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Chemistry 3.1 - Periodic table > 3.1.1 periodicity > Flashcards

3.1.1 periodicity Flashcards

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1
Q

How was the old period table arranged?

A
  • Mendeleev ordered the first 60 elements in increasing atomic number
  • left gaps for unknown elements and swapped elements that didnt have fitting properties
  • he also organised the elements in groups of similar properties
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2
Q

How is the period table arranged now?

A
  • It is now organised in seven periods and 18 vertical groups
  • ordered in increasing atomic number
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3
Q

What do elements in the same group contain?

A
  • the same number of electrons in their highest energy electron shell
  • all have similar properties
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4
Q

What do elements in the same period have in common?

A
  • The number of the highest energy electron shell in an elements atoms
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5
Q

to do with electronic configuration

What is the trend across a period?

A
  • each period starts with a new higher energy level
  • for example, period 2 starts with 2s then p and period 3 starts with 3s and fills into the p shell
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6
Q

How are elements divided into blocks?

A
  • elements can be divided into s,p,d and f blocks
  • the blocks represent the shell their outer electrons are located in
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7
Q

What is ionisation energy?

A

Ionisation energy is the measure of how easily an atom loses an electron to form ions

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8
Q

What is the definition of first ionisation energy?

A
  • the energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1 plus ions
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9
Q

Write an eqaution for 1st ionisation energy of Helium?

A

He (g) - He+ + e-

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10
Q

What is the definition of ionisation energy?

A

the energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions

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11
Q

What factors affect nuclear attraction and therefore ionisation energy?

A
  • atomic radius
  • nuclear charge
  • sheilding
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12
Q

How does atomic radius affect ionisation energy?

A
  • The grester the distance between the nucleus and the outer electrons, the less the nuclear attraction
  • this means that ionisation energy decreases
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13
Q

How does nuclear charge affect ionisation energy?

A
  • the greater the nuclear charge, the more protons in the nucleus of an atom
  • this increases the attraction between the nucleus and outer electrons
  • the greater the nuclear charge, the greater the ionisation energy
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14
Q

How does sheilding affect ionisation energy?

A

-The shielding effect is the repulsion between the inner electrons and outer electrons, reducing the attraction between the nucleus and outer electron
- this reduces ionisation energy

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15
Q

The second ionisation energy of helium is greater than the first ionisation energy.
Explain why.

A
  • the second ionisation energy of helium is greater than the first as in a helium atom there are two protons attracting two electrons
  • after the first ionisation energy, there are 2 protons attracting one electron
  • ## This increases the nuclear charge therefore increasion the attraction between the outer electron and nucleus so more ionisation energy is needed to remove the second electrons
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16
Q

What is the trend in ionisation energy across a period?

A
  • first ionisation energy increases due to a decreasing atomic radius, increasing nuclear charge which increases nuclear attraction
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17
Q

What is the trend in ionisation energy down a group?

A
  • ionisation energy decreases down a group
  • This is because atomic radius increases, shielding increases due to more inner shells
  • this decreases the nuclear atrraction on the outer electrons
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18
Q

What are the exceptions across period 2?

A
  • there is a increase from lithium to beryllium, a fall from beryllium to boron
  • A rise from boron to carbon to nitrogen
  • there is then a fall to oxygen and a rise to flourine and neon
19
Q

Why is there a fall from beryllium to boron?

A
  • boron has an electronic configuration of 1s2, 2s2 howeber beryllium has an electronic configuration of 1s2 2s2 2p1
  • this means that the outer electron in the p sub shell is easier to remove as it is in a higher energy sub shell, reducing its nuclear attraction
20
Q

Why is there a fall in ionisation energy from nitrogen from oxygen?

A
  • in nitrogen, each electron in the p sub shell is in a different orbital however in oxygen there is an orbital that contains two electrons
  • these two electrons repel eachother which makes it easier to remove an electron from an oxygen atom
  • this reduces ionisation energy
21
Q

What are the exceptions across period 3?

