AS Chemistry Chapter#1.3 Atomic Structure

8/11/2019 AS Chemistry Chapter#1.3 Atomic Structure

1/14

H. M. Sulman Munir ([email protected]) Bright Future Pakistani Intr. School

O / A Level Chemistry Teacher Doha, Qatar

Notes Chapter 1.3 (Atomic Structure) AS Level 1

Atomic Structure

Atomic structure meanswhichparticles and howthese particles are arranged in an atom.

Q-1Define relative isotopicmass, relative atomicmass, atomic mass unit, relative

formulamass and relative molecularmass. (You can consult your book or notebook but bespecific while using the terminology)

In these definitions do notuse weightinstead of mass.

When the word relativeor mass of single particle(atom, molecule, ion or formula unit)

is there, do not use a unit. Use gmol-1when the words molaror mass of one moleis

there.

Measuring the masses of atoms------------Mass Spectrometer

1. Sample should be vaporized firstso that

particles can move easily in the machine.

2. Ionization: These particles are bombarded

with high speed electrons. These in-

coming electrons knock out one or more

electrons from the particles of sample,

making the particles positively charged.

3. The positively charged ions are now

acceleratedthrough electric field.

4. They passed through a velocity selector,

which allows only those ions to go to the

next chamber, which are moving at the

same velocity. Purpose: we will separate

the ions, next in the magnetic field, on the

basis of the differences in their masses

and charges(m/e).

5.

Magnetic field deflectsthe beam of ionson the basis of their m/e. Higher is the m/e lesser

will be the deflection, shown in diagram of mass spectrometer. So the beam of ionsis the now

separated into groups of ions of different masses. By setting the magnetic, ions of the same m/e

are made to fall on the detector through an outlet. This falling of ions on the detector takes

place turn by turnin groups.

6. The detector detects that how many types of the ions are coming. An electric current is

produced, directly proportional to the ions falling, in the detector. This strength of electric

current give the relative or percentage abundances of the ions.


2/14




7. The current is of very lowstrength, so first current is amplifiedin each case.

8. This current is fed, after amplification, to the recorder.

Mass spectrum: The resultsfrom the mass spectrometer are obtained in the form of a graphcalled mass

spectrum. On x-axism/eof the ions or simply the relative mass (because most of the ions have +1

charge) and y-axis relative abundancesof the ions are taken.

Q-2The fig. opposite shows the mass

spectrum of atomic chlorine.

(a) Label the x-axis and y-axis.

(b)Explain how many isotopes chlorine has.

Suggest a reason for your answer.

(c) Write the formulae of the ions responsible

for the peaks at 35 and 37.

(d)

Find the relative abundance of each isotope.

(You will use a ruler scale)

(e) From the relative abundance, calculate the

percentage abundance of each isotope.

(f) Calculate relative atomic massof chlorine up to four significant figures.

In many exam questions you are given the percentage abundances directly.

Uses andAdvantagesof Mass Spectrometer:

It needs very small amount of the sample for analysis being very sensitive.

It determines the structure formulaeof the compounds discovered or synthesized. So isomers

(they look very similar to each other) can be differentiated.

Homework: Note: For question 1 you are to write only the names of all the 8 parts of mass

spectrometer and for question two read carefully HSW on page 53. Write answers in the notebooks of

questions on page 53 of the book.

Mass spectrums of molecular substances:

If a sample of bromine (Br2) is analyzed in

the mass spectrometer we will get two sets

of peaks as shown. First set is for atomic

ions 79Br+ and 81Br+, other for molecular

ions 79Br79Br+,79Br81Br+and 81Br81Br+.

The two isotopes of bromine have almost

equal relative abundances.

The last peak in molecular spectrum gives

relative molecular mass of the molecule.


3/14




Q-3

Chlorine has two isotopes 35Cl and 37Cl. The

fig. opposite shows the mass spectrum of

molecular chorine.

(a)

Suggest the masses at which the three

peaks appeared.

(b)Write the formulae of the particles

responsible for each peak.

(c)Calculate relative abundances of the

particles. (use ruler scale and statistical

way)

The living organisms contain carbon, which they take from the air in the form of CO 2, as one

of the main elements in them. Carbon has three isotopes 12C, 13C and 14C. Carbon-14 is

radioactive and slowly decays and a living organism maintains (from their foods) its 12C : 14C

in its life. This ratio in a living organism is just equal to this ratio in the air which is constant.

