Noble Gases

Discovered He 1868; Ar 1894 Ne, Kr and Xe in 1898 Rn 1900 Unreactive because they have a full outer shell of electrons Called Noble then Inert now Noble

Uses Helium Airships as it is lighter than air Not as light as hydrogen [twice as heavy per volume] but does not burn [Hindenberg] Used in atmosphere of deep sea divers Less likely to cause the “bends” Gives the diver “Mickey Mouse” voice because it has such a low density compared to air.

Argon The most common Noble Gas. Used to fill normal incandescent light bulbs to stop them imploding, to reduce evaporation of the filament and because it is unreactive Used as inert atmosphere for welding

Trends in Noble gases Don’t react in general because they have a full outer shell. Down the Group – molecules get larger Makes it easier to form temporary dipoles due to larger electron clouds So van der Waal’s forces are greater More energy required to break these. So MP and BP get higher Same trend in halogens

Octet Rule A chemical rule of thumb Octet Rule says atoms with 8 electrons in their outer shell are stable Atoms tend to combine in such a way that they each have eight electrons in their valence shells, giving them the same electron configuration as the nearest Noble gas

The rule applies to the main-group elements, especially carbon, nitrogen, oxygen, the halogens, and also to metals such as sodium or magnesium. In simple terms, molecules or ions tend to be most stable when the outermost electron shells of their constituent atoms contain 8 electrons.

Limitations Doesn’t allow for H, He or Li Stable with 2 e- in their outer shells - Duet Rule Transition elements - 18 electron rule BF3 which only has 6 e- in its outer shell

How can atoms get a full outer shell? Give away electrons Take in electrons Share electrons Don’t bother getting a full outer shell at all – (Rare but does happen [BF3])

Ionic Bonding

Ar Noble Gas - Also Ne, He, Xe, Kr, Rn Stable or Unreactive 40 18 Protons 18 18+ Electrons 18 18 - Overall Charge ATOM Noble Gas - Also Ne, He, Xe, Kr, Rn Stable or Unreactive Full outer shell 18+ All atoms want to be like this i.e. have a full outer shell - Octet Rule

Na Protons 11 11+ Electrons 11 11- Overall One electron in outer shell 23 11 Protons 11 11+ Electrons 11 11- Overall ATOM One electron in outer shell Wants to get a full outer shell 11+ To have the same pattern as a Noble Gas Loses electron from outer shell

Na 23 11 Protons 11 11+ Electrons 10 11 11- 10- Overall 1+ ATOM 11+

Na+ Where has the electron gone? ION Protons 11 11+ Electrons 10 10- Overall 1+ ION Full outer shell Same pattern as Neon 11+ Stable Now has a positive charge Called an ION Where has the electron gone?

Cl Protons 17 = 17+ Electrons 17 = 17 - Overall = 0 35 17 Protons 17 = 17+ Electrons 17 = 17 - Overall = 0 ATOM 7 electrons in outer shell Wants to get a full outer shell 17+ Takes an electron into its outer shell

Cl- ION Full outer shell Same pattern as Argon 17+ Stable Protons 17 17+ Electrons 18 18 - Overall 1 - ION Full outer shell Same pattern as Argon 17+ Stable Now has a - charge Called an ION

Unlike charges attract Cl- Na+ Na+ Cl-

Summary Sodium atoms lose an electron to become sodium ions Sodium ions have a +ve charge Chlorine atoms gain an electron to become Chloride ions Chloride ions have a − ve charge Opposite charges attract so the Na+ and Cl- come together and stick to each other This is called an ionic bond

General Points Ionic Bonding involves a transfer of electrons It happens when the difference in electronegativity between atoms > 1.7 Metals lose electrons to form positive ions They gain one plus charge for each electron lost Non-metals gain electrons to form negative ions They gain one minus charge for each electron gained Total electron loss must equal electron gain

Keeping Track of Electrons Atoms in the same group [column] Have the same outer electron configuration. Have the same valence electrons. Easily found by looking up the group number on the periodic table. Group 2A - Be, Mg, Ca, etc.- 2 valence electrons

