Module 3.1 The Periodic Table Flashcards Preview

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Flashcards in Module 3.1 The Periodic Table Deck (53)
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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



The trend in properties that is repeated across each period


The 4 blocks of the periodic table



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



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



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


*They'll react with acids to form a salt and water*


Uses of calcium hydroxide


Used by farmers and gardeners as "lime" - reduces the acidity levels of soil