3.1 Periodic table and energy 7.1-8.3 Flashcards
Atomic radius
The greater the distance between the nucleus and the outer electron the less the nuclear attraction
The force of attraction falls off sharply with increasing distance, so atomic radius has a large effect
Nuclear charge
The more protons there are in the nucleus of an atoms, the greater the attraction between the nucleus and the outer electrons
Electron shielding
Electrons are negatively charged and so inner shell electrons repel outer shell electrons
This repulsion, called the shielding effect, reduces the attraction between the nucleus and the outer electrons
Trend in first ionisation energy down a group
Atomic radius increases
More inner shells so shielding increased
Nuclear attraction on outer electrons decreases
First ionisation energy decreases
General trend in first ionisation energy across period two
Nuclear charge increases Same shell: similar shielding Nuclear attraction increases Atomic radius decreases First ionisation energy increases
Ionisation
Ionisation is the reaction where one electron is removed from an atom.
There will be one electron in the product of the equation.
So we will always have an e− term on the right-hand side of the equation.
Ionisation equation
The first ionisation energy equation of Na:
Na(g) → e− + Na +(g)
The second ionisation energy equation of Mg:
Mg+(g) → e− + Mg 2+(g)
First ionisation energy
The first ionisation energy is the energy required to remove an electron from every atom in a mole of atomic gas, to produce a mole of unipositive gaseous ions.
Second ionisation energy
The second ionisation energy is the energy required to remove an electron from every ion in a mole of unipositive gaseous ions, to produce a mole of dipositive gaseous ions.
Factors effecting ionisation energy
The main factors affecting ionisation energies are the nuclear charge, the distance from the nucleus and electron shielding.
Distance from nucleus- ionisation energy
Electrostatic attraction decreases sharply with distance.
This means that less energy is needed to remove electrons which are further away.
This means that as distance increases, ionisation energy decreases.
In practice, this means that the higher the principal quantum number of an electron, the lower its ionisation energy.
Nuclear charge- ionisation energy
The greater the number of protons in the nucleus, the greater the attraction of the electron to the nucleus.
A greater attraction of the electron means more energy is needed to remove the electron.
This means that ionisation energy is greater.
Shielding- ionisation energy
The greater the number of electrons between the nucleus and the outer electrons, the lower the effective nuclear charge.
This is because the positive charge felt by the electron is reduced by the electrons in between.
This means that the greater the number of electrons, the lower the ionisation energy.
Experimental evidence for shells
Ionisation energies can be used to show the presence of shells in atoms.
Atomic radius and ionisation energy in group two
Atomic radius increases as you go down Group 2.
This is because each extra electron shell is further away.
Ionisation energy decreases as you go down Group 2.
This is because the outer electrons are further away and experience less attraction to the nucleus.
Melting point in group 2
Melting points decrease as you go down Group 2.
This is because the ion cores have larger radii down the group.
The free electrons experience less attraction to the nuclei because of the larger radii.
The bonding is weaker, so the melting point is lower.
Magnesium has an anomalously low melting point.
This is because it has a different crystal structure to the rest of Group 2.
Group two reactions with water
Group 2 metals react with water to form metal hydroxides. For example:
Mg + 2H2O → Mg(OH)2 + H2
Reactivity increases as you go down Group 2. This is because the lower elements have lower ionisation energies.
Beryllium is an exception. If beryllium were to lose two electrons it would be tiny and have a very high charge density. This would make it unstable and so beryllium doesn’t react with water.
Solubility of salts in group two
The solubility of hydroxides increases as you go down Group 2.
Magnesium hydroxide is very insoluble. The phrase to use for this is ‘sparingly soluble’.
The solubility of sulfates decreases as you go down Group 2.
Barium sulfate is completely insoluble.
Trends in reactivity group 2
The reactivity of group 2 metals increases down the group.
The ionisation energies needed decrease down the group because increased shielding from extra shells of electrons and a greater distance from the nucleus makes it easier to remove an outer electron.
Both first and second ionisation energies decrease down the group. The second ionisation energy is greater than the first ionisation energy within an element.
Group 2 reaction with oxygen
Group 2 elements react with oxygen in the air to form metal oxides, For example:
2Mg + O2 → 2MgO
The elements react more readily with oxygen down the group because of the decrease in ionisation energies.
Group 2 reaction with dilute acids
Group 2 metals react with dilute acids to produce a salt and hydrogen. For example:
Ca + H2SO4 → CaSO4 + H2
Reactivity increases down the group because of the decrease in ionisation energies.
The salts produced become less soluble as you go down the group.
Group 2 oxides and water
Group 2 oxides steadily react with water.
The reaction produces a metal hydroxide salt. The salt dissociates into metal ions and hydroxide ions, giving the solution an alkaline pH.
