Properties and Structure of MatterPart 2

Atomic structure

Atomic Structure and Atomic Mass

● investigate the basic structure of stable and unstable isotopes by examining:

– their position in the periodic table

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– the distribution of electrons, protons and neutrons in the atom– representation of the symbol, atomic number and mass number (nucleon number)

● model the atom’s discrete energy levels, including electronic configuration and spdf notation (ACSCH017, ACSCH018, ACSCH020, ACSCH022)

● calculate the relative atomic mass from isotopic composition (ACSCH024) ● investigate energy levels in atoms and ions through:

– collecting primary data from a flame test using different ionic solutions of metals (ACSCH019)

– examining spectral evidence for the Bohr model and introducing to the Schrödinger model

● investigate the properties of unstable isotopes using natural and human-made radioisotopes as examples, including but not limited to:– types of radiation– types of balanced nuclear reactions

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Properties and Uses of Radioisotopes1. The Basic Structure of Stable and Unstable Isotopes.

What are isotopes?Radioisotopes are widely used in medicine, industry and scientific research, and new applications for their use are constantly being developed.

Background:

Atoms contain protons and neutrons in a nucleus surrounded by electrons in energy

level shells. Isotopes of an element are atoms of that element containing the same

number of protons but different numbers of neutrons.

If the nucleus of an atom contains excess energy the nucleus is unstable and can emit

radiation. The radiation emitted is characteristic of the nucleus. The emitted radiation

can be used in many ways in industry and medicine.

https://phet.colorado.edu/en/simulation/isotopes-and-atomic-mass

Position on the Periodic Table.All elements with atomic number greater than 83 (Bismuth) are unstable

(radioactive).

Only 279 of about 2000 known isotopes are stable. In a stable isotope nucleus, the

protons and neutrons are in a low energy level and are unable to emit radioactivity.

The emission of radiation is unaffected by factors such as pressure,

temperature and the presence of catalyst that influence the reaction rate.

Radioactivity is the spontaneous disintegration of an unstable isotope leading

to the emission of radiation.

Radioactive isotopes are unstable. They emit radiation as they spontaneously

release energy. This is called radioactive decay.

An unstable isotope can be called a radioisotope, an abbreviation of the term

radioactive isotope.

The time for the radioactivity level from a given amount of radioactive isotope to

be halved is called its half-life. Each radioactive isotope has a characteristic half-

life.

Radioactive isotopes occur when the atom has a large atomic mass and a high

neutron to proton ratio.

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Chemistry in FocusCheck your understanding 3.2, pages 49 – 50 Check your understanding 3.3, page 54

Radioactive isotopes can emit three types of radiation:

Properties of three radioactive emissions.

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Alpha EmissionAlpha decay involves the emission of a particle (4

2He) from the unstable nucleus.

When an atom loses a particle, the atomic mass decreases by 4 and the atomic

number decreases by 2. For example the radioactive decay of uranium-238 into

thorium-234 by a emission:

Radioactive decay

23892U → 234

90Th + 42He

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Beta Emission Beta (β) decay involves the conversion of a neutron into a proton and a beta

particle (an electron).

10n → 1

1p+ + 0-1e

Note: A hydrogen ion, H+ is a proton. Hence the above equation can be written as:

10n → 1

1H+ + 0-1e

The β decay of carbon-14 is shown in the following nuclear equation:

146C → 14

7N + 0-1e

Carbon-14 is used in carbon dating.

Chemistry in FocusCheck your understanding 3.3, page 54

Research:

1. What is a Transuranic Element

2. How are they made : Nuclear Reactor and Cyclotron

3. One use of a radioisotope in Medicine and one in Industry

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● model the atom’s discrete energy levels, including electronic configuration and spdf notation1. Electron configuration, orbital and energy levels.

Electrons are found in orbitals (or shells) around the nucleus of an atom.

Each orbital represents an energy level.

An electron in an orbit has a specific energy level.

