Atomic structure

Question 1

SIMPLE BACKGROUND ( from cards 1 to 8 )

3 Sub-atomic particles ?

Answer 1

-Proton
-Neutrons
-Electrons

Question 2

Relative mass and Relative charge of PROTON ?

Answer 2

-Relative mass = 1
-Relative charge = +1

Question 3

Relative mass and Relative charge of NEUTRON ?

Answer 3

-Relative mass = 1
-Relative charge = 0

Question 4

Relative mass and Relative charge of ELECTRON ?

Answer 4

-Relative mass = 1/1840
-Relative charge = -1

Question 5

Working out the numbers of protons and neutrons ?

Answer 5

No of protons = ATOMIC NUMBER of the atom

The atomic number is also given the more descriptive name of proton number.

No of protons + no of neutrons = MASS NUMBER of the atom

The mass number is also called the nucleon number.

Question 6

What are isotopes ?

Answer 6

Isotopes are atoms which have the same atomic number but different mass numbers. They have the same number of protons but different numbers of neutrons.

Question 7

Working out the number of electrons ?

Answer 7

no of electrons = no of protons

Question 8

Where are electrons found ?

Answer 8

The electrons are found at considerable distances from the nucleus in a series of levels called energy levels

Question 9

ATOMIC ORBITALS ( from cards 9 to 26 )

Where do electrons inhabit ?

Answer 9

electrons in fact inhabit regions of space known as orbitals

Question 10

What shape is the S orbital ?

Answer 10

s orbitals are spherically symmetric around the nucleus

Question 11

Each orbital has a name.

The orbital occupied by the hydrogen electron is called a ?

Answer 11

1s orbital

Question 12

The “1” represents ?

Answer 12

orbital is in the energy level closest to the nucleus

Question 13

The “s” tells you ?

Answer 13

the shape of the orbital

Question 14

2s orbital ?

Answer 14

This is similar to a 1s orbital except that the region where there is the greatest chance of finding the electron is further from the nucleus - this is an orbital at the second energy level.

Question 15

“Electron density” ?

Answer 15

how likely you are to find an electron at a particular place.

Question 16

At the first energy level, the only orbital available to electrons is the 1s orbital, but at the second level, as well as a 2s orbital, there are also orbitals called 2p orbitals.

Shape of P orbital ?

Answer 16

dumbbell shape , consisting of two lobes , with a nodal plane ( zero probability of finding an electron ) .

Question 17

All levels except for the first level have p orbitals.

At the higher levels the lobes get more elongated, with the most likely place to find the electron more distant from the nucleus.

Answer 17

( just background info )

Question 18

What are the other 2 orbitals ?

Answer 18

d and f orbitals

Question 19

d orbitals ?

Answer 19

At the third level, there is a set of five d orbitals

Question 20

f orbitals ?

Answer 20

At the fourth level there are an additional seven f orbitals

Question 21

Fitting electrons into orbitals ?

Answer 21

Each orbital can only hold 2 electrons.

Question 22

A convenient way of showing the orbitals that the electrons live in is to draw ?

Answer 22

“Electrons-in-boxes”

Orbitals can be represented as boxes with the electrons in them shown as arrows. Often an up-arrow and a down-arrow are used to show that the electrons are in some way different.

Question 23

The order of filling orbitals - the Aufbau Principle ?

Answer 23

Electrons fill low energy orbitals (closer to the nucleus) before they fill higher energy ones.

Question 24

Hund’s rule ?

Answer 24

Where there is a choice between orbitals of equal energy, they fill the orbitals singly as far as possible.

This filling of orbitals singly where possible is known as Hund’s rule. It only applies where the orbitals have exactly the same energies (as with p orbitals, for example) .

Question 25

Why does we apply hunds rule ?

Answer 25

helps to minimise the repulsions between electrons and so makes the atom more stable.

Question 26

When you are using the Aufbau Principle, you fill the 4s orbital before the 3d orbitals. The same thing happens at the next level as well - you fill the 5s orbital before the 4d orbitals .

Answer 26

( background info )

Question 27

ELECTRONIC STRUCTURES ( from card 27 to 35 ) The first period ?

Answer 27

Hydrogen has its only electron in the 1s orbital - 1s1, and at helium the first level is completely full - 1s2.