General Points Group I lose 1 e- to become M+ Group II lose 2 e- to become M2+ Group III lose 3 e- to become M3+ Group V gain 3 e- to become X3- Group VI gain 2 e- to become X2- Group VII gain 1 e- to become X- There are exceptions

Group I with Group VII Li F Na Cl K Br Rb I Cs At Fr Formula MX M+1 M

Group II with Group VI Be O Mg S Ca Se Te Sr Ba Po Ra Formula MX M2+ M

Group III with Group V B N Al P Ga As In Sb Tl Bi Formula MX M3+ M X3-

Group II with Group VII Be F Mg Cl Ca Br Sr I Ba At Ra Formula MX2 X

Group III with Group VII X X1- B Al Ga In Tl F Cl Br I At X1- X M+ M2+ M3+ M X1- X Formula MX3

Group I with Group VI Be O Mg S Ca Se Sr Te Ba Po Ra Formula M2X M M+

Group I with Group V Li Na K Rb Cs Fr N P As Sb Bi Formula M3X M+ M X

Group III with Group VI O B S Al Se Ga Te In Po Tl Formula M2X3 X2- X

Group II with Group V Be Mg Ca Sr Ba Ra N P As Sb Bi Formula M3X2 M

Covalent Bonding Peter Jackson

General Rules Involves a sharing of electrons Electrons are shared in pairs In each shared pair one electron comes from each atom This type of bond occurs between elements of Groups IV, V, VI and VII If there is a choice between ionic and covalent - covalent is preferred There are exceptions

H2 Hydrogen H — H — Represents a shared pair of electrons. 1+ 1+ H — H — Represents a shared pair of electrons. Called a Single Bond

Hydrogen - Alternative View H — H H2 Peter Jackson

Chlorine Only draw the outer shell electrons Cl Cl Cl — Cl Cl2

Chlorine – Alternative View Cl Cl Cl — Cl

Oxygen O O Peter Jackson

Oxygen O O O = O O2 Two pairs of electrons shared called a Double Bond

Oxygen – Alternative View O = O O2

N ≡ N Nitrogen N2 Three pairs of electrons shared called a Triple Bond

Nitrogen – Alternative View N ≡ N N2 Peter Jackson

Methane H CH4 H C H H C H H H H 4 single C- H Bonds Peter Jackson

Peter Jackson Ozone O3 O O O O O O

Bonding

Electronegativity SCC Science Dept

Is the relative attraction that an atom in a molecule has for the shared pair of electrons in a single covalent bond between two atoms of the same element In a bond between two identical atoms the pair of electrons are shared equally Chemists have found that in many bonds the pair of electrons are attracted to one of the atoms more than to the other. SCC Science Dept

Hydrogen and Chlorine Electrons attracted to chlorine more than to hydrogen [bigger, but more +ve nucleus] Electrons spend more time near the chlorine than near the hydrogen This gives the chlorine a slightly negative charge δ- delta minus This gives the hydrogen a slightly positive charge δ+ delta plus SCC Science Dept

H Cl SCC Science Dept

Linus Pauling measured the electronegativity of each element and put them in a table Noble gases are not in the table because they do not form bonds H Cl +  – SCC Science Dept

Differences and Bond Type Difference 0 to 0.45 [Pure] Covalent Bond Difference > 0.45 but < 1.7 Polar Covalent Bond Difference is = or > 1.7 then Ionic Bond SCC Science Dept

Trends in Electronegativity Across a period Goes up Bigger nuclear charge Smaller atomic radius [distance from nucleus] Increased effective nuclear charge SCC Science Dept

Increased Atomic Radius - electron is much further from nucleus Down a Group Gets less Bigger nuclear charge But! But! But! But! Increased Atomic Radius - electron is much further from nucleus Increased shielding by more inner electron shells Decreased effective nuclear charge SCC Science Dept

Intra and Intermolecular Forces Intramolecular forces are the forces within a molecule They are essentially Ionic Bonds and Covalent Bonds. They are powerful forces

Intermolecular Forces Forces of attraction between molecules They are generally weak Three types Dipole – Dipole [permanent] Hydrogen Bonding [special type of permanent dipole – dipole] Van der Waals’ Forces [transient or temporary dipole]

