CHEM SPM Atomic Structure

8/8/2019 CHEM SPM Atomic Structure

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ALevel Chemistry Chapter: Physical Chemistry

Millennium Education Centre 1 Prepared by Anandh

Atomic Structure

Contents:

1. The structure of the atoma. The nucleus

b. Arrangement of electrons around the nucleus2. Ionisation energy

a. Ionisation energy as evidence for sub-shells3. Chemical bonding

a. Ionic and covalent bondingb. Ionic bondingc. Covalent bonding

4. Shapes of moleculesa. The electron pair repulsion theory

b. How to calculate the shape of a moleculec. Distorted shapes due the presence of lone pairs


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The structure of the atom

The nucleus

1. The protons and neutrons in each atom are tightly packed in a positively charged nucleus.Negatively charged electrons move around the nucleus. The number of protons in a

nucleus defines the type of atom and is the same as the atomic number. The number of

neutrons is found by subtracting the atomic number from the mass number. In an atom

because there is no overall charge the number of electrons equals the number of protons.

2. In chemical reactions the nucleus remains unchanged.3. Geiger and Marsden bombarded a thin gold foil with a beam of alpha particles.4. Most of the particles passed through the foil without deflection and were detected by a

flash of light when the alpha particle struck a zinc sulphide screen, surrounding the gold

foil.

5. A few were deflected and some of these were deflected at angles greater than 900,suggesting they had been repelled by large positive charges within the foil - nuclei of

atoms of gold.

Arrangement of electrons around the nucleus

1. The Bohr model of electrons arranged around a nucleus. The electrons are in certainenergy levels and each energy level can hold only up to a maximum number of electrons.

2. This is summarised in the table below:Energy level or 'shell' Max no of electrons

1st

2

2nd

8

3rd

18

4th

32

5th

50


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3. A sodium atom containing 11 electrons has an electron arrangement of 2,8,1. Twoelectrons filling the first shell, eight electrons filling the second shell and one electron in

the outer third shell.

4. However, these models of electron arrangement are simple and a more advance done cannow be used. It is possible to break these energy levels into sub-shells.

5. Electrons are impossible to locate exactly at any one time. It is however, possible toindicate a region or volume where the electron is most likely to be found. This region is

called an Orbital.

6. Each orbital is capable of holding a maximum of 2 electrons. Orbitals can be divided intos, p, d, andftypes. Each type has its own characteristic shape.

7. The shape ofs and p orbitals are shown below:


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8. The first energy level holds a maximum of 2 electrons in one s type orbital called 1s.There are nop, d, orforbitals available at this energy level.

9. The second energy level consists of ones type orbital and threep type orbitals: 2s, 2px,2py, 2pz.

10.There are 3p orbitals of identical energy, one along the x-axis, one along the y-axis andone along the z-axis.

11.These four orbitals can hold a total of 8 electrons (i.e. 2 electrons each). There are no 2dor 2forbitals.

12.The third energy level consists of: one s type orbital, three p type orbitals and 5 dtypeorbitals. These nine orbitals can hold a maximum of 18 electrons altogether (i.e. two

electrons each).

13.There are sevenftype orbitals holding a maximum of 14 electrons in total.14.When filling the available orbitals with electrons two important principles should be

followed:

a. Electrons fill the lowest energy orbitals first and the other orbitals in order ofascending energy. It is incorrect to assume that an energy level is always completely

filled before electrons enter the next energy level. The order of filling orbitals as

shown below is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p.

b. Where there are several orbitals of exactly the same energy e.g. three 2p orbitals,electrons will occupy different orbitals whenever possible.

15.For example: nitrogen is 1s2 2s2 2px1 2py1 2pz1 and not 1s2 2s2 2px2 2py1.16.This principle is Hund's rule. When an orbital only contains 1 electron then this electron

is said to be unpaired.

a. The small number above the orbital refers to the number of electrons in the orbital:1s

2means 2 electrons in a 1s orbital.

b. The electron arrangements are sometimes abbreviated. For example: the electronarrangement for calcium may be written as 1s

22s

22p

63s

23p

64s

2.


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Ionisation energy

Ionisation energy as evidence for sub-shells

1. Ionisation energy is a measure of the ease in which atoms lose electrons and becomepositive ions.

2. The first ionisation energy is the energy required to remove one electron from each atomof a mole of gaseous atoms.

M(g) - e- M+(g)

3. Further electrons may be removed giving successive i.e.:M

+(g) - e

- M

2+(g)

4. This energy is usually quoted in units ofkilojoules per mole (kJ mol-1).5. Energy is required to remove an electron from any atom because there is an attractive

force between the nucleus and the electron being removed which has to be overcome.

6. The value of the first ionisation energy depends upon: a. The effective nuclear charge

b. The distance between the electron and the nucleusc. The 'shielding' produced by lower energy levels

7. Shielding involves the repulsion between electrons in inner, filled orbitals and electronbeing removed from the outer orbital.

8. The graphs of atomic number against first ionisation number show that across each periodthere is an increase in ionisation energy.


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9. Beryllium (Group II) has an extra electron and proton compared with lithium. The extraelectron goes into the same 2s orbital. The increase in ionisation energy (I.E.) can be

attributed to the increased nuclear charge.

