section 2 atomic structure_edited

7/29/2019 Section 2 Atomic Structure_edited

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Atomic Structure

Cambridge A-level Chemistry

Centre of Pre-U Studies

Section

1


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Overview of Section 2

Content:

Describe the atom in terms of

protons and neutrons in the nucleusExplain and describe the electrons

in terms of energy levels, ionisation

energy, atomic orbitals andextranuclear structure


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Defining the Atom

The Greek philosopherDemocritus (460B.C. 370 B.C.) was among the first to

suggest the existence of atoms (from

the Greek word atomos) He believed that atoms were indivisible and

indestructible

His ideas did agree with later scientifictheory, but did not explain chemical

behavior, and was not based on the

scientific method but just philosophy


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Viewing Surface of metal using STM

Blue platinum

Blue Nickel


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Daltons Atomic Theory (experiment based)

3) Atoms of different elements combine in

simple whole-number ratios to formchemical compounds

4) In chemical reactions, atoms are combined,

separated, or rearranged but never

changed into atoms of another element.

1) All elements are composed of

tiny indivisible particles calledatoms

2) Atoms of the same element are

identical. Atoms of any one

element are different fromthose of any other element.

John Dalton

(1766 1844)


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Section 2

Structure of the Nuclear Atom

Review:

Discovery of subatomic

particles.


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Thomsons Atomic Model

Thomson believed that the electronswere like plums embedded in a

positively charged pudding, thus it

was called the plum pudding model.

J. J. Thomson


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Ernest Rutherfords

Gold Foil Experiment - 1911

Alpha particles are helium nuclei -

The alpha particles were fired at a thinsheet of gold foil

Particle that hit on the detecting

screen (film) are recorded


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Rutherfords Findings

a) The nucleus is smallb) The nucleus is dense

c) The nucleus is positively

charged

Most of the particles passed right through

A few particles were deflected VERY FEW were greatly deflected

Like howitzer shells bouncing

off of tissue paper!

Conclusions:


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The Rutherford Atomic Model Based on his experimental evidence:

The atom is mostly empty spaceAll the positive charge, and almost all

the mass is concentrated in a small area

in the center. He called this a nucleusThe nucleus is composed of protons

and neutrons

The electrons distributed around thenucleus, and occupy most of the volume

His model was called a nuclear model


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Section 2

Structure of the Nuclear Atom

One change to Daltons atomic

theory is that atoms are divisible into

subatomic particles:Electrons, protons, and neutrons are

examples of these fundamental

particles

There are many other types of

particles, but we will study these three


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Discovery of the ElectronIn 1897, J.J. Thomson used a cathode ray

tube to deduce the presence of a negativelycharged particle: the electron


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Question May 2002

Sir James Jean who was a great populariser ofscience, once described an atom of carbon asbeing like six bees buzzing around a space thesize of a football stadium.

(a)(i) suggest what were represented by the sixbees in this description.

(ii) explain(in terms of an atom of carbon) whatstopped the bees from flying away from the

space of the football stadium.(iii) what is missing from Jeans description whenapplied to an atom of carbon?


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Atomic particles Section 2(a),(b),(c)

Identify and describe

protons,neutrons and electrons interms of their relative charges and

relative masses.Deduce the behavior of beams of

protons, neutrons and electrons in

electric fields.

Describe the distribution of mass

and charges within an atom


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Mass of the Electron

1916 Robert Millikan determines the mass

of the electron: 1/1840 the mass of a

hydrogen atom; has one unit of negative

charge

The oil drop apparatus

Mass of the

electron is

9.11 x 10-28 g


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Conclusions from the Study

of the Electron:

a) Cathode rays have identical properties

regardless of the element used to

produce them. All elements must contain

identically charged electrons.

b) Atoms are neutral, so there must be

positive particles in the atom to balance

the negative charge of the electronsc) Electrons have so little mass that atoms

must contain other particles that account

for most of the mass


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Conclusions from the Study

of the Electron:

Eugen Goldstein in 1886 observed

what is now called the proton -

particles with a positive charge, anda relative mass of 1 (or 1840 times

that of an electron)

1932James Chadwick confirmedthe existence of the neutron a

particle with no charge, but a mass

nearly equal to a proton


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Summary Table

Therefore

Modern picture of an atom, then, consist of three types of particles-electrons, protons

and neutron.

