Module 3.1 - The Periodic Table Flashcards
Describe how the English chemist, John Newlands first tried to arrange a table of elements?
In 1863, he noticed that if he arranged the elements in order of mass, similar elements appeared at regular intervals - every eighth element was similar. He called this the law of octaves and he listed the known elements in rows of 7, so that similar elements lined up in columns, but the pattern didn’t always work such as with the transition metals.
Who was the Russian chemist that created the first accepted version of the periodic table?
Dmitri Mendeleev.
Describe how Dmitri Mendeleev produced a better table?
In 1869, he arranged all the known elements by atomic mass but left gaps in the table where the next element didn’t fit so he could keep elements with similar chemical properties in the same group. He predicted the properties of these undiscovered elements that would go in the gaps.
Was Mendeleev’s predictions and leaving gaps beneficial?
When elements were later discovered (e.g. germanium, scandium, gallium) with properties that matched his predictions.
What did Henry Moseley in 1914 arrange the elements by in the modern periodic table?
Increasing atomic (proton) number.
What do elements in the same period have in common with each other?
They have the same number of electron shells.
What is periodicity?
The repeating trends in the physical and chemical properties of the elements across each period.
What do elements within the same group have in common?
Have the same number of electrons in their outer shell.
If an element is in the s-block what does this mean?
If an element is in the s-block it means that its outer shell electrons will be in the s orbital.
Name an example of a group in the p-block?
Group 7.
Give a definition of first ionisation energy?
The first ionisation energy is the energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms.
Is ionising an atom or molecule an exothermic or endothermic process?
Endothermic process.
What are the 3 factors that affect ionisation energy?
Nuclear charge, atomic radius and shielding.
Describe how atomic radius affects the 1rst ionisation energy?
Attraction falls off very rapidly with distance and atomic radius decreases as you go across the period as nuclear charge increases pulling the shells closer. An electron close to the nucleus will be much more strongly attracted than one further away.
Describe how shielding affects the 1rst ionisation energy?
As the number of electrons between the outer electrons and the nucleus increases, the outer electrons feel less attraction towards the nuclear charge. This lessening of the pull of the nucleus by inner shells of electrons is called shielding.
Why does ionisation energy decrease as you go down the group?
> Elements further down a group have extra electron shells compared to ones above and these extra shells means that the atomic radius is larger, which greatly reduces their attraction to the nucleus.
The extra inner shells shield the outer electrons from the attraction if the nucleus.
Describe the reason for the drop in ionisation energy between Groups 2 and 3?
Due to sub-shell structure:
>The outer electron in Group 3 elements is in a p orbital rather than an s orbital.
>A p orbital has a slightly higher energy than an s orbital in the same shell, so the electron is, on average, to be found further from the nucleus.
>The p orbital also has additional shielding provided by the s electrons.
>These factors override the effect of the increased nuclear charge, resulting in the ionisation energy dropping slightly.
Describe the reason for the drop in ionisation energy between Groups 5 and 6?
Due to orbital repulsion:
>In the Group 5 elements, the electron is being removed from a singly-occupied orbital.
>In the Group 6 elements, the electron is being removed from an orbital containing 2 electrons.
3)The repulsion between 2 electrons in an orbital means that electrons are easier to remove from shared orbitals.
When you remove additional electrons, it’s called the.?
Successive ionisation energies.
Why do successive ionisation energies increase as each electron is removed?
This is because electrons are being removed from an increasingly positive ion, and there’s also less repulsion amongst the remaining electrons so are held more strongly by the nucleus.
What structure does diamond, graphite and graphene all have?
Giant covalent lattices.
What are different forms of the same element in the same state called?
Allotrope.
State and explain the properties of diamond?
Because of it’s lots of strong covalent bonds:
>Has a very high melting point.
>Extremely hard.
>Vibrations travel easily through the stiff lattice, so it’s a good thermal conductor.
>Can’t conduct electricity - all the outer electrons are held in localised bonds.
>Won’t dissolve in any solvent.
Describe diamond’s structure?
In diamond, each carbon atom is covalently bonded to four other carbon atoms. The atoms arrange themselves in a tetrahedral shape - its crystal lattice structure. Its the hardest known substance.
Describe silicon’s structure?
Also forms a crystal lattice structure with similar properties to carbon. Each silicon atom is able to form 4 strong covalent bonds.
Why graphite used as dry lubricant and in pencils?
The weak forces between the layers in graphite are easily broken, so the sheets can slide over each other and therefore, graphite feels slippery.
