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310417-Unit 1- Higher Chemistry > Periodicity > Flashcards

Periodicity Flashcards

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1
Q

What is the name if group 1 in the periodic table?

A

Alkali metals

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2
Q

What is the name of group 2 in the periodic table?

A

Alkali Earth metals

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3
Q

What is the name of group 7 in the periodic table?

A

Halogens

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4
Q

What is the name of group 0/8 in the periodic table?

A

Noble gases

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5
Q

Describe metallic bonding

A

positive charged ions with delocalised electrons.

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6
Q

Explain why metals have high melting and boiling point.

A

STRONG ATTRACTION force between these positive ions and delocalised electrons. A lot of energy is required to break metallic bonds.

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7
Q

What are monoatomic elements

A

Consist of single atoms, they are not bonded

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8
Q

Why monoatomic elements have a very low melting and boiling point?

A

Monoatomic elements can form very weak forces between the atoms when cooled. These forces are called London Dispersion Forces and are very weak.
They can be easily broken, therefore monoatomic element have a very low MP and BP.

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9
Q

Which electrons can move about easily in the metallic lattice?

A

The electrons from the outershell.

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10
Q

What happen to the melting point of alkalis metals going down the group? (use databook)

A

Melting point decrease as we go down group 1.

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11
Q

Which type of bonding is holds the atoms together in solid Neon?

A

London dispersion forces

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12
Q

At what Kelvin temperature does the solid lattice Neon break up and turn liquid? (Note: 273 K = 0 °C)
(use databook)

A

At -246°C Neon goes from Gas to Liquid. -246°C = 27K

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13
Q

What state of matter would Argon be at -187 °C

A

Liquid

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14
Q

What do the δ+ and δ- signs stand for?

A

Slightly positive and slightly negative

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15
Q

What elements can form covalent network?

A

Carbon, Boron and Silicon

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16
Q

Give an example of a colalent network

A

Diamond, Graphite, Boron and Silicon

17
Q

Why covalent network have high melting point and boiling point?

A

Structures consist of many thousands of atoms joined together by covalent bonds.
High MP and BP as strong covalent bonds must be broken for the solid to melt/boil.

18
Q

What are the diatomic molecules in the first 20 elements of the periodic table?

A

H2; N2; O2; F2; Cl2

19
Q

What is a diatomic molecule?

A

2 atoms joined together by covalent bond.

20
Q

What is the definition of Covalent Radii?

A

Is a measure of the size of an atom.

It is half the distance between the nuclei of 2 covalently bonded atoms.

21
Q

What happen to the atomic radius going down a group and why?

A

In groups, covalent radius increases going down the group.

More electron layers = higher covalent radius

22
Q

What happen to the atomic radius across periods and why?

A

In a period, covalent radius decreases
going across the period.
All atoms have the same number of electron layers
Number of protons increases across the period.
Greater positive charge = greater attraction for electrons
Electron shells are pulled closer => Covalent radius is smaller.

23
Q

Give the definition of ionisation energy.

A

Is the energy required to remove 1 mole of electrons from one mole of gaseous atom.

24
Q

Why the second ionisation energy of sodium is much higher than the first?

A

Because we need to break into a new, complete electron shell which is closer to the nucleus. Attraction between the nucleus is therefore greater.

25
Q

What happen to the ionisation energy going down a group and why?

A

it decreases because there are extra electron shells meaning that the electron (-ve) being removed is further from the nucleus (+ve) so the force of attraction is lower.

26
Q

What happen to the ionisation energy across periods and why?

A

It increases.
The electron is being removed from the same electron shell. Across the period there are more protons (+ve) and the covalent radius is decreasing.
The electron (-ve) being removed is closer to the nucleus, which has more protons (+ve) so the force of attraction is higher.

27
Q

Explain why only the 1st ionisation energy is quoted for hydrogen.

A

Hydrogen have only one electron it can get rid of.

28
Q

Would the melting point of Li be higher or lower that that of sodium? Why?

A

The atoms in a metal are held together by the attraction of the nuclei to electrons which aredelocalized over the whole metal mass. As the atoms increase in size, the distance between the nuclei and these delocalized electrons increases; therefore, attractions fall. The atoms are more easily pulled apart to form a liquid, and then a gas. Li has a smaller covalent radius than Na, so it has more attraction to the nuclei, therefore stronger force and more energy is needed to break these forces. Therefore, Li will have a higher MP than Na.

29
Q

Xenon was the first noble gas to be made into a compound, explain why this would be the case.

A

Xenon have electrons not as tightly bond to the nuclei as the other noble gas, with smaller atomic number. Its inner electrons screen the outer electrons from the nucleus.

30
Q

A potassium ion is larger than a sodium ion because….

A

it has an extra shell of electrons.

31
Q

Explain why the Al3+ ion is much smaller than the P3- ion even though they are very close to each other in the periodic table.

A

Al3+ have 13 protons and 10 electrons and therefore will have 2 shell of electrons, P3- have 15 protons and 18 electrons and therefore will have 3 shells of electrons, therefore P3- will be larger than Al3+.

32
Q

Explain why potassium has a lower first ionisation energy than lithium.

A

K has a nuclear charge of 19+ and electron arrangement 2,8,8,1
Li has a nuclear chare of 3+ and electron arrangement 2,1
Potassium has two extra filled layers of electrons compared with lithium. These extra layers screen the outer electrons from the nucleus. Consequently, the outer electron in K is further from the nucleus and is not as strongly attracted. It therefore takes less energy to remove the outer electron from K than it does to remove the outer electron from Li.

33
Q

Why there is no electronegativity value provided for Argon?

A

Argon is a noble gas and do not form bonds.

34
Q

Al and P are close in the periodic table but P3- ion is much larger than Al3+ ion, why?

A

P3- had an extra occupied energy level compared to the aluminium ion.

35
Q

The P3- ion and the Ca2+ ion have the same electron arrangement but Ca2+ ion is smaller the P3- ion, why?

A

Calcium ion has a higher nuclear charge, so the electrons are pulled closer to the nucleus, making the calcium ion smaller than the phosphorus ion.

310417-Unit 1- Higher Chemistry (1 decks)

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