital is a region of space where electrons with “opposite spin” are found.

Each orbital holds a maximum of 2 electrons so:

• The s sub-shell can hold 2 electrons
• The p sub-shell can hold 6 electrons
• The d sub-shell can hold 10 electrons
• The f sub-shell can hold 14 electrons

The table below gives a summary of the information above but also relates it to how the periodic table is built up. For example, electron shell 1 is relating to the elements along period 1 of the periodic table, electron shell 2 is relating to the elements along period 2, and so on. Please note that it only goes up to the element Krypton, atomic number = 36.

Electron shell	Maximum number of electrons	Sub-shells
		1 x s	3 x p	5 x d	7 x f
1	2	2	-	-	-
2	8	2	6	-	-
3	18	2	6	10	-
4	32	2	6	10	14

Electrons are fed into the shells of an atom with the lowest energy sub-levels filled up first. The superscripts show the maximum number of electrons that sub-level can hold. Note that the 4s sub-level is filled before the 3d:

            1s²          2s²          2p⁶          3s²          3p⁶          4s²          3d¹⁰..           

            Low energy                                                           High energy

Lets take the magnesium atom, which has 12 electrons and an electron configuration of 2.8.2. The electrons are then fed into the sub-levels, giving an electronic configuration of:

            1s²          2s²             2p⁶            3s²

Another example would be krypton, which has 36 electrons and its electron configuration is 2.8.18.8. The electronic configuration is:

            1s²             2s²          2p⁶         3s²          3p⁶         4s²         3d¹⁰         4p⁶

Electrons fill sub-shells singly before pairing up due to lower repulsion when unpaired (more stable). This can be shown by using the ‘electrons in boxes’ notation to help make this clearer:

            e.g. The oxygen atom has an electronic configuration of 1s² 2s² 2p⁴, the 8 electrons occupy the sub-shells in the following way:

                                        1s           2s                    2p

Note! The electrons fill up the 2p sub-shells singly before pairing up.

Below shows the ‘electrons in boxes’ notation from the lowest energy sub-level, 1s, up to higher energy sub-level, 4d:

Atomic Structure...simplified

Atom                smallest part of an element. Cannot be broken down into anything simpler in a chemical reaction.

Molecule          two or more atoms covalently bonded together.

Ion                    a charged particle. Usually formed when atoms lose or gain electrons.

Compound      substance containing two or more elements bonded together.

Proton             a particle found in the nucleus. It has a mass of 1 and a charge of +1.           

Neutron           a particle found in the nucleus. It has a mass of 1 and a charge of 0.

Electron          found in shells around the nucleus. It has a mass of ¹/₂₀₀₀ and a charge of -1.

The nucleus is tiny compared to the size of the atom but it contains nearly all of the mass of the atom.

In an atom the number of protons is equal to the number of electrons, therefore, the charges cancel each other out so the atom has no charge.

If an atom has gained an electron(s) it becomes a negative ion and so has more electrons than protons.

If an atom has lost an electron(s) it becomes a positive ion and so has more protons than electrons.

• ‘12’ is representing the atomic number, which tells us the number of protons.

• ‘6’ is the mass number that equals the number of protons + the number of neutrons.

• So the number of neutrons  =  mass number  -  atomic number

Isotopes            atoms of the same element with the same number of protons (atomic number) but different number of neutrons (mass number).

They have the same number of electrons in their outer shell and therefore do the same chemical reactions.

They can be detected by using a mass spectrometer, which I will talk about in one of my later blogs. 

Relative Atomic Mass (Ar)

The actual mass of an atom is approximately 10^-23 g (now that is small!) and working with it when doing calculations can be a bit difficult and to be honest a waste of time! Instead we use what is known as its relative atomic mass.

The ¹²C isotope is used as a standard to which all other atoms are compared. One atom of ¹²C has a mass of 12 so ¹/₁₂ of an atom of ¹²C has a mass of 1….so now it is a lot more easier to work out how many times heavier an atom is compared to ¹/₁₂ of the mass of an atom of ¹²C.

If you are still confused, it can be put into even simpler terms…

The average mass of an atom of the element

¹/₁₂ of the mass of one atom of ¹²C

We also use the same idea for the masses of individual isotopes – Relative isotopic mass:

The average mass of an atom of the isotopes

¹/₁₂ of the mass of one atom of ¹²C

As you may or may not know, all elements contain isotopes (isotopes are atoms of the same element with the same number of protons but different number of neutrons), so when calculating the relative atomic mass of an element we need to take this into account. There are two things you need to know when doing these calculations:

1.     the relative isotopic mass of each isotope

2.   the percentage of each isotope in the element

Below are a couple worked examples to calculate the relative atomic mass (There are a few ways to set your working out but I find this way really simple):

1.    Chlorine consists of two isotopes – 75.5% of ³⁵Cl and 24.5% of ³⁷Cl.

³⁵Cl  x  75.5%  à  (75.5 ÷ 100)  x  35  =  26.425

³⁷Cl  x  24.5%  à (24.5 ÷ 100)  x  37  =  9.065

26.425  +  9.065  =  35.49 (35.5 to 3sf)

2.    Magnesium consists of three isotopes – 78.6% of ²⁴Mg, 10.1% of ²⁵Mg and 11.3% of ²⁶Mg.

²⁴Mg  x  78.6%  à  (78.6 ÷ 100)  x  24  =  18.864

²⁵Mg  x  10.1%  à  (10.1 ÷ 100)  x  25  =  2.525

²⁶Mg  x  11.3%  à  (11.3 ÷ 100)  x  26  =  2.938

18.864  +  2.525  +  2.938  =  24.3 to 3sf

Salts

Below you will find my attempt at writing a GCSE resource topic on SALTS for students and teachers.