Principal Energy Levels or Shells

Principal energy levels or

Shells

Number of subshells Name of subshells

K = 1 1 s

L = 2 2 s

p

M = 3 3 s

P

d

N = 4 4 s

p

d

f

Electron Configuration.

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The electron configuration of an element is the number of electrons in each shell of

the atom according to energy levels.

Electrons fill the orbital: 2,8,18,32.

For the first twenty elements, the maximum number of electrons in each shell is

2,8,8,2.

Therefore, Hydrogen, atomic number 1, has the electron configuration of 1.

Sodium, atomic number 11, has the electron configuration: 2,8,1

Homework: Complete the table below filling in all the electron configurations.

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Element Electron configuration

Hydrogen 1

Helium

Lithium

Beryllium

Boron

Carbon

Nitrogen

Oxygen

Fluorine

Neon

Sodium 2,8,1

Magnesium

Aluminium

Silicon

Phosphorus

Sulfur

Chlorine

Argon

Potassium

Calcium

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Orbitals and Sublevels

An orbital is a volume of space surrounding the nucleus of an atom through which

one or two electrons may randomly move.

Filling Subshells

Hund's rule: every orbital in a subshell is singly occupied with one electron before

any one orbital is doubly occupied, and all electrons in singly occupied orbitals have

the same spin.

The order in which each shell is filled is as follows:

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Homework: Complete the following table, giving the electron configurations of the first

twenty elements in terms of their s,p,d subshells.

Element Electron configuration

Hydrogen 1s1

Helium

Lithium

Beryllium

Boron

Carbon 1s2, 2s2, 2p2

Nitrogen 1s2, 2s2, 2p3

Oxygen

Fluorine

Neon 1s2, 2s2, 2p6

Sodium 1s2,2s2,2p6,3s1

Magnesium

Aluminium

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Silicon

Phosphorus

Sulfur

Chlorine

Argon

Potassium

Calcium

● investigate energy levels in atoms and ions through:– collecting primary data from a flame test using different ionic solutions of metals

(ACSCH019) – examining spectral evidence for the Bohr model and introducing to the

Schrödinger model

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The Bohr model (in 1913), depicts the atom as a small, positively charged nucleus

surrounded by electrons that travel in circular orbits around the nucleus.

Each orbit represents an energy level. When an electron is in its normal energy level

it is said to be in its ground state.

Electrons can absorb energy and move to a higher energy orbital. Electrons that do

this are said to be in an excited state.

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Bohr incorporated Planck’s quantum theory in that electrons could only have certain

discrete energies. Each atom had a set of distinct energy levels. The Bohr could not

be quantitatively used to interpret more complex spectra.

Schrodinger Equation

Schrodinger, in 1932, incorporated that electrons had wave properties as well as

particle properties. He used this to develop the Schrodinger Equation.

He used an expression for the wavelength of an electron in terms of its energy and

applied the mathematics of waves to develop an equation to calculate the probability

of finding an electron at any particular location around the nucleus.

This is called the Schrodinger Equation.

The Schrodinger Equation was successfully used to interpret the emission spectrum

of atoms with many electrons. (Chemistry in Focus)

Every line on an emission spectrum corresponds to an electron falling from an

excited state to its ground state.

In their normal state atoms do not emit light, but if an atom is given extra energy by

either heat or electricity, the electrons within the atom are excited into a higher

energy level.

When the electrons fall back down to their ground state (lowest energy level, the

atoms are normally in this state) the atom will emit light (energy).

The wavelength and frequency of light emitted is related to the energy released, with

the greater the energy, the shorter the wavelength and the greater the frequency.

If a sample of an element in a gas discharge tube and if the light emitted as the atoms

fall back from excited states to ground states is examined through a spectroscope we

observe a series of bright coloured lines on a black background.

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This is called an emission spectrum of the element (atomic emission spectrum).

Different elements have different spectra because the transition between electron shells

involve different energy changes and so light emission are at slightly different wavelength.

Flame Test Practical

Chapter Review Questions P73

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