What is dipole? Charge caused by the separation of the centres of + and - charge within an atom or molecule Dipole-Dipole [Polar Bonding] caused by difference in electronegativity. Dipole permanent e.g. δ+ C=Oδ- Hydrogen Bonding special dipole-dipole where H bonded to N, O or F. e.g. δ- O-H δ+ in water [1/10th] Van der Waals’ Forces occurs in atoms or molecules where the difference in electronegativity is 0 and there should be no polarity. They are caused by temporary dipoles produced by random movement of electrons . E.g. H-H or He [about 1/1000 of covalent]

Why is NH3 soluble in water but PH3 is not? NH3 is polar, PH3 is not. NH3 forms hydrogen bonds with the water, PH3 does not. Metallic Bonds a lattice of +vely charged nuclei in a sea of electrons

Let us look at a Helium atom + 1. Centre of + and – charges are in the same place [centre] so there is an even distribution of charge

2. Random movement of electrons results in two electrons being on one side of the atom. This side is now slightly negative δ- and the other side is therefore δ + + δ+ δ -

3. Random movement of electrons results in two electrons being on another side of the atom. This side is now slightly negative δ- and the other side is therefore δ + + δ+ δ - + + 4. It is unlikely that the two electrons will be exactly opposite each other so the atom spends most of it time with one part slightly negative δ-and the other side slightly positive δ + This is called a Temporary Dipole

Let us now look at a Neon atom The bigger the atom the more electrons it has so there is more scope for a bigger charge differential between the two sides. This is due to the possibility of more electrons being at one side.

+ 1. Centre of + and – charges are in the same place [centre] so there is an even distribution of charge

+ 2. Random movement of electrons results in more electrons being on one side of the atom. This side is now more slightly negative δ- and the other side is more δ + δ + δ - This is a bigger Temporary Dipole This is why neon has a higher boiling and melting point than helium

What is the effect of these Temporary Dipoles? These atoms haves NO temporary dipoles at the moment. They have a limited attraction to each other

δ - δ + This atom has developed a temporary dipole This atom still has NO temporary dipole δ - δ +

δ - δ + δ - δ + This atom NOW has an induced temporary dipole The attractive force between the two atoms is now greater due to the two dipoles

Since the force of attraction is greater it is more difficult to remove one of the atoms from the other. The heat needed to give one enough energy to evaporate is greater Thus the boiling point is greater than if there were no dipoles Thus neon has a higher boiling point than helium. Melting points are also greater for the same reason

Neon has a higher boiling point than helium because the atoms are bigger and therefore can produce greater temporary dipoles. The boiling point of O2 is higher than H2 for the same reason Molecules with dipoles tend to be soluble in water – like dissolves like

Everday experience of Van der Walls’ Forces

Graphite is a form of carbon [allotrope] Graphite is pencil lead Made up of sheets of carbon atoms in hexagons Carbons are joined to each other by covalent bonds Sheets are held to each other by van der Walls’ forces.

When you put the pencil on the paper the force of attraction between the paper and the graphite is greater than the attraction between the grapite sheets. Thus when you move the pencil the bottom graphite sheet sticks to the paper The pencils slides along and drops to the paper to repeat the process. Pencils can write underwater The american spent millions developing a pen that could write in space The Russians were smart !!!!!!!!!!!!!!! They brought pencils

Paper / graphite attraction Pencil moves Van der Waals’ Forces Since the paper/grapite attraction is greater than the graphite/graphite attraction the bottom layer stays put as the upper layers moves to the side and there is now a black mark on the paper Paper / graphite attraction

[ ] Dative Bonding + H3O+ H2O Special case where one atom supplies both electrons Two examples H2O → H3O+ and NH3 → NH4+ H+ H3O+ H2O Proton [ ] + H O H O H Hydroxonium Ion

[ ] Dative Bonding + NH4+ NH3 Ammonia Ammonium Ion Special case where one atom supplies both electrons Two examples H2O → H3O+ and NH3 → NH4+ Ammonia H+ NH3 NH4+ Proton [ ] + H N H N H H Ammonium Ion