A

-There is a fall from magnesium to Aluminium
- There is a fall from phosphorus to sulfur

22
Q

Why is there a fall from magnesium to aluminium?

A
  • Aluminiums electronic configuration is 1s2 2s2 2p6 3s2 however magnesiums electronic configurstion is 1s2 2s2 2p6 3s2 3p1
  • this means in aluminium, the distance from the nucleus and outer electron increases, increasing shielding
  • this means it is easier to remove an electron from aluminium
23
Q

Why is there a fall in ionisation energy from phosphorous to sulfur?

A
  • sulphur contains two electrons in an orbital, repelling eachother
  • this reduces the ionisation energy of sulfur
24
Q

What is metallic bonding?

A
  • metallic bonding is the strong electrostatic attraction between cations and delocalised electrons
25
Q

What are the properties of metals?

A
  • solid at room temp exempting mecury
  • ## good conducters of electricity and heat
26
Q

Why do metals have these properties?

A
  • the electrostatic attraction between the cations and delocalised electrons requires a lot of energy to overcome
  • metals conduct electricity as it contains delocalised electrons that can carry charges
27
Q

Why are metals insouble?

A
  • although metals are polar, any interaction between metals and solvents would lead to a reaction and not dissolving
28
Q

What are the two structures non metallic elements make?

A
  • simple molecular compounds
  • Giant covalent structures
29
Q

How are simple molecular compounds held together and what are the properties?

A
  • weak intermolecular forces
  • this causes them to have low melting and boiling points
30
Q

How are giant covalent structures held together and what are the properties?

A
  • held by billions of atoms with a network of strong covalent bonds which form a giant covalent lattice
  • have high melting and boiling points
  • insouble in all solvents
  • non - condutors of electricity with some exceptions
31
Q

What giant covalent structures can conduct electricity?

A

-graphene and graphite which are forms of carbon

32
Q

What is the structure of graphite?

A
  • each carbon atom forms three covalent bonds with three other carbon atoms
  • organised into sheets of hexagons, these sheets are arranged in layers and are joined by weak intermolecular forces called van der waals
33
Q

What are properties of graphite?

A
  • soft, slipperly due to its layers that easily slide
  • high melting and boiling point
  • insouble in both water and organic solvents
  • good conductor of electricity
  • low density
34
Q

Why does graphite have these properties?

A
  • delocalised electron which are free to move between the sheets in graphite (good conductor)
  • high melting and boiling points due to its giant covalent structure and strength of covalent bonds
  • insouble in water and organic solvents due the covalent bonds and the attraction between carbon atoms
  • low density due to distance between layers as they are held apart by weak intermolecular forces (van der waals)
35
Q

What is the structure of graphene?

A
  • Graphene is a single layer of graphite which is composed of hexagonal arranged carbon atoms held by strong covalent bonds
36
Q

What are the properties of graphene?

A
  • high melting and boiling point
  • conductor due to the delocalised electron
  • flexible and strong
37
Q

What is the trend in melting point across period 2?

A
  • There is an increase from lithium to carbon, then there is a sharp decrease from carbon to nitrogen
  • From nitrogen to neon, it remains constantly low
38
Q

Why is there an increase in melting point from lithium to carbon?

A
  • This is because lithium and beryllium are giant metallic structures and boron and carbon are giant covalent structures meaning a lot of energy is required to overcome the bonds
39
Q

Why is there a decrease in melting point from carbon to nitrogen?

A

Nitrogen to neon are simple molecular structures held by weak London forces meaning they have low melting points

40
Q

What is the trend in melting point across period 3?

A
  • there is an increase from sodium to silicon
  • there is a decrease in melting point for, silicon to argon
41
Q

Why does the melting point across period three increase from sodium to silicon?

A
  • There is an increase from sodium to aluminium as they are giant metallic structures
  • There is a greater increase to silicon as silicon is a giant covalent structure
42
Q

Why is there a decrease in melting point from silicon to argon?

A
  • Phosphorus to argon are simple molecular structures held by weak London forces meaning they require little energy to overcome its bonds
43
Q
A

Chemistry 3.1 - Periodic table (4 decks)

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