Question 28

The second period ?

Answer 28

E.g Lithium's electron goes into the 2s orbital because that has a lower energy than the 2p orbitals. Lithium has an electronic structure of 1s22s1. Beryllium adds a second electron to this same level - 1s22s2.

Question 29

You can lump all the inner electrons together using ?

Answer 29

-Short hand using the group 8 noble gases For example: Mg 1s2 2s2 2p6 3s2 [Ne]3s2

Question 30

On the periodic table where is the S , P , D , F block elements ?

Answer 30

( use a periodic table )

Question 31

s- and p-block elements ?

Answer 31

The elements in Group 1 of the Periodic Table all have an outer electronic structure of ns1 (where n is a number between 2 and 7). All Group 2 elements have an outer electronic structure of ns2. Elements in Groups 1 and 2 are described as s-block elements. Elements from Group 3 (the boron group) across to the noble gases all have their outer electrons in p orbitals. These are then described as p-block elements.

Question 32

d-block elements ?

Answer 32

d-block elements are elements in which the last electron to be added to the atom using the Aufbau Principle is in a d orbital.

Question 33

A transition element ?

Answer 33

defined as one which has partially filled d orbitals either in the element or any of its compounds. Zinc (at the right-hand end of the d-block) always has a completely full 3d level (3d10) and so doesn't count as a transition element.

Question 34

2 exceptions for electronic structure ?

Answer 34

-Chromium -Copper

Question 35

Summary Writing the electronic structure of an element from hydrogen to krypton Use the Periodic Table to find the atomic number, and hence number of electrons. Fill up orbitals in the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p - until you run out of electrons. The 3d is the awkward one - remember that specially. Fill p and d orbitals singly as far as possible before pairing electrons up. Remember that chromium and copper have electronic structures which break the pattern in the first row of the d-block.

Answer 35

( background info )

Question 36

IONISATION ENERGY ( from card 36 to 54 ) Defining first ionisation energy ?

Answer 36

Definition The first ionisation energy is the energy required to remove one mole of the most loosely held electrons from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+.

Question 37

This is more easily seen in symbol terms ?

Answer 37

X(g) ---> X+(g) + e- It is the energy needed to carry out this change per mole of X.

Question 38

What state symbol should ionisation energies always be in ?

Answer 38

The state symbols - (g) - are essential. When you are talking about ionisation energies, everything must be present in the gas state.

Question 39

What is ionisation energies measured in ?

Answer 39

Ionisation energies are measured in kJ mol-1 (kilojoules per mole). They vary in size from 381 (which you would consider very low) up to 2370 (which is very high).

Question 40

All elements have ... what ?

Answer 40

All elements have a first ionisation energy -Even atoms which don't form positive ions in test tubes. -The reason that helium (1st I.E. = 2370 kJ mol-1) doesn't normally form a positive ion is because of the huge amount of energy that would be needed to remove one of its electrons.

Question 41

Patterns of first ionisation energies in the Periodic Table What does first ionisation energy show ?

Answer 41

periodicity That means that it varies in a repetitive way as you move through the Periodic Table.

Question 42

Factors affecting the size of ionisation energy ?

Answer 42

-The charge on the nucleus. -The distance of the electron from the nucleus. -The number of electrons between the outer electrons and the nucleus. -Whether the electron is on its own in an orbital or paired with another electron.

Question 43

Explain The charge on the nucleus ?

Answer 43

The more protons there are in the nucleus, the more positively charged the nucleus is, and the more strongly electrons are attracted to it.

Question 44

Explain The distance of the electron from the nucleus ?

Answer 44

Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away.

Question 45

Explain The number of electrons between the outer electrons and the nucleus ?

Answer 45

Consider a sodium atom, with the electronic structure 2,8,1. (There's no reason why you can't use this notation if it's useful!) If the outer electron looks in towards the nucleus, it doesn't see the nucleus sharply. Between it and the nucleus there are the two layers of electrons in the first and second levels. The 11 protons in the sodium's nucleus have their effect cut down by the 10 inner electrons. The outer electron therefore only feels a net pull of approximately 1+ from the centre. This lessening of the pull of the nucleus by inner electrons is known as screening or shielding.

