Module 3.1 The Periodic Table Flashcards
Before the periodic table
Aristotle believed the world was made up of 4 elements (earth, water, air and fire) - similar to solid, liquid and gas - “fire” represented weird things like plasma
Materials were extracted (some from ores) - deepening knowledge of how substances behaved
Antoine-Laurent de Lavoisier
1789 - produced first modern chemical textbook - contained a list of elements “substances that could not be broken down further”
He devised a theory about the formation of compounds
His list distinguished between metals and non-metals
His list included mistakes e.g. compounds, mixtures, light and heat
Jöns Jakob Berzelius
1828 - published a table of atomic weights
He determined the composition by mass of many compounds
He introduced letter based symbols for elements (previously signs were used)
Johann Wolfgang Döbereiner
Noticed certain groups of 3 elements (“triads”) ordered by atomic weight would have a middle element with the weight and properties that were roughly the average of the other 2 elements
John Newlands
Created a periodic table in order of relative atomic mass
1865 - suggested elements show similar properties to the element 8 places after - “law of Octaves”
Dmitri Mendeleev
1869 - published his periodic table
Elements ordered by atomic masses and also periodically
Gaps helped predict new elements
Henry Moseley and Glenn Seaborg
1913 - Moseley determined the atomic number of all known elements
Moseley modified Mendeleev’s table so the elements were organised by atomic no.
He corrected the order in some instances
Seaborg discovered the transuranic elements and placed the actinides below the lanthanides
Periodicity
The trend in properties that is repeated across each period
The 4 blocks of the periodic table
s-block
d-block
p-block
f-block
First ionisation energy
The energy required to remove one electron from each atom in one mole of the gaseous element to form one mole of gaseous 1+ ions
Factors affecting ionisation energy
Atomic radius
Nuclear charge
Electron shielding
Each successive ionisation energy is higher than the one before. Why?
As each electron is removed, there is less repulsion between the remaining electrons and each shell will be drawn slightly closer to the nucleus
The positive nuclear charge will outweigh the negative charge each time an electron is removed
As the distance of each electron from the nucleus decreases slightly, the nuclear attraction increases. More energy is needed to remove each successive electron
Why ionisation energy generally increases across a period
Decrease in atomic radius - bc of increased nuclear charge
The number of protons increase - higher attraction on the electrons
Electrons are added to the same shell - outer shell drawn inwards slightly
Same no. of inner shells - shielding barely changes
Attraction between the nucleus and the outer electrons increases - more energy needed to remove an electron
Small decrease in the first ionisation energy between the group 2 and group 13 elements
Group 13 elements have their outermost electrons in a p-orbital
Group 2 elements have their in an s-orbital
p-orbitals have slightly higher energy than s-orbitals - marginally further from the nucleus
Electrons in these orbitals are slightly easier to remove = lower I.Es
Small decease in the first ionisation energy between group 15 and 16
As you move from group 13 towards group 18, outer electrons are found in p-orbitals
In groups 13,14 and 15, each of the p-orbitals contains only a single electron
In group 16, however, the outermost electron is now spin-paired in the Px orbital
Electrons that are spin paired experience some repulsion - making the first outer electron slightly easier to remove, so the 1st I.E is slightly lower
The first ionisation energy decreases as you move down a group. Why?
