Atomic Physics

Welcome to the attractive world of atomic structure! In this topic, we’ll explore the various atomic models that have been proposed to explain the behaviour of atoms. You’ll also learn how to calculate the energy of a photon produced during an electron’s transition between energy levels.

You will learn how to compute simple nuclear equations. Nuclear reactions involve changes to the nucleus of an atom, resulting in a new element or isotope. These reactions can be either natural or artificial and have numerous applications in fields like medicine, energy, and industry. Balancing nuclear reactions is crucial to understand the changes that occur during these reactions. It helps us to identify the reactants and products, to determine the number of atoms involved, and to understand the conservation of mass and energy.

By the end of this section, you will be able to:

• Explain various atomic models and their limitations.

• Calculate the energy of a photon during a transition.

• Describe the structure of the nucleus of the atom.

• Explain radioactivity.

• Balance basic nuclear reactions.

KEY IDEAS

• Dalton’s Atomic Model (1803): Atoms are indivisible, indestructible particles.

• Thompson’s Atomic Model (1897): Atoms have a positive nucleus surrounded by electrons.

• Rutherford’s Atomic Model (1911): Atoms have a small, dense nucleus with electrons orbiting around it. Bohr’s Atomic Model (1913): Electrons occupy specific energy levels around the nucleus.

• One piece of evidence of Bohr’s atomic model is the emission and absorption spectra of different elements, produced when electrons excite and de-excite, either absorbing or emitting a photon.

• The nucleus of the atom can undergo changes that make it more stable, which involves the emission of an alpha particle, beta particle, or gamma photon. In this section we will study this effect, known as nuclear radioactivity, in more detail.

Atomic Physics

Atomic physics is the field of physics that studies the structure of an atom, atomic models, and the constituent particles in an atom.

Atoms An atom is the smallest particle of an element that retains the properties of that element. All pure (uncombined) elements are composed of atoms. For example, iron is composed of iron (Fe) atoms, helium gas is composed of helium (He) atoms, etc.

Fig. 8.1: The structure of the atom The figure above gives a simplified artist’s impression of our modern understanding of the structure of the atom, although it is much more complicated!

We have not always believed that the atom is built this way. The model of the atom has developed significantly since Dalton’s model in 1803. Read the dialogue below to gain insight as to the contribution of three very important scientists.

Activity 8.1: A debate on atomic models Participants:

• Emmanuella (representing J.J. Thompson)

• Love (representing Ernest Rutherford)

• Nana Poku (representing Niels Bohr) Moderator: Welcome everyone to today’s debate about atomic models! Each participant will explain their ideas and respond to each other. Let’s start with Emmanuella.

Fig. 8.2: High School Debate Competition

Round 1: Opening Statements

Emmanuella (Thompson):

“Thank you! I want to present J.J. Thompson’s atomic model, known as the ‘plum pudding model.’ This model suggests that atoms are made up of a positively charged ‘soup’ with negatively charged electrons mixed throughout.

This idea helps explain how atoms can be neutral and shows that they contain smaller particles.”

Love (Rutherford):

“Thank you, Emmanuella. While I appreciate your contributions, I believe the ‘plum pudding model’ lacks a clear structure. My gold foil experiment demonstrated that atoms have a small, dense centre called the nucleus, which contains most of the atom’s mass. Electrons orbit around this nucleus, showing that atoms are mostly empty space. This was a significant advancement in understanding atomic structure.”

Nana Poku (Bohr):

“Thank you, Love. While I respect Rutherford’s findings, his model doesn’t explain why electrons stay in their orbits without spiralling into the nucleus.

My Bohr model introduces energy levels, meaning electrons can only occupy certain paths or orbits. This explains why atoms are stable and how they emit light when they gain energy.”

Round 2: Responses and Critiques

Emmanuella (Thompson):

“Love, your model shows a nucleus, which is a great step forward. However, it doesn’t explain how electrons can stay in orbit without losing energy. My model introduces subatomic particles, which is essential for understanding atomic structure.”

Love (Rutherford):

“Emmanuella, your ‘soup ‘ analogy is interesting, but it doesn’t account for the concentrated mass in the nucleus. Nana Poku, your fixed orbits are groundbreaking, but they can be misleading since electrons don’t follow set paths. They exist in areas of probability around the nucleus.”

Nana Poku (Bohr):

“Love, I appreciate your emphasis on the nucleus. My model explains electron behaviour more accurately. The fixed orbits represent areas where we are likely to find electrons, addressing some limitations of earlier models.”

Round 3: Closing Statements

Moderator: Thank you all for your insightful arguments! This debate highlights the evolution of atomic models and how each contribution has shaped our understanding of atoms.

Modern atomic structure Discovery An atom consists of a positive nucleus, which is found at its centre and surrounded by negative electrons that orbit (move around) the nucleus. The nucleus also contains protons and neutrons.

Most great scientists have contributed in various ways that led to the discovery of the modern atomic structure. The contributions of some of these scientists are highlighted below.