Electric Charge MassParticle SI (C ) Atomic SI (g) amu Located

Electron -1.602x10-19 -1 9.109x10-28 5.49x10-4 outside nucleus

Proton +1.602x10-19 +1 1.673x10-24 1.0073 in nucleus

Neutron 0 0 1.675x10-24 1.0087 in nucleus


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Subatomic Particles

Particle RelCharge

Rel Mass Location

Electron

(e-

) -1 1/1840 Electron

cloud

Proton

(p+)+1

1

NucleusNeutron

(no) 0 1 Nucleus


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Deflection of atomic particles

Interaction of charged particles in electricfield

Non-interaction of Non-charged particles

in electric field Direction of deflection depends on the

charge of the particle

Magnitude of deflection depends on themass of the atomic particles when themagnitude of charge is the same


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Question 1.Beams of particles traveling at the samespeed from different sources are subjected to an electricfield as shown in the diagram below. A beam of neutronshas already been drawn.

Sketch on the diagram above how beams of each of the following

particles are affected by the electric field: (i) protons; (ii) electrons; (iii)2

1H

Label each of the beams.

Explain briefly the position and shape of each beam: (i) protons;(ii)electrons and (iii) 21H

+-

n


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Section 2(d)

Learning Outcomes

Deduce the numbers of protons, neutrons and

electrons present in both atoms and ions

given proton and nucleon numbers( andcharge)


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Proton Number

Atoms are composed of protons,neutrons, and electrons

How then are atoms of one element

different from another element? Elements are different because they

contain different numbers ofPROTONS

The proton number of an element is

the number of protons(#) in the nucleus

# protons in an atom = # electrons


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Proton Number

Proton number (Z) of an element isthe number of protons in the nucleus

of each atom of that element.

Element # of protons Z

Carbon 6 6

Phosphorus 15 15

Gold 79 79

N l N b (A)


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Nucleon Number(A)

Nucleon number is the number of

protons and neutrons in the nucleus

of an isotope: A = p+ + n0

Nuclide p+ n0 e- A

Oxygen - 10

- 33 42

- 31 15

8 8 1818

Arsenic 75 33 75

Phosphorus 15 3116


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Complete Symbols

Contain the symbol of the element,the nucleon number and the proton

number.

XNucleonNumber(A)

Proton

Number(Z)Subscript

Superscript


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Symbols

Find each of these:a) number of protons

b) number ofneutrons

c) number of

electronsd) Proton number

e) Nucleon Number

Br8035


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Symbols

If an element has an protonnumber of 34 and a nucleon

number of 78, what is the:

a) number of protons

b) number of neutrons

c) number of electrons

d) complete symbol


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Symbols

If an element has 91protons and 140 neutrons

what is thea) Proton number

b) Nucleon numberc) number of electrons

d) complete symbol


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Symbols

If an element has 78electrons and 117 neutrons

what is thea) Proton number

b) Nucleon numberc) number of protons

d) complete symbol


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Note

In a neutral atom , the number of electronsis equal to the number of protons.

In a negative ion (anion), the number of

electrons is greater than the number ofprotons since a negative ion is obtained by

adding electron(s) to a neutral atom.

In a positive ion(cation), the number ofelectrons is less than the number of

protons since a positive ion is obtained by

removing electron(s) from a neutral atom.


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Ions

is an electrically charged particle obtained from an atom or

chemically bonded group of atoms lose or gain electrons.

The charge on an ion is equal to the # of protons minus the #of electrons. An atom that gains extra electrons becomes a

negatively charged ion, called an anion. An atom that loses

electrons becomes positively charged ion, called a cation.

EAZ

number p + number n

number p

+ ? number p - number

E.g. Determine numbers of electrons in Mg2+ cation and the S2-

anion?

Mg2+ 12-number e = +2 number e =10

S2- 16-number e = -2 number e =18


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May/June 2008

Skin cancer can be treated using a radioactive isotope ofphosphorus,3215P . A compound containing the phosphideion 3215P

3- , is wrapped in a plastic sheet, is strapped to theaffected area.

What is the composition of the phosphide ion, 3215P3 ?

P32

15

332

15P

332

15P

Protons Neutrons electrons

A 15 17 18

B 15 17 32C 17 15 17

D 32 17 15


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Section 2.3(e)(II)

Distinguishing Among Atoms

Learning Outcome:

Distinguish between isotopeson the basis of different

numbers of neutrons present


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Isotopes Dalton was wrong about all elements of

the same type being identicalAtoms of the same element can have

different numbers ofneutrons.( see

next slide) Thus, different nucleon numbers.

These are called isotopes.

Isotopes have the same chemicalproperties but different properties(e.gdenstiy, mass)


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IsotopesElements

occur innature as

mixtures of

isotopes.


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Naming Isotopes

We can also put the nucleonnumberafterthe name of the

element:carbon-12

carbon-14uranium-235


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Isotopes are atoms of the same element having

different masses, due to varying numbers ofneutrons.

Isotope Protons Electrons Neutrons Nucleus

Hydrogen1

(protium) 1 1 0Hydrogen-2

(deuterium) 1 1 1

Hydrogen-3

(tritium)

1 1 2


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-

-

i) Define the term isotopes.

[2]

ii) Br has 2 main isotopes. Complete the table below show the atomic structure of these

isotopes. [2]

Isotope Protons Neutrons Electrons

79Br81Br

iii) The percentage composition by mass of the isotopes in part (b) is 50.5% 79Br and

49.5%81

Br. Use this data to calculate the relative atomic mass of bromine to threesignificant figures.

[2]


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Other types of particles

Our understanding of atom is still

in growing. The unknowns are

yet to be discovered.


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Question May/June 2007

John Daltons atomic theory, published in 1808,contained four predictions about atoms.

Which of his predictions is still considered to be

correct

A Atoms are very small in size

B No atom can be split into simpler parts

C All the atoms of a particular element have the

same massD All the atoms of one element are different in

mass from all the atoms of the other elements


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Question31 May/June 2007

Use of the Data Bookletis relevant to this

question.

The technetium-99 isotope( 99Tc) is radioactive

and has been found in lobsters and seaweedadjacent to nuclear fuel reprocessing plants

Which statements are correct about an atom of99Tc?

1. It has 13 more neutrons than protons.

2. It has 43 protons.

3. It has 99 nucleons.


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Orbitals Section 2(f)

Learning Outcomes

Describe the number and relative

energies of the s,p, and d orbitalsfor the principal quantum numbers

1, 2, and 3 and also the 4s and 4p

Describe the shapes of s and porbitals


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What is an orbital?

An atomic orbital is a region of space(3dimensional) round the nucleus in which theprobability of finding a particular electron( in afree atom) is the greatest.

Electrons can occupy four types of orbital, whichdiffer from each other in shape and in theirorientation in space. These are called s,p, d andf orbitals. s orbitals are spherical

p orbitals are dumb-bell-shaped and can be arrangedin different directions.


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Shapes of the orbitals

s,p and d- orbitals


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Question

Sketch the shapes of and spatial distributions of all

the occupied orbitals in nitrogen


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Answers

1s

2s

2pz 2px 2py


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Modern Atomic Theory

The Principal Quantum Number (n)

Represents the energy level of highest

probability.

n = 1

n = 2

n = 3


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Analogy

A ladder

A bookshelf

View video on atom


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What is a shell?

A shell is a group of orbitals that are about

the same distance out from the nucleus.

Shells are numbered starting with the shell

nearest to the nucleus and working outwards

Each successive shell has a different number

of orbitals in it. In the nth shell, there are n

sub-shells, n

2

orbitals and a maximum of 2n

2

electrons.


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No. of shells and No. of electrons

shell No. of

orbitals(n2)

s p d f Max. no. of electrons

in shell( 2n2)

n=1 1 1 2

n=2 4 1 3 8

n=3 9 1 3 5 18

n-=4 16 1 3 5 7 32


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What is a sub-shell?

A sub-shell is a group of orbitals with the

same energy level, but differ in their

orientation in space, e.g. the second shell

(n=2) contains two sub-shells: a sub-shell containing one s orbital, and

a sub-shell containing three p orbitals:

2px,2py,2pz


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Section 2(h)

Learning Outcome

State the electronic configuration of atoms

and ions given the proton number( and

charge)

Wh t i t b l t i


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What is meant by electronic

configurations?

The electronic configuration of an element

describes how the element of its atoms

are arranged in their shells, sub-shells and

orbitals.


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Filling of orbitals

In an atom, the orbitals are filled in order

of increasing energy(starting from 1s)

according to the following rules.


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Electronic Energy Level

An Energy Diagram

n = 1 1s

n = 2

2s ___

2p ___ ___ ___

N = 3

3s ___

3p ___ ___ ___

3d ___ ___ ___ ___ ___

Rules for assigning electrons

1. Occupy lowest levels

first.(Aufbau Principle)

2. Allowed 2 electrons in

each orbital. If 2, must

have opposite spins. [Pauli

Exclusion Principle]

3. Fill sub-level before

moving up.(Hunds Rule of

Multiplicity)

4. 1 electron in each orbital of

the sameenergy(degenerate) before

doubling up. To keep

electron spins the same

before thay are occupied in

pair)


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Electronic Configurations

The Electron Configuration

A short hand method for notating the electronic

structure of atoms.

Write the electron configuration for sodium.

Sodium has 11 electrons. We know this

because it has an atomic number of 11 which

means it has 11 protons and thus must have 11

electrons.

1s22s2p63s1 or [Ne]3s1


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Modern Atomic Theory

An Energy Diagram

n = 1 1s

n = 2

2s ___

2p ___ ___ ___

n = 3

3s ___

3p ___ ___ ___

3d ___ ___ ___ ___ ___

1s22s22p63s1


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Question Oct 2001

Give the full electron configuration of the

following

(i) Mg

(ii) Mg 2+

(iii) O

(iv)

O

2-


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Question

Construct a schematic diagram representing the

number and relative energies of the orbitals of the

first three principal quantum numbers.


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Answers

Energy

3d

3p

3s

2p

2s

1s


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Question

Show the electronic configuration of

nitrogen using the diagram in the previous

answer


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Answer

Energy 2p

2s

1s


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Isoelectronic

Atoms that have the same number of

electrons are known as isoelectronic.


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Oct/Nov 2007

Use the Data Booklet is relevant to this question.In forming ionic compounds, elements

generally form an ion with the electronic

structure of a noble gas.

Which ion does not have a noble gas

electronic structure?

A. I- B Rb+ C Sn 2+ D Sr2+

Representing electronic


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Representing electronic

configuration

By using electrons-in-boxes

Using a noble gas core

Using the s, p, d and f notation

Using energy levels


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Ionisation Energy Sect.2(j)

Learning Outcomes

(i) explain and use of the term ionisation

energy

(ii) explain the factors influencing theionisation energies of elements

(iii) explain the trends in ionisation energies

across a period and down a group of the

Periodic Table.


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Definition of IE

The energy to remove an electron from

each atom/ion in a mole of atoms/ions of

an element in the gaseous state

Refer to data booklet for successive IEenergies

Consecutive IE of a single


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Consecutive IE of a singleelement or atom

Generally increase of IE is due to the

consecutive electron removed because it

is harder to remove electrons from the

more positive ion. Large rise of IE implies that electron are

removed from a different principal shell.

Small increase of IE represent change insubshell.


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Explanation of IE Explain large rise in IE. When all valence shell

electrons are removed, only the inner shellremains. The inner shell experience much lessshielding , resulting in a higher nuclear attractionand therefore lead to a much higher IE big

jump Explain small increase in IE. Electrons in higher

level subshell removed first; after whichelectrons from lower energy subshell removed.

For example, 3p electrons are removed before3s. Lower subshell electron experience slightlyhigher nuclear attraction and therefore slightlyhigher IE small jump


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Factors influencing IE

The size of the positive nuclear charge

The distance of the electron from the

nucleus

The shielding effect by electrons in filled

inner shells


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Explaining the IE trends

Increase across the period

Across the period, nuclear attraction on the

valence electron increases. This is due to the

increase in proton number,Z but a fairlyconstant shielding

Decrease down the group

Down the group, nuclear attraction on thevalence electron decreases. This is due to

great increase in Screening Effect,SE with

atom having more inner shells.


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IE trend ACROSS period

IE depends on the effect of the nuclear attraction on thevalence electron.

Across the period, IE increases since the proton numberincreases but the shielding effect remains fairly constant.

IE do not increase uniformly. Small kink may beproduced, where the IE is lower than normal. 2 possiblereasons(i) electron is removed from the higher energysubshell. Less energy is required to remove p-electron(ii) paired vs unpaired electron is removed. Paired

electron experiences repulsion makes a removal easier(use only in p orbital)

Always write out the electronic configuration to figure outthe reason for non-uniformity.