Group 2 hydroxide solubility down the group
Group 2 hydroxides become more soluble down the group.
This means that the solutions become more strongly alkaline the further down Group 2 you go.
Group 2 trends in pH
Because Group 2 oxides get more soluble as you go down Group 2, the pH increases when reacting Group 2 oxides with water as you go down the group.
E.g. Calcium oxide (CaO) will have a higher pH than magnesium oxide (MgO).
milk of magnesia
a suspension of white magnesium hydroxide in water
an antacid
Magnesium hydroxide
used in medicine
slightly soluble in water
Magnesium hydroxide (Mg(OH)2) can neutralise stomach acid (HCl) in the following reaction:
Mg(OH)2(s) + 2HCl(aq) → MgCl2(aq) + 2H2O(l)
There are no changes in the oxidation numbers so this is NOT a redox reaction
Calcium carbonate
used in medicine
Calcium carbonate is an insoluble base.
Calcium carbonate can also be used as an antacid to treat indigestion.
CaCO3(s) + 2HCl(aq) → CaCl2(aq) + 2H2O(l) + CO2(g)
There are no changes in oxidation numbers so this is NOT a redox reaction.
Calcium hydroxide
used in agriculture
Calcium oxide is soluble enough to dissolve in water.
Calcium oxide can be added to water and used to neutralise acidic soil.
Ca(OH)2(s) + 2HCl(aq) → CaCl2(aq) + 2H2O(l)
Halogens
Halogens are located in Group 7 of the periodic table.
They are non-metals.
They tend to form negative ions called halide ions.
The halogens are made up of fluorine, chlorine, bromine, iodine and astatine.
Halogen phases
Fluorine - gas.
Chlorine - gas.
Bromine - liquid.
Iodine - solid.
Halogen formulae
Halogens all tend to exist as diatomic molecules. Fluorine - F2 Chlorine - Cl2 Bromine - Br2 Iodine - I2
Halogen electronegativity
Electronegativity is how strongly an element attracts a bonding pair of electrons.
You should already know that fluorine is the most electronegative element.
Electronegativity decreases as you go down Group 7.
This is because atomic radius increases and so electrons experience less attraction to the nucleus.
Electrons become further away from the nucleus and so experience a lower attraction.
This means that the electronegativity is lower.
Halogen boiling points
Boiling points increase down the group. At room temperature:
Fluorine and chlorine are gases.
Bromine is a liquid.
Iodine is a solid.
This is because elements have more electrons as you go down Group 7.
This causes an increase in London dispersion forces.
Stronger London dispersion forces means greater intermolecular forces.
This means a higher boiling point.
Silver halide solubility in NH3
All silver halides precipitate out of aqueous solution.
When ammonia is added, they may dissolve again.
Solubility in ammonia decreases as you go down Group 7.
Halogen electron configuration
Halogens’ electron configuration has a full s subshell and a p subshell with five out of a possible six electrons.
Formation of halide ions
Halide ions are formed by the addition of one electron to the P subshell, in order to complete both the subshell and the entire outer shell.
Halide ions have a charge of -1.
The oxidation state of halide ions is -1, whereas the oxidation state of halogen atoms is 0.
So, the formation of halide ions involves a reduction.
This reduction can be observed in many redox reactions involving halogens.
Trends in reactivity of halogens
The reactivity of halogens decreases as you go down the group.
This is due to features such as electronegativity and atomic radius.
This means that the further down the group you go, the less readily halogen atoms form halide ions.
Atomic structure of halogens
Atomic radius increases down the group.
This means electrons in the outer shell are further from the pull of the nucleus. -Due to the further distance, there is lower electronegativity (attraction) to pull other electrons in and form ions.
Electron shielding is also higher further down the group.
This also reduces electronegativity and makes it less favourable for ion formation.
Reactivity and displacement reactions of halogens
This trend of reactivity decreasing as you go down the group can be observed in displacement reactions involving halide salts and diatomic halogens.
If the halogen is more reactive in diatomic state (two atoms together) than in halide form, then it will displace it and form its own halide salt.
Disproportionation
Disproportionation is a type of reaction in which a substance is both oxidised and reduced. Chlorine displays this characteristic in reactions involving both water and sodium hydroxide.
Disproportionation of Cl2
Chlorine atoms in diatomic chlorine possess the oxidation state 0.
For disproportionation to occur one atom must be oxidised to a higher oxidation state and one atom be reduced to a lower one.
Reaction between Cl and H2O
The reactions of chlorine and water is a disproportionation reaction.
They react to form a green solution of hydrochloric acid and chloric(I) acid.
The oxidation state of chlorine in hydrochloric acid is -1, it has been reduced.
The oxidation state of chlorine in chloric(I) acid is +1, it has been oxidised.