Question 46

Explain Whether the electron is on its own in an orbital or paired with another electron ?

Answer 46

Two electrons in the same orbital experience a bit of repulsion from each other. This offsets the attraction of the nucleus, so that paired electrons are removed rather more easily than you might expect.

Question 47

Which element has highest FIRST IONISATION ENERGY ?

Answer 47

helium

Question 48

Explaining the general trend across periods 2 and 3 ?

Answer 48

The general trend is for ionisation energies to increase across a period. in the whole of period 2, the outer electrons are in 2-level orbitals - 2s or 2p. These are all the same sort of distances from the nucleus, and are screened by the same 1s2 electrons. The major difference is the increasing number of protons in the nucleus as you go from lithium to neon. That causes greater attraction between the nucleus and the electrons and so increases the ionisation energies. In fact the increasing nuclear charge also drags the outer electrons in closer to the nucleus. That increases ionisation energies still more as you go across the period. In period 3, the trend is exactly the same. This time, all the electrons being removed are in the third level and are screened by the 1s2 2s2 2p6 electrons. They all have the same sort of environment, but there is an increasing nuclear charge.

Question 49

Why the drop between groups 2 and 3 (Be-B and Mg-Al) ?

Answer 49

The increased distance results in a reduced attraction and so a reduced ionisation energy. The 2p orbital is screened not only by the 1s2 electrons but, to some extent, by the 2s2 electrons as well. That also reduces the pull from the nucleus and so lowers the ionisation energy. The 3p electron in aluminium is slightly more distant from the nucleus than the 3s, and partially screened by the 3s2 electrons as well as the inner electrons. Both of these factors offset the effect of the extra proton.

Question 50

Why the drop between groups 5 and 6 (N-O and P-S) ?

Answer 50

The screening is identical (from the 1s2 and, to some extent, from the 2s2 electrons), and the electron is being removed from an identical orbital. The difference is that in the oxygen case the electron being removed is one of the 2px2 pair. The repulsion between the two electrons in the same orbital means that the electron is easier to remove than it would otherwise be. The drop in ionisation energy at sulphur is accounted for in the same way.

Question 51

Trends in ionisation energy down a group ?

Answer 51

As you go down a group in the Periodic Table ionisation energies generally fall. nuclear charge is much greater greater distance from the nucleus and more screening.

Question 52

Trends in ionisation energy in a transition series ?

Answer 52

Apart from zinc at the end, the other ionisation energies are all much the same. All of these elements have an electronic structure [Ar]3dn4s2 (or 4s1 in the cases of chromium and copper). The electron being lost always comes from the 4s orbital. As you go from one atom to the next in the series, the number of protons in the nucleus increases, but so also does the number of 3d electrons. The 3d electrons have some screening effect, and the extra proton and the extra 3d electron more or less cancel each other out as far as attraction from the centre of the atom is concerned.

Question 53

The rise at zinc is easy to explain.

Answer 53

In each case, the electron is coming from the same orbital, with identical screening, but the zinc has one extra proton in the nucleus and so the attraction is greater. There will be a degree of repulsion between the paired up electrons in the 4s orbital, but in this case it obviously isn't enough to outweigh the effect of the extra proton.

Question 54

Ionisation energies and reactivity ?

Answer 54

The lower the ionisation energy, the more easily this change happens: X(g) ---> X+(g) + e- You can explain the increase in reactivity of the Group 1 metals (Li, Na, K, Rb, Cs) as you go down the group in terms of the fall in ionisation energy. Whatever these metals react with, they have to form positive ions in the process, and so the lower the ionisation energy, the more easily those ions will form. The danger with this approach is that the formation of the positive ion is only one stage in a multi-step process. For example, you wouldn't be starting with gaseous atoms; nor would you end up with gaseous positive ions - you would end up with ions in a solid or in solution. The energy changes in these processes also vary from element to element. Ideally you need to consider the whole picture and not just one small part of it. However, the ionisation energies of the elements are going to be major contributing factors towards the activation energy of the reactions. Remember that activation energy is the minimum energy needed before a reaction will take place. The lower the activation energy, the faster the reaction will be - irrespective of what the overall energy changes in the reaction are. The fall in ionisation energy as you go down a group will lead to lower activation energies and therefore faster reactions.