When an organism dies, its 12C : 14C starts increasing. 14C has a very long half-life (the time

required for a substance to become half of its original amount in a reaction, for carbon-14 it

is 5730 years). So by comparing its 12C : 14C when it was alive (it is a known fixed value)

and at the time when we are calculating this ratio to calculate the age of the organism, the age

of the organism can be calculated. Mass spectrometer is usedto find12C : 14C. See fig 1.3.8

also see the table 1.3.4

Similarly uranium-235 has half-life 704 million years and uranium-238 billion years, these

can help us to calculate the ages or the rocks and even the age of the earth.

After first, second and third half-life the substance remains 50%, 25% and 12.5%

respectively and so on.

Revise the concept of isotopes and relative properties of the subatomic particles from

chapter#1.1.

mass abundance = relative mass x relative abundance or percentage abundance

Homework: Write answers in the notebooks of questions on page 55 of the book and question 2 page 59.

The chemical propertiesof the elements depend upon the electronic configurationsof the elements.

Emission line spectrum:

If a gas is heated or electric currentis given to it, a light is given by this gas. When this light is

passed through a prism, it splits to form a spectrum. This spectrum is a characteristicof an

element.


4/14




This is the evidence of the fact that an atom has different energy levels in it. The energy gaps

between the energy levels decrease as we move away from nucleus. The electrons when de-excite

in these shells give light and hence spectrum.

The energy and the color(if visible) depend upon the frequencyof this light. Only a few

frequencies out of million suggest that a few energy levels are there in an atom.

The electron going awayfrom the nucleus absorbs energyand releaseswhen it comes back.

Energy levels and electron shells:

The electrons around the nucleus are arranged in a series of shells, which resemble the layers

of an onion. The number, n, for a shell is also called principle quantum number.

n = 1,2,3,..

Subshells

Shells are split further into subshells which are s,p,d,f.

Shell Subshells

1 1s

2 2s, 2p

3 3s, 3p, 3d

4 4s, 4p, 4d, 4f

The energy order of the subshells of a shell is

s < p < d < f

The number 1, 2,3,.. with the subshells is

the shell number or principal quantum

number.

Arranging the subshells in the increasing order of their energies :

Q-4Rank the following subshells in the increasing order of energies. 3p, 1s, 2s, 3d, 4s.

Orbitals: The subshells contain orbitals in them further. orbital is a 3D region around the nucleuswhere the electrons are most likely to be found.


5/14




Subshell Orbitals

s s

p px, py, pz

d dxy, dyz, dxz, dx2

-y2, dz

2

f f has seven orbitals. (names are complicated, not in your syllabus)

Q-5Write names and number of subshells and orbitals in the fourth shell.

Electron density maps/Orbitals:

According to the modern theorieselectrons and other microscopicparticles moving at very

high speedsbehave like wavesso electron around the nucleus looks like a cloud. Calculations

from quantum mechanics(mathematical models for solving the electrons behavior around

the nucleus) help us to draw electron density maps or probabilitiesof finding the electron

around the nucleus.

The figure opposite shows the electron

density map for 1s, 2s and 3s (1,2 and 3

are the principal Quantum numbers)

around the nucleus. We can notice that the

s orbitalof each quantum number has

same shape (spherically symmetrical), it

is only the size which is different.It may be called as a subshell or an orbital.

Nodeis a plane where the probability of finding the electron is almost zero.

The shapes and orientation of p and d orbitals:


6/14




Electronic configuration in subshells and orbitals:

Before knowing how to write the electronic configurations when the electrons are in their normal state

(ground state), we should learn the following few points.

An orbital like s, px, dxyetc can accommodate two electrons maximum.

So s, p, d, and f subshells can accommodate maximum 2, 6, 10 and 14 electrons respectively.

All orbitals of a subshell have equal energies but the

subshells of a shell have slightly different energies as their

energy arrangement has been shown on page 4.

Electrons first fill the low energy subshells. The order of

their increasing energy is shown on page 4 or, for shortcut,

write each row, given below, twice.

s

p s

d p s

f d p s

After writing this sequence, assign principal quantum number values for each subshell as:

s from 1, p from 2, d from 3, f from 4

Q-6Arrange the subshells, 4s3s3p3d, in the increasing order of their energies.

An electron in an orbital orbits and spins around its axis. The two electrons in an orbital have

opposite spins (two electrons of the same orbital are spin paired---Paulis Exclusion Principle).