Cl2 + H2O → HCl +HOCl
Reaction between Cl and cold NaOH
The reaction of chlorine with cold sodium hydroxide is a disproportionation reaction.
They react to form sodium chloride, sodium chlorate(I) and water.
The chloride ion in sodium chloride has been reduced to an oxidation state of -1.
In sodium chlorate(I), the chlorine has an oxidation state of +1.
2NaOH + Cl2 → NaCl + NaClO + H2O
Uses of chlorine
Chlorine is commonly used to treat water sources, for example, swimming pools. Chlorine is used a sanitation method as it kills potentially harmful bacteria.
Limitations of sanitising water with chlorine
Chlorine gas is toxic so the chlorine needs to remain in the water. You often need to use a chlorine stabiliser such as cyanuric acid.
Chlorine can react with hydrocarbons that are found in organic materials (like leaves in a swimming pool).
This can produce trihalomethanes which is potentially carcinogenic (it causes cancer), although there is no conclusive evidence for this claim.
Chlorine sanitation
Chlorine has been used to treat domestic water supplies for decades.
It is effective in killing bacteria and preventing the spread of waterborne diseases such as cholera.
Chlorine reaction with water
The equation for this reaction is:
Cl2(g) + H2O(l) ⇌ HCl(aq) + HClO(aq) ⇌ 2H+(aq) + ClO−(aq) + Cl−(aq)
This reaction is used in water treatment to kill bacteria.
Chlorine reaction with water and in sunlight
The same reaction happens with water. But in sunlight, a further reaction can happen.
In sunlight, the chlorate(I) ion produced will decompose to produce hydrochloric acid and oxygen. The equation for this reaction is:
2HClO(aq) ⇌ 2HCl(aq) + O2(g)
So the overall equation can be written as:
2Cl2(g) + 2H2O(l) ⇌ 4HCl(aq) + O2(g)
Chlorine as water treatment
Chlorine is added to the water supply to try to make it safe to drink.
But the treatment has risks:
Chlorine gas is toxic.
Liquid chlorine on the skin or eyes cause burns.
Water contains trace organic compounds. Chlorine will chlorinate hydrocarbons and many of these chlorinated hydrocarbons are carcinogenic.
Benefits of chlorine as water treatment
Water treatment kills harmful bacteria.
The increased cancer risk is small, but a cholera epidemic would kill thousands.
Some people complain about this forced mass medication, but this prevents thousands of deaths every year and on balance is good for society.
Making bleach with chlorine
When Cl2 is added to sodium hydroxide, you get sodium chlorate(I) and sodium chloride.
This is an example of a disproportionation reaction.
This means chlorine is both oxidised and reduced at the same time.
Cl2 has oxidation state 0, Cl- has oxidation state 1- and chlorate(I) has oxidation state 1+.
Halogen reaction with silver nitrate
If we add silver nitrate to a solution of halide ions, the halides will react with the silver cations:
Ag+(aq) + X−(aq) → AgX(s)
The product of the silver nitrate reaction is a solid.
We call this solid the precipitate.
The colour of this solid depends on what halide is involved:
AgCl → white precipitate.
AgBr → cream precipitate.
AgI → yellow precipitate.
Using the silver halide product, we can test that their ability to redissolve in ammonia is as expected.
AgCl should redissolve.
AgBr should redissolve slowly and needs a lot of ammonia.
AgI should NOT dissolve.
Testing for ions
Qualitative tests can be used to identify the presence of certain anions or cations. Qualitative tests rely on the observation of a characteristic outcome, instead of quantitative measurement.
Test for anions
Anions are negative ions.
Carbonate ions react with acids (e.g. dilute nitric acid) and release carbon dioxide, which characteristically turns limewater milky.
Sulfate ions react with barium ions (e.g. barium nitrate) in solution to produce barium sulfate, a white precipitate.
Halide ions (chloride, bromide, iodide) react with silver ions (aqueous silver nitrate)in solution to produce precipitates of silver chloride (white), silver bromide (cream) and silver iodide (yellow) respectively.
Test for cations
Cations are positive ions.
Ammonium ions are an example of a cation and can be tested for by mixing with warm sodium hydroxide solution.
Ammonia, a noxious gas is given off.
test gas with moist pH indicator paper
ammonia is alkaline and will turn the paper blue
Metallic bonding
strong electrostatic attraction between cations (positive ions) and delocalised electrons
What has a giant metallic lattice structure
All metals
Giant covalent lattices
Atoms bonded by strong covalent bonds
What has a giant covalent structure
B
C
Si
sequence of qualitative tests
- carbonate
- sulfate
- halides
halogen colours in water
chlorine- pale green
bromine- orange
iodine- brown
halogen colours in cyclohexane
chlorine-pale green
bromine- orange
iodine-violet
