Topic 2 Atomic structure

IB CHEMISTRY

2.1 The nuclear atom

• Atoms contain a positively charged dense nucleus composed of protons and neutrons (nucleons). • Negatively charged electrons occupy the space outside the nucleus.• The mass spectrometer is used to determine the relative atomic mass of an element from its isotopic composition.• Use of the nuclear symbol notation to deduce the number of protons, neutrons and electrons in atoms and ions.• Calculations involving non-integer relative atomic masses and abundance of isotopes from given data, including mass spectra.

OBJECTIVES

Atomic Structure

Atoms are very small ~ 10-10 metres

All atoms are made up of three sub-atomic particles:protons, neutrons, and electrons

• The protons and neutrons form a small positively chargednucleus

• The electrons are in energy levels outside the nucleus

Label this diagram:

• The actual values of the masses and charges of the sub-atomic particles are shown in your data booklet:

• A meaningful way to consider the masses of the sub-atomic particles is to use relative masses

Subatomic particles

• Element (X)

• Atomic number (Z) is the number of protons in the nucleusof an atom. It is also known as the proton number. No. ofprotons always equals the no. of electrons in any neutralatom of an element.

• Mass number (A) is the sum of the number of protons andthe number of neutrons in the nucleus of an atom. Someperiodic tables have Z above A. Remember A will always bethe biggest number.

Problem:

Calculate the number of protons and neutrons in:

Z = number of protons = 17 protons

A = number of protons and neutrons = 35

Number of neutrons = A - Z= 35 - 17 = 18 neutrons

Isotopes

Isotopes are atoms of the same element with the same atomic number, but different mass numbers, i.e. they have different numbers of neutrons.

Each atom of chlorine contains

the following:

Cl Cl35

17

37

17

17 protons

17 electrons

18 neutrons

17 protons

17 electrons

20 neutrons

The isotopes of chlorine are often referred to as chlorine-35 and chlorine-37

Some Isotopes of Carbon

Properties of isotopes• Isotopes of an element have the same chemical

properties because they have the same number ofelectrons. When a chemical reaction takes place, it isthe electrons that are involved in the reactions.

• However isotopes of an element have the slightlydifferent physical properties because they have differentnumbers of neutrons, hence different masses.

• The isotopes of an element with fewer neutrons willhave:

• Lower masses

• Faster rate of diffusion

• Lower densities

• Lower melting and boiling points

Radioisotopes

Radioisotopes are isotopes that have unstable nuclei andtherefore emit radiation when then break up.

• Radioisotopes are very useful in society:– 14C is used in radiocarbon dating14C is used in radiocarbon dating– Detecting gas leaks– Industrial quality control– 60Co is used in radiotherapy– 131I and 125I are used a medical tracers– Nuclear power

• Radioisotopes can also be very dangerous to living things:– Radioactive contamination of the environment– Radiation poisoning

Industrial use: detecting blockages in underground pipes

A radioactive isotope which is a gas gets passed down the pipe, where it concentrates the blockage is present.

Industrial use: Quality Control

The radioactive isotope is used as a source of radiation and the amount penetrating the material gives a measure of it’s thickness.

Radiotherapy

A cobalt-60 source can be rotated around the patient. The gamma rays emitted are focussed on the tumour. Healthy surrounding tissue receives a much smaller dose. The cells in the tumour are damaged while surrounding tissue is not.

A radioactive sample can be swallowed. The chemical chosen will be one that concentrates in a particular area. For example, cancer of the thyroid can be treated using iodine-131.

Mass Spectrometer

When charged particles pass through a magnetic field, the particles are

deflected by the magnetic field, and the amount of deflection depends upon

the mass/charge ratio of the charged particle.

Problem1: Determine the relative atomic mass of boron from

the following spectrograph:

m/z value 11 10

Relative abundance %

18.7 81.3

Ar of boron = (amu1 x %1) + (amu2 x %2)

total %

= (11 x 18.7) + (10 x 81.3)

(18.7 + 81.3)

= 205.7 + 813

100

Ar = 1018.7 = 10.2

100

Problem 2: A mass spec chart for a sample of neon shows that it contains 90.9% 20Ne, 0.17% 21Ne, and 8.93% 22Ne. Calculate the relative atomic mass of neon.