The no. of shells inc. - distance of the outer shell electrons from the nucleus inc. - weaker force of attraction on the outer electrons
More inner shells - inc. shielding - weaker attraction
No. of protons inc. - the resulting inc. attraction is far outweighed by the inc. in distance and shielding
Attraction dec. as you move down a group - less energy needed to remove an electron
Metallic bonding structure
Cations fixed in the lattice
Delocalised electrons - can move
Giant metallic lattices properties: High melting and boiling point
Strong electrostatic attraction between the positively charged metal ions and the negative delocalised electrons
A high temperature is needed to overcome the metallic bonds
Giant metallic lattices properties: good conductors
Delocalised electrons can move freely within the lattice
This allows metals to conduct electricity, even when in the solid state
Giant metallic lattices properties: Malleability and ductility
Delocalised electrons can move - the metallic structure has a degree of “give” - allows atoms or layers to slide past each other
Trend in melting points in the periodic table
Group 1 to 14:
Melting points increase steadily
Giant structures
For each successive group:
If metallic lattice - nuclear charge increases, no. of electrons in outer shell inc., stronger attraction
If giant covalent lattice - each successive group has more electrons with which to form covalent bonds
Group 14-15:
Sharp dec. in melting point
Simple molecular substances
Relatively weak intermolecular forces
Group 15-18:
Relatively low melting points
Simple molecular structures
Graphene
2D giant lattice
One carbon atom thick
Interlocking hexagonal carbon rings
V strong
V light
Can conduct electricity
Many uses in the field of nanotechnology
Physical properties of the group 2 elements
Reasonably high melting points and boiling points
Low densities
Form white compounds
Reactivity of the group 2 elements
Reactive metals
Strong reducing agents
Group 2 elements are oxidised in their reactions
Form 2+ ions
Reactions between group 2 elements and oxygen
React vigorously
Redox reaction
Ionic oxide formed
Reactions between group 2 elements and water
Form hydroxides (except beryllium) - general formula M(OH)2
Hydrogen gas formed
Moving down the group = more vigorous reactions
Redox reaction - the metal is oxidised and one hydrogen atom from each water molecule is reduced
Reactions between group 2 elements and dilute acids
All except Be react with dilute acids to form a salt and hydrogen gas
Move down group = more vigorous
Redox
Solubility of group 2 metal hydroxides
Inc. down the group - will release more OH- ions, will make more alkaline solutions - higher pH
The oxides, carbonates and hydroxides of group 2 metals are
Basic
They’ll react with acids to form a salt and water
Uses of calcium hydroxide
Ca(OH)2
Used by farmers and gardeners as “lime” - reduces the acidity levels of soil
Use of Magnesium hydroxide
Mg(OH)2
“Milk of magnesia”
Used to treat indigestion
Neutralises excess stomach acid to produce a salt and water
Use of group 2 metal carbonates
Building materials
A major drawback of using group 2 metal carbonates as building materials
They react readily with acids so corrosion will occur
E.g. CaCO3 (s) + 2HCl (l) –> CaCl2 (aq) + H2O (l) + CO2 (g)
Properties of the halogens
Low melting points
Low boiling points
Diatomic
Trend in boiling points of the halogens
As you go down the group, the boiling point increases
As you go down the group, the physical state changes from gas, to liquid, to solid
Because each successive element has an extra shell - higher level of London forces between molecules
Reactivity of the halogens
V reactive
V electronegative
Strong oxidising agents
During reactions they form -1 ions and obtain noble gas configuration
The reactivity and oxidising power of the halogens decreases down the group. Why?
The atomic radius increases - less nuclear pull
Electron shielding increases
The ability to gain an electron in the p sub-shell and form a 1- ion decreases
What do you observe when Cl- reacts with acidified silver nitrate solution?
A white precipitate forms
What do you observe when Br- reacts with acidified silver nitrate solution?
A cream precipitate forms
What do you observe when I- reacts with acidified silver nitrate solution?
A yellow precipitate forms
What do you observe when Cl- reacts with dilute ammonia?
The precipitate dissolves to give a colourless solution
What do you observe when Br- reacts with dilute ammonia?
The precipitate doesn’t dissolve
What do you observe when I- reacts with dilute ammonia?
The precipitate doesn’t dissolve
What do you observe when Br- reacts with concentrated ammonia?
It goes colourless and the precipitate dissolves
What do you observe when I- reacts with concentrated ammonia?
The precipitate doesn’t dissolve
NaCl (aq) + AgNO3 (aq) –>
AgCl (s) + NaNO3 (aq)
NaBr (aq) + AgNO3 (aq) –>
AgBr (s) + NaNO3 (aq)
NaI (aq) + AgNO3 (aq) –>
AgI (s) + NaNO3 (aq)
Reaction of chlorine with water (used in water purification)
Cl2 (aq) + H2O (l) –> HClO (aq) + HCl (aq)
Disproportionation! Chlorine has been both oxidised and reduced
Reaction of chlorine with cold dilute aqueous sodium hydroxide to make bleach
Cl2 (aq) + 2NaOH (aq) –> NaCl (aq) + NaClO (aq) + H2O (l)
Disproportionation! Chlorine has been both oxidised and reduced
Testing for carbonate ions
Method:
Add a dilute strong acid
Collect any gas and pass through limewater
Positive test observations:
Fizzing
Colourless gas produced
Limewater is turned cloudy by the gas
Testing for sulfate ions
Method:
Add dilute HCl and barium chloride
Positive test observations:
The white ppt of barium sulfate is produced
Testing for ammonium ions
Method:
Add sodium hydroxide solution and warm v gently
Test any gas produced with red litmus paper
Positive test observations:
Ammonia gas turns red litmus paper blue