This spin of electrons creates a magnetic field. The two electrons in an orbital have opposite

spin and opposite magnetic fields. Clockwise ( ) and anticlockwise ( )

Hunds Rule: As orbitals like px, pyand pzof a subshell like p have equal energies, so one

electronshould be given to each orbital with same spin firstthen the other electron to each

orbital with the opposite spin.

So keeping into mind the above points we can write electronic configurations of the elements as

Q-7 Which subshell will come after each given subshell: 2s, 3p,6s,4f,..

Q-8Write electronic configurations in the above mode (in the boxes) for Carbon, oxygen, fluorine,

Neon, phosphorus, calcium, manganese and zinc. See atomic number from periodic table.


7/14




Q-9

Writing electronic configurations of mono-atomic ions:

Atoms or mono-atomic ions having the equal number of electrons have same electronic

configurations. Ga3+, Ca2+, Ar, Cl-, P3-each has 18 electrons. So their electronic configuration is same

i.e 1s2, 2s2, 2p6, 3s2, 3p6, 4s

Exceptions or anomalies in writing the electronic configurations:

(1)

The expected electronic configuration of 24Cr is 1s2, 2s2, 2p6, 3s2, 3p6, 4s , 3d 3d 3d 3d 3d but

actually 3d contains 5 electrons instead of 4. One electron is promoted/excited (high energy state) from

4s to 3d so it becomes 1s2, 2s2, 2p6, 3s2, 3p6, 4s , 3d 3d 3d 3d 3d . The same behavior is

shown by Cr and other elements of its group.

(2)Similarly the expected electronic configuration of 29Cu should be

1s2, 2s2, 2p6, 3s2, 3p6, 4s , 3d 3d 3d 3d 3d but again in this case one electron is promoted

from 4s to 3d. So its actual electronic configuration is 1s2, 2s

2, 2p

6, 3s

2, 3p

6, 4s

, 3d

3d

3d

3d

3d

Same do the other elements of group of copper.

Reasonfor both above caseslies in the different stabilities of a subshell when it contains different number

of electrons in it. The stability order for a subshell when it contains different number of electrons is:

Full filled subshell > half filled subshell > scarcely filled (any other number of electrons except full

and half). So 3d5and 3d

10are more stable than 3d

4and 3d

9respectively. The answer of whyfor this order is

beyondthe scope of AS syllabus.

(3) To make atoms ions we add or remove electrons normally from the highest energy subshell means the last

one. So keeping this thing into minds the expected electronic configuration of

Fe2+

is 1s2, 2s2, 2p6, 3s2, 3p6, 4s , 3d

3d 3d

3d

3d (two electrons from d are lost) but actually the

electrons are lost from 4s rather 3d. 1s2, 2s2, 2p6, 3s2, 3p6, 4s, 3d 3d 3d 3d 3d

But d-block (later we will discuss) elements of all the periods lose electrons first from s than from d.

This is due to some stability measures; this is again beyond the scope of your syllabus.

Q-10Write electronic configurations of the following atoms or ions (Show the orbitals of last two

subshells and the spins of electrons). Au, Cu+, Mo, Ti3+.

Q-11


8/14




Q-12 In which group of the periodic table an element X with the following electronic configuration is likely to

be found? Explain your answer.

Ionization Energy:

The amount of energy requiredto remove 1 moleof electrons from 1 mole of gaseous atoms

to make 1 mole gaseous ions understandard conditions is called first ionization energy. OR

H=positive

States,which is always gas,must be given.

Equations for the first ionization energies of some elements

H(g) H+(g) + e- O(g) O+(g) + e- Na(g) Na+(g) + e-

Ne(g) Ne+(g) + e- Fe(g) Fe+(g) + e- U(g) U+(g) + e-

Equations for the second ionization energies of some elements

Al+

(g) Al2+

(g) + e-

Cl+

(g) Cl2+

(g) + e-

Sn+

(g) Sn2+

(g) + e-

Q-13Write equations which show the third ionization energies of the following elements.

Chromium, fluorine, nitrogen, potassium

Trend of ionization energies down a group: It decreasesdown a group

The atomic radius increases down a group, nuclear attraction decreases so ionization

energy decreases down a group.

No doubt number of protons in the nucleus increase down a group but effectiveness of the

nuclear charge decreases due to increase in size of atom, so ionization energy decreases.