Ø  Salts

You have already been introduced to acids and alkalis, but what actually happens when an acid and an alkali react together?

The general word equation for this is: acid + alkali --> salt + water

This reaction is also known as neutralisation as the alkali has neutralised the acid by removing the H+ ions and turning them into water:

H⁺(aq) + OH^-(aq) --> H₂O(l)

Lets look at an example for the reaction between an acid and alkali:

Hydrochloric acid + sodium hydroxide --> sodium chloride + water

           HCl(aq)        +          NaOH(aq)   -->        NaCl(aq)      + H₂O(l)

You will notice that the salt, sodium chloride is dissolved in water (aq); later on in this chapter we will talk about how to get the salt from its solution.

Q. Can you remember what the difference is between an alkali and a base?

Ø  Preparation of soluble salts

Below are some general methods of preparation of a salt that can be made from acids:

1.     Acid + metal  -->salt + hydrogen

E.g. hydrochloric acid + magnesium --> magnesium chloride + water

               2HCl(aq)       +       Mg(s)   -->      MgCl₂(aq)           + H₂O(l)

Most metals react with acid giving off hydrogen gas; however the metal copper does not react with acid.

       Q. Which acid and metal would you start with to make aluminium sulphate?

       Q. Do you remember how to test for whether hydrogen gas has been given off?

2.     Acid + metal carbonate -->salt + carbon dioxide + water

E.g. copper carbonate + hydrochloric acid  --> copper chloride + carbon dioxide + water

                CuCO₃(s)    +          2HCl(aq)       -->       CuCl₂(aq)      +        CO₂(g)       + H₂O(l)

       Q. How could you test for the gas given off?

       Calcium carbonate, CaCO₃, is also a common carbonate which can be found in chalk, limestone and marble.

NOTE! Metal oxides and metal hydroxides DO NOT dissolve in water.

Ø  Naming salts

When naming a salt, it has two parts, like naming a person. The first part comes from the metal in the base or the metal itself, and the second part is from the acid used.

Look at the common examples in the table below:

Acid                                                  Its salts	Example
Hydrochloric acid, HCl    -->            chlorides	sodium chloride, NaCl
Sulphuric acid, H2SO4        -->            sulphates	magnesium sulphate, MgSO4
Nitric acid, HNO3                -->            nitrates	potassium nitrate, KNO3

Ø  Preparing salts

You now know how to make a salt by reacting an acid with an alkali. But an alkali is just a base that dissolves in water.

Q. Do you think bases will react with acids, the same way that acids react with alkalis?

Experiment 1: Making a salt from an acid and a soluble base

1.     Put 25 cm³ of sodium hydroxide in a small beaker. Gradually add hydrochloric acid to the base to make a neutral solution.

How could you test to see when all the acid has reacted?

2.     Put your solution into an evaporating dish and slowly heat till you see crystals. The solution now has to be left for a few days to crystallise.

Tip! Heating the solution slowly will leave you with larger sized crystals.

What is the name of the salt formed?

You will remember that metal oxides and metal hydroxides do not dissolve in water; so when making a salt from either of these two compounds, a couple of steps have to be added to the start of the experiment above:

1.     Add the insoluble base to the acid until no more will dissolve or react. Some of the base will be left behind – this is called excess.

2.     Filter the mixture into a separate beaker to remove the excess base.

3.     Evaporate the water left behind in the filtrate, leaving the salt behind.

Ø  Preparation of insoluble salts

When preparing an insoluble salt, this involves mixing two solutions of two soluble salts. The insoluble salt is then precipitated – a precipitation reaction.

E.g. The insoluble salt, lead carbonate can be prepared by reacting a solution of soluble lead salt (lead nitrate) and a solution of soluble carbonate (sodium carbonate).

Lead nitrate + sodium carbonate   --> lead carbonate + sodium nitrate

    Pb(NO₃)₂(aq)  +     Na₂CO₃(aq)      -->      PbCO₃(s)    +   2NaNO₃(aq)

      We can also show what happens by writing an ionic equation:

                  Pb²⁺(aq) + CO₃^2-(aq)    -->    PbCO₃(s)

     From this we can now see which ions stick together – the partners that make up the individual compounds have swapped partners and formed a precipitate.

The nitrate ions, NO₃^-(aq) and the sodium ions, Na²⁺(aq) have not disappeared; they are still in the solution and do not change so they don’t appear in the ionic equation. They are known as spectator ions.

Remember! When writing chemical or ionic equations, especially in the exam, always remember to include state symbols. Marks can sometimes be lost if they are not included, unless told otherwise.

·       Solid (s)

·       Liquid (l)

·       Gas (g)

·       Dissolved in water (aq)

·       Questions

1.     Copy and complete:

Acid + base (or alkali)   -->   ................ + water

Acid + metal   -->  salt + ................

Acid + carbonate   -->    salt + .............. + ...............

When an acid reacts with a base (or alkali) to ‘cancel’ each other out, what type of reaction do we call this?

2.     Copy and complete the following word equations:

a)     sodium hydroxide + hydrochloric acid   -->  ....................... + water

b)     magnesium + sulphuric acid   --> ............................. + water

c)     ......................... + nitric acid   --> copper nitrate + water

d)     calcium oxide + ........................   --> calcium chloride + water

e)     copper .............. + hydrochloric acid   --> copper ................. + carbon dioxide + water

3.     Explain the difference in procedure when using a soluble base or an insoluble base to make a salt.

4.     Copy and complete the following chemical equation by filling in the missing state symbols:

Na (s) + H₂O (  )   --> NaOH (  ) + H₂(  )