Ar of neon = (amu1 x %1) + (amu2 x %2) + (amu3 x %3)

total %

= (90.9 x 20) + (0.17 x 21) + (8.93 x 22)

100

= 1818 + 3.57 + 196.46

100

Ar = 2018.03 = 20.18

100

2.2 Electron configuration

• Emission spectra are produced when photons are emitted from atoms as excited electrons return to a lower energy level.• The line emission spectrum of hydrogen provides evidence for the existence of electrons in discrete energy levels, which converge at higher energies. • The main energy level or shell is given an integer number, n, and can hold a maximum number of electrons, 2n2.• A more detailed model of the atom describes the division of the main energy level into s, p, d and f sub-levels of successively higher energies.• Sub-levels contain a fixed number of orbitals, regions of space where there is a high probability of finding an electron.• Each orbital has a defined energy state for a given electronic configuration and chemical environment and can hold two electrons of opposite spin.• Description of the relationship between colour, wavelength, frequency and energy across the electromagnetic spectrum. • Distinction between a continuous spectrum and a line spectrum.• Description of the emission spectrum of the hydrogen atom, including the relationships between the lines and energy transitions to the first, second and third energy levels. • Recognition of the shape of an s atomic orbital and the px, py and pz atomic orbitals.• Application of the Aufbau principle, Hund’s rule and the Pauli exclusion principle to write electron configurations for atoms and ions up to Z = 36.

OBJECTIVES

Electromagnetic Spectrum

short long

Wavelength (λ)In the data booklet:

E=hν and c=νλ

(where ν is the frequency, and h and c are constants)

(ν is the small greek letter N, not v for velocity, pronounced nu, λ

is pronounced lambda)

• It follows then that the shorter the wavelength, the higher the frequency of the wave, and the more energy it contains.

Emission spectra

When electrons are excited to a higher energy level, and then return to a lower energy level, they release a photon of a specific energy, as shown by a specific frequency of light.

SpectroscopyEmission (line) spectrum

Absorption spectrum

Continuous spectrum

Spectral line series

Emission Spectrum of Hydrogen - convergence

Electrons moving back to the lowest energy states and over the longest distances release the highest E (short λ).

In each series the lines converge meaning higher levels/shells get closer together.

Actual values

Convergence

Outer shells become closer together, so spectral lines get closer together –called convergence.

Spectral fingerprint of the elements

Spectrophotometer with discharge lamps

Periodic table of element emission and absorption spectra

Electron Shells

Although simplistic, a useful way to look at shells is to use the periods in the Periodic table.

Electron Configuration

• Electrons go in shells or energy levels. The energy levels are called principle energy levels, 1 to 4. The maximum number of electrons an energy level (n) can hold is 2n2.

• The energy levels contain sub-levels.Principle energy

level

Maximum number of electrons

Number of sub-levels

1 2 1

2 8 2

3 18 3

4 32 4

These sub-

levels are

assigned the

letters,

s, p, d, f

Sublevels

• Each type of sub-level can hold a differentmaximum number of electron.

Sub-level

Maximum number of electrons

s 2

p 6

d 10

f 14

• The energy of the sub-levels increases from s to p to d to f. The electrons fill up the lower energy sub-levels first.

Electron Configuration

• Let’s take a look at the Periodic Table to see how this fits in.

Electron Configuration

5f

4f

6d

5d

4d

3d

6p

5p

4p

3p

2p

7s

6s

5s

4s

2s

1s

7p

3s

This diagram helps you to work out the order in which orbitals fill: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, ….. However, it can be easier to read across the periodic table, but remember that the first transition metal row is 3d:

1s 1s

2s 2p

3s 3p

4s 3d 4p

(One of these needs to be memorized)

• So how do you write it?

1s2

Energy levelSub-level

Number of

electrons

Example

For magnesium:

1s2, 2s2, 2p6, 3s2

Electron Configuration

• The electronic configuration follows a pattern – the orderof filling the sub-levels is 1s, 2s, 2p, 3s, 3p…

• After this there is a break in the pattern, as the 4s fillsbefore 3d (The electrons fill up the lower energy sub-levels first)

• Taking a look at the table below can you work out why thisis?

• This is because the 4s

sub-level is of

lower energy than the

3d sub-level.

Electron Configuration

Electrons and Sub-Levels

4s of "lower" energy than 3d

Distance from nucleus

Energy

1s

2s

2p

3s

3p

3d 4s

4p

4d

4f

Ionisation energy

Writing Electronic Configurations

• The order in which the energy levels are filled is called the Aufbau Principle.