Fig. 8.3: Evolution of Atomic Models by J.J. Thomson, Ernest Rutherford, and Niels Bohr J. J. Thompson’s description of the atom is that an atom is a solid sphere of positive charge with the negative electrons stacked uniformly in it. The positive charges occupy the greatest volume, while the electrons occupy the smallest volume, and the whole atom becomes neutral. I.e. the electron charge balanced the positive charge.

Fig. 8.4: J. J. Thompson’s description of the atom Limitations of Thompson’s Model of the Atom An experiment designed some years later by Ernest Rutherford heralded results that disagreed with Thompson’s model. Complete Activity 3 to learn more about this experiment, and be sure to review the solutions in Annex 1.

Activity 8.2: Constructing a 3D Plum Pudding Atomic Model Create a 3D model of the plum pudding, labelling its key features.

Rutherford’s Alpha Scattering Experiment

In 1909, Ernest Rutherford came up with a new model of the atom based on observations he made in his alpha scattering experiment, which were:

1. Most of the alpha particles went through the thin metal foil without any change in path.

2. A few alpha particles were deflected through small angles.

3. Very few alpha particles were deflected backwards.

Activity 8.3: Simulating Rutherford’s Gold Foil Experiment

Use the simulation linked above to replicate Rutherford’s experiment, and explain how he came to the following conclusions:

1. The atom is mostly made up of empty space.

2. The atom holds at its centre a small, positively charged nucleus, which contains the majority of the mass of the atom.

3. Electrons orbit the atom very distantly.

You should also explain how Rutherford’s observations disagreed with Thompson’s Plum Pudding model.

Ernest Rutherford’s Model of the Atom (The

Nuclear Model)

Rutherford proposed that the atom has a positively charged nucleus, which contains most of the mass of the atom. The nucleus is surrounded by electrons in orbit, just as the planets orbit the sun in our solar system. He said that the nucleus, together with the electrons, constitutes an electrically neutral unit.

Fig. 8.5: Ernest Rutherford’s Model of the Atom

Rutherford proposed the following:

1. Nucleus: The atom has a centre called the nucleus, which contains positively charged particles called protons. Most of the atom’s mass is in the nucleus.

2. Electrons: Around the nucleus there are small, negatively charged particles called electrons. These electrons move in orbits at a relatively large distance from the nucleus.

3. Orbits and Forces: Electrons move in circular paths around the nucleus due to an attractive force between the positively charged protons and the negatively charged electrons.

4. Empty Space: Most of the atom is empty space, which is why many particles can pass through materials like gold foil without being deflected.

5. Deflection: When an alpha particle (a type of positive particle) gets close to the nucleus, it can experience a strong repulsive force, causing it to change direction or bounce back if it’s heading straight toward the nucleus.

This model helped us understand that atoms are mostly empty space with a dense centre and electrons orbiting around it.

Limitations of Rutherford’s Model of the Atom

1. Electrons Should Spiral In: Rutherford’s model shows electrons moving in orbits around the nucleus. However, when charged particles like electrons move, they should lose energy and spiral into the nucleus, which would destroy the atom. But we don’t observe atoms being destroyed this way; they stay stable.

2. Radiation Emission: If electrons were continuously emitting radiation as they moved, we would expect them to produce a mix of different frequencies of light. Instead, studies show that atoms only emit light at specific frequencies, meaning they have fixed energy levels.

3. Surprising Deflections: Rutherford’s experiments with alpha particles (which are much heavier than electrons) showed that these particles were sometimes deflected at very large angles when passing through gold atoms.

This was unexpected because, based on earlier models, we would expect only small deflections. Rutherford compared it to firing a cannonball at a piece of tissue paper and having it bounce back!’

Activity 8.4 Constructing a 3D Model of Rutherford’s Atomic Structure Create a 3D model of Rutherford’s atom, labelling its key features.

Neils Bohr’s Atomic Model

Niels Bohr proposed a new model of the atom in 1913 that built on Rutherford’s ideas but addressed some of the limitations highlighted above. Here are the key points of Bohr’s model:

1. Electrons in Orbits: Electrons move in specific circular paths or “orbits” around the nucleus. Each orbit corresponds to a specific energy level.

2. Quantized Energy Levels: Electrons can only occupy certain energy levels and cannot exist between these levels. When an electron jumps from a higher energy level to a lower one, it emits energy in the form of light.

3. Stable Orbits: Electrons in these orbits do not lose energy and spiral into the nucleus. They remain stable due to the balance between the attractive force of the nucleus and the motion of the electrons.

4. Emission of Light: When electrons move between energy levels, they absorb or release energy, resulting in the emission or absorption of light.

This explains how atoms emit specific colours of light.

5. Hydrogen Atom: Bohr’s model was particularly successful in explaining the hydrogen atom’s spectrum, where the light emitted corresponds to specific energy transitions of electrons.

Protons and the neutrons are found in nucleus and the electrons move round the nucleus in selected orbits.