The number of shells between the nucleus and the outer shell increases (which shield outer

electrons from the attraction of the nucleus, this is called shielding effect), less attraction

between the shell and the nucleus, so ionization energy decreases.

Trend of ionization energies along a period:

It increasesalong a period due to increase in the effective nuclear chargewhich is due to

the number of electronic shells remains same.

Shielding effect remains same.

Due to increase in effective charge, size of the atom decreases.

Exceptions/Anomalies in ionization energies along the periods:

According to the above discussions ionization energy should increase regularly along the period.

But if we have look on the graphs of ionization energies on the next page, there are two

anomalies: while moving from group II and V to group III and IV ionization energy decreases.

Reasons: The two anomalies can be explained on the basis of the relative energies of the

subshells and their stability with different number of electrons.

1.

Electronic configuration of group II ns

Electronic configuration of group III ns np The electron to be removed from group II is spin-paired which is stable, difficult to remove.

The electron to be removed from group III is in 2p which is at higher energy level and unstable

than 2s, so at the cost of less energy electron can be removed.


9/14




2.

Electronic configuration of group V ns

np

np

np

Electronic configuration of group VI ns

np

np

np

The p subshell of group V is half-filled which is more stable than the p subshell of the group VI

which is scarcely-filled and unstable. So electron from group VI elements can be removed easily.

(OR) the electrons in p-subshell of group VI are paired and repulsion between the two makes

removal of one electrons easier.

Q-14Explain why 11Na has smaller value of first ionization energy than 10Ne?

Higher ionization energies:

12Mg has total twelve values of ionization energies. The every successive ionization energy has more

value than the previous. Second ionization energy value is higher than the first because once first

electron is removed the nuclear charge becomes more effective, size of the atom decreases,it gives

stronger pullto the electrons left and hence ionization energy increases.

There is a very large gap of the values between 2ndand 3rdand 10thand 11thvalues of ionization

energies. After the removal of second electron, third electron is to be removed from the new shell i.e

2ndshell. So a lot of energy is required.

Q-15

Q-16


10/14




Q-17

Q-18


11/14




Note: A copy of periodic table is given on page 12.

Electronic configuration and the periodic table

The chemical propertiesof the elements depend upon their electronic configurations.

The elements in the modern periodic table are arrangedin the increasing order of atomic

numbers.

All the elements of a grouphave same number of electrons in their outer shells and

hence the same outer shell electronic configurations.

Elements of a period have equal number of electronic shells e.g all elements of 3rd

period

have three electronic shellseach.

Noble gases have complete outer shells (octets or duplets) and exceptionally high ionization

energies, so nonreactive.

Classification of the periodic tableon the basis of the subshell in which highest energy

electron of the element is put

1) s-Block 2) p-Block 3) d-Block 4) f-Block

s-Block: The elements of group 1, 2 hydrogen and heliumare called s-block elements because

their highest energy electronis filled in the s-subshell. e.g 11Na, group 1, [Ne]3s1. Elements of

group 1 and 2have properties of typical metalsbut they are very much reactiveand have low

melting points. On the other hand hydrogen and heliumare nonmetals.

d-Block: The elements between group 2 and 3because their highest energy electron goes in d

subshell. These elements except zinc group and scandium group are called transition

elements. These have very high melting points (except mercury which is a liquid), relatively less

reactive than s-block metals, can be converted into wires and sheets and give coloured

compounds.


12/14




f-Block: There are two rowseach containing 14 elements at the bottomof the periodic table are

f-block elements. First row is lanthanide seriesin which 4f subshell is filledat the end and the

second row is called actinide seriesin which 5f subshellis filled at the end.

p-Block: The elements of the groups 3-8 (except He) are p block because their highest energy

electron goes in p-subshell. This block contains non-metals (oxygen, halogens etc),

metalloids (silicon and germanium etc) and metals (tin and lead etc) as well.

Q-19How the atomic sizes vary along the (a) periods and (b) groups? Give two reasons for

these trends in each case.


13/14




Elements of the same group have similar chemical properties: It is due to their similar

outer shell electronic structures. How atoms held in covalent bonds? Electrons are shared between two nuclei. Negative

electrons are tightly attracted by the positive nuclei from both sides.


14/14




Arrange the ions in the increasing order of their angle of deflection from the straight path in a

magnetic or electric field.(mass spectrometer)

Transition of electrons involve energy changes

The concept of spectrum should be after the energy levels in the next edition.

Documents

AS Chemistry Chapter#1.3 Atomic Structure