10.The ionisation energy of Boron is less than that of Beryllium because in Boron there is acomplete 2s orbital. The increased shielding of the 2s orbital reduces the ionisation

energy.

11.Similarly, the I.E. of Oxygen is less than that of Nitrogen because the extra electron isshielded by the half-filled 2p orbital.

12.The break in the graph between N-O can be explained by the increased repulsionproduced when two electrons are in the same orbital.

13.Within a group the first I.E. decreases down the group as the outer electron becomesprogressively further from the nucleus. Also there is more shielding because of the extra

filled orbitals.

14.The graph below shows the successive I.E. for sodium:

15.The electronic structure for sodium is 1s2 2s2 2p6 3s1. The energy required to remove thefirst electron is relatively low. This corresponds to the loss of one 3s electron.

16.To remove the second electron needs a much greater energy because this electron iscloser to the nucleus in a 2p orbital.

17.There is a steady increase in energy required as electrons are removed from 2p and then2s orbitals.

18.The removal of the tenth and eleventh electrons requires much greater amounts of energy,because these electrons are closer to the nucleus in the 1s orbital.


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19.Finally an alternative way of expressing electron configuration ass, p,d, andfis to usebox notation as shown below for silicon:

Chemical bonding

Ionic and covalent bonding

1. Theories of chemical bonding are based on the knowledge that:a. Metallic elements from Groups I, II, III tend to lose electrons and form positive ions

that have a noble gas configuration.

b. Non-metallic elements in Groups VI and VII gain electrons top form negative ionswith a noble gas configuration.

c. Elements in groups IV and V do not form charged ions.d. Noble gases do not form chemical bonds. A full outer shell of electrons has extra

stability.

Ionic bonding

1. This involves the transfer of electrons from metal atoms to a non-metal atom to formcharged ions.

2. The oppositely charged ions are held together by electrostatic attractions.


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3. Formation of cations is governed by ionisation energies, with Group I elements formingions most readily and Group III elements forming ions with difficulty. Group IV elements

never form ions because the ionisation energy is too great.

4. Formation of anions is governed by electron affinities.This is the energy change involvedwhen a mole of uni-negative ions from a mole of gaseous atoms. Group VII and Group

VI both form negative ions. Group VII form ions more readily due to greater electron

affinity.

Covalent bonding

1. A covalent bond is made when atoms share one or more electrons to form a molecule. Asingle covalent bond is made when each atom donates one electron to the bond, it is also

possible to form double and triple bonds where two and three electrons are donated.

2. The two atoms come close together so that their outer orbitals overlap. Both nuclei areattracted to the shared pair of electrons and this attraction binds the atoms together.

3. Each atom has been stabilised as it gains a full outer shell (2, 8, 8).


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Co-ordinate or dative covalent bonding:

1. In a normal covalent bond, each atom donates one electron to the shared pair. In a co-ordinate bond electrons come from the same atom.

Shapes of molecules

The electron pair repulsion theory

1. In 1940 Sidgwick and Powell pointed out that the shape of molecules could be explainedin terms of electron pair repulsions.

2. Electron pairs whether bonding or non-bonding repel each other and will arrangethemselves in space to be as far apart as possible. Hence, the shape of a molecule is

related to the number of outer electron pairs.

3. Here are the different structures:

y No. of outer electrons pairs:2y Bond angle: 1800y Example:BeCl2

y No. of outer electrons pairs: 3


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y Bond angle: 1200y Example:BF2,BCl3

y No. of outer electrons pairs: 4y Bond angle: 1090 and 270y Example: CH4

y No. of outer electrons pairs: 5y Bond angle: 1200 and 900y Example: PF5, PCl5

y No. of outer electrons pairs: 6y Bond angle: 900y Example: SF6

Valence shell electron pair repulsion theory(VSEPR):

1. Valence shell electron pairs are arranged to minimise repulsions between themselves.2. Order of Repulsion Strength: Lone pair - lone pair lone pair - bonding pair bonding

pair - bonding pair


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3. When there are 5 or more electrons pairs neglect repulsion between electron pairs at anangle greater than 90

0

4. Basic geometry is still determined by number of electron pairs.5. Lone pair repulsion Double bond repulsion Bonding pair repulsionHow to calculate the shape of a molecule

1. Example: CH4

a. How many electrons around the central atom C = 4b. How many valence electrons shared by H = 4c. Number of bonding pairs = 4 + 4 / 2 = 4d. Number of lone pairs = 0e. Total number of electron pairs = 4f. Arrangement = Tetrahedral

2. The same calculation can be used for all molecules with single bonds.Distorted shapes due the presence of lone pairs

Ammonia

1. This molecule has three bonding pairs of electrons and one lone pair. Although, based onthe tetrahedral shape, due to the extra repulsion from the presence of a lone pair, the 3 N-

H bonds bend further away from the lone pair, in order to minimise the repulsion.

2. The shape is described as pyramidal and has a bond angle of 1070.


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Water

1. This molecule consists of two lone pairs and two bonding pairs. Again, these repel eachother towards the corners of the tetrahedron, due to the extra repulsion from the presence

of two lone pairs. This shape is described as bent and has a bond angle of 1040.

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