Fig. 8.6: Neils Bohr’s Atomic Model

Let’s observe the fluorescent light bulb… Fig. 8.7: Fluorescent light bulb Did you know that a fluorescent bulb works by excitation and subsequent de- excitation of mercury atoms to produce light? Read on for more information!

Energy Levels

An atom’s energy levels are the different energy states that electrons can have within an atom; any other magnitudes of energy are forbidden! The energy levels are often represented as a series of parallel lines, a little like a ladder, with the lowest energies at the bottom and the highest at the top. An example of this is shown below.

Fig. 8.7: Energy Levels

“Ground state” The ground state is defined as the lowest allowed energy state of an atom, molecule, or ion. In other words, the ground state represents the most stable configuration.

Excited states A particle in an excited state is an atom, ion, or molecule in which an electron is at a higher energy level than its ground state.

Transitions of electrons Excitation Electrons can move from one energy level to another. When an electron absorbs energy, it can jump from a lower energy level (orbit) to a higher energy level. This process is called excitation. The energy can come from various sources, such as photons of electromagnetic radiation, particle collisions or electrical energy.

Fig. 8.8: Electron excitation due to energy absorption De-excitation After an electron has been excited to a higher energy level, it is unstable and eventually returns to a lower energy level. This process is called de-excitation. As it loses energy, it releases it in the form of a photon. This may be a visible light photon or may be a photon belonging to another region of the electromagnetic spectrum.

Fig. 8.9: Electron De-excitation due to energy emission

Activity 8.5: Exploring the Electromagnetic Spectrum

List the seven types of electromagnetic wave and research their corresponding frequencies and wavelength.

Table 8.1: Table to complete activity 8.5 Wave Frequency range Wavelength range Calculating the frequency of photons absorbed or emitted during electron energy transitions The figures below represent the movement of an electron from a low to a higher energy level (left) and vice versa (right).

Fig. 8.10a: Electron from a low to a higher energy level Fig. 8.10b: Electron from a high to a lower energy level The energy carried by a photon is calculated using the formula:

E = hf = hc___ λ Where:

E is the energy of the photon (J) h is Planck’s constant (Js⁻¹) f is the frequency of the photon (Hz) c is the speed of light (ms⁻¹) λ is the wavelength of the photon (m) By using this equation, we can calculate the frequency or wavelength of a photon absorbed or emitted by an electron when it transitions from one energy to another.

First, we need to find the change in energy of the electron by reading values form the diagrams and converting electron volts (eV) to Joules (J).

1eV = 1.6 × 10⁻¹⁹J

Activity 8.6: Energy Transitions

From the diagram show that when an electron jumps from orbit n = 3 to orbit n = 2 it produces a photon of red light with an energy of 1.89 eV and a wavelength of 656m.

Activity 8.7: Calculation involving the use of the formula E = hf

1. What is the energy of a photon of frequency 3.2 × 10¹⁰Hz ?

( h = 6 . 63 × 10⁻³⁴J s )

2. What is the wavelength of a photon of energy 6 . 5 × 10⁻¹⁹J ?

3. Find the energy of a photon of wavelength 550 nm . h = 6.63 × 10⁻³⁴J s, c = 3.0 × 10⁸m s⁻¹4. Calculate the frequency of radiation emitted when electrons in hydrogen atom lose 10.21eV.

(h= 6.63×10⁻³⁴Js, leV = 1.6 x 10⁻¹⁹J )

Activity 8.8: Challenge questions

1. What is meant by the term ‘ionisation energy’?

2. What is the frequency of radiation emitted when electrons in hydrogen atom moves from n=2 to n =1 states. Given the information:

(Eₙ = - 13.6/n²eV, h= 6.5 × 10⁻³⁴Js) leV = 1.6 x10⁻¹⁹J; c = 3.0 x10⁸ms⁻¹3. The energy levels of hydrogen atom are given by the expression Eₙ = - 2.16 × 10⁻¹⁸_________ n²J What is the wavelength of radiation, which arises from transitions between n = 3 and n = 2 levels? (h = 6.5 × 10⁻³⁴Js; leV = 1.6 × 10⁻¹⁹J;

c = 3.0 × 10⁸ms⁻¹)

Review Questions 8.1 (Atomic Physics)

1. What are the main limitations of Dalton’s Atomic Model?

2. Describe the key features of Thompson’s Atomic Model.

3. What is the major limitation of Rutherford’s Atomic Model?

4. How does Bohr’s Atomic Model explain electron energy levels?

5. What formula relates the energy of a photon during a transition?

6. What is the significance of Planck’s Constant in calculating photon energy?

Review Questions 8.2 (Nuclear Physics)

1. Describe the structure of alpha, beta, and gamma radiation.

2. Radiation is detected from inside a thick concrete bunker. What type of radiation is it likely to be?

3. Describe the structural difference between carbon-12 and carbon-14.

4. Research and summarise the use of carbon-14 in the aging of fossils.

Review Questions 8.3 (Nuclear decay equations)

1. Some nuclides were obtained in some nuclear bombardment processes.

Determine the values of x, y and z in the following nuclear reaction equation below:

i. ₈₆ ²²²