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IE down the group

IE down the group will decreasedramatically, since the effect of the nuclear

attraction on the valence electron

decreases significantly due to increase innumber of principal quantum electron

shells(inner shells) . Outer shell

experiences much higher shielding effectwhich outweighs(greater weight attached

to the effect) increase of nuclear charge.

SE > Z


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Importance of IE

To predict or confirm the simple electronicconfigurations of elements

An example. the first four IE of an element

are, in kJ per mol: 590,1150,4940,6480.Suggest the Group in the Periodic table to

which this element belongs

Oct/Nov 2007


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The graph show the first thirteen IE ionisation energies for

element X

What can be deduced about element X from the graph?

A It is in the second period (Li-Ne) of the Periodic Table

B It is a d-block element

C It is in the Group II of the Periodic Table

D It is in Group III of the Periodic Table

i) Define the term first ionisation energyand write an equation to represent the first


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) gy q p

ionisation energy of chlorine.

[3]

Definition

ii) Write an equation to show the fourthionisation energy of nitrogen.

iii) State and explain the trend in first ionisation energies down group 2 in the Periodic

Table.

Equation [2]

Equation [2]

[4]

The successive ionisation energies for nitrogen are show below.


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g g

Ionisation number 1st 2nd 3rd 4th 5th 6th 7th

Ionisation energy / kJmol1 1402 2856 4578 7475 9445 53268 64362

i) Explain why the successive ionisation energies show a general increase.

ii) What can be deduced about the electronic structure of the nitrogen atom? Explain

your deductions carefully.

[2]

[4]

i) Define the term isotope.

ii) B h 2 i i C l h bl b l h h i f h


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Mark schemei) Substances consisting of atoms with same atomic number / same number of protons / of the

same element 1but with different mass number / different numbers of nucleons / different neutron numbers

1

ii)

2 marks all correct 1 mark1 row or 2 columns correct

iii) Ar= (79 x 0.505) + (81 x 0.495) 1

= 79.99% 80.0% to 3 sig. figs. 1

Isotope Protons Neutrons Electrons

79Br 35 7935 = 44 35

81Br 35 8135 = 46 35

ii) Br has 2 main isotopes. Complete the table below show the atomic structure of these

isotopes.

iii) The percentage composition by mass of the isotopes in part (b) is 50.5% 79Br and

49.5% 81Br. Use this data to calculate the relative atomic mass of bromine to three

significant figures.

i) Define the term first ionisation energyand write an equation to represent the first


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ionisation energy of chlorine. [5]

ii) Write an equation to show the fourthionisation energy of nitrogen. [2]

iii) State and explain the trend in first ionisation energies down group 2 in the Periodic

Table. [3]

Mark schemei) Definition: Energy required to remove first electron from each atom 1

in one mole of 1

gaseous atoms (to form one mole of gaseous+1 ions) 1Equation: Cl (g)Cl + (g) + e 1 for species, 1 for state symbols

ii) N +3 (g)+4 (g) + e 1 for species, 1 for state symbolsiii) First ionisation energies decrease as group 2 is descended 1

Number of shells and hence atomic radii increase down the group 1

Greater number of shells means shielding increases 1

Overall, the outer shell electrons are further from the nucleus and less stronglyattracted to it 1

These factors are more important than the increase in nuclear charge 1

Maximum 4 marks

The successive ionisation energies for nitrogen are show below.


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Mark scheme

i) Negative electrons are being removed from an increasingly positive ion / there is agreater proton to electron ratio. 1

The attraction between the nucleus and the remaining electrons increases. 1

Also, the repulsion amongst the remaining electrons decreased 1

Maximum 2 marks

ii) There is a large increase in ionisation energy when the 6th electron is removed. 1 The 6th

electron is removed from a new shell 1 that is closerto the nucleus with less shielding. 1 Within a

shell, the increase in ionisation energy is much smaller. 1 So nitrogen

has 5 electrons in outer shell and 2 in shell closer to the nucleus. 1

Maximum 4 marks

i) Explain why the successive ionisation energies show a general increase. [2]

ii) What can be deduced about the electronic structure of the nitrogen atom? Explain

your deductions carefully. [4]

643625326894457475457828561402Ionisation energy / kJmol1

7th6th5th4th3rd2nd1stIonisation number

section 2 atomic structure_edited

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