• Example (Sodium: 2, 8, 1)

• There are two exceptions to the Aufbau principle.

• The electronic configurations of chromium andcopper do not follow the pattern – they areanomalies!

• Chromium – 1s2, 2s2, 2p6, 3s2, 3p6, 3d5, 4s1

• Copper – 1s2, 2s2, 2p6, 3s2. 3p6, 3d10, 4s1

Writing Electronic Configurations

• When an atom loses or gains electrons toform an ion, the electronic configurationchanges:

- Positive ions: formed by the loss of e-

- Negative ions: formed by the gain of e-

1s2 2s2 2p6 3s1 1s2 2s2 2p6

1s2 2s2 2p4 1s2 2s2 2p5

Na atom Na+ ion

O atom O- ion

Writing Electronic Configurations for ions

• With the transition metals it is the 4selectrons that are lost first when they formions:

– Titanium (Ti) - loss of 2 e-

– Chromium (Cr) - loss of 3 e-

1s2 2s2 2p6 3s2 3p6 3d2 4s2 1s2 2s2 2p6 3s2 3p6 3d2

Ti atom Ti2+ ion

Cr atom Cr3+ ion

1s2 2s2 2p6 3s2 3p6 3d5 4s1 1s2 2s2 2p6 3s2 3p6 3d3

Writing Electronic Configurations for transition metals

• Abbreviations can also be used in electronconfiguration for simplicity sake.

– Titanium (Ti):

1s2 2s2 2p6 3s2 3p6 3d2 4s2 Ar] 3d2 4s2 or [Ar] 4s2 3d2

[Ar] always represents 1s2 2s2 2p6 3s2 3p6

Other noble gases in the VIII group of theperiodic table are used as well, such as, [He], [Ar],[Kr], [Xe], etc.

Writing Electronic Configurations – Condensed form

Orbitals• The energy sub levels are made up of orbitals,

each which can hold a maximum of 2 electrons.

• Different sub-levels have different number oforbitals:

Sub-levelNo. of orbitals

Max. no. of electrons

s 1 2

p 3 6

d 5 10

f 7 14

Main energy

levelSub-levels

Max. no. of

electron pairs

in sub-level

Max. no. of

electrons in

sub-level

Max. no. of

electrons in

main level

1 s 1 2 2

2s 1 2

8p 3 6

3

s 1 2

18p 3 6

d 5 10

4

s 1 2

32p 3 6

d 5 10

f 7 14

Energy Levels and Sub-levels

Shap

es o

f th

e o

rbit

als

Orbital diagrams

• Within a sub-level, the electrons occupy orbitalsas unpaired electrons rather than paired electronsand these all spin in the same direction. (This isknown as Hund’s Rule).

• We use boxes to represent orbitals:

1s

2s

2p

Electronic configuration of carbon, 1s2, 2s2, 2p2

Orbital diagrams

• The arrows represent the electrons in the orbitals.

• The direction of arrows indicates the spin of theelectron.

• Paired electrons will have opposite spin, as thisreduces the mutual repulsion between the pairedelectrons (This is known as the Pauli exclusionprinciple)

Electron configuration of carbon: 1s2, 2s2, 2p2 1s

2s

2p

Notice how the arrowsin each box of the 1s2

& 2s2 are oppositewhich means oppositespin

Problem: Using boxes to represent orbitals, give thefull electronic configuration of the following atoms:

a) lithium

b) fluorine

c) potassium

d) nitrogen

e) oxygen

1s

2s

2p

Lithium: 1s2, 2s1

1s

2s

2p

Fluorine: 1s2, 2s2, 2p5

1s

2s

2p

Nitrogen: 1s2, 2s2, 2p3

1s

2s

2p

Potassium: 1s2, 2s2, 2p6, 3s2, 3p6, 4s1

1s

2s

2p

3s

3p

4s

Oxygen: 1s2, 2s2, 2p4

1s

2s

2p

Orbital diagram review• Electrons enter the lowest energy orbital

available (Aufbau principle). • Electrons prefer to occupy orbitals on their own,

spin in the same direction, and only pair up when no empty orbitals of the same energy are available (Hund's Rule).

• Paired electrons have the opposite spin (Pauli exclusion principle).

• In ions, the electrons in the highest energy levels are lost first, but when losing electrons, electrons are lost from 4s before 3d (the energy levels are very close, and when electrons fill them, 4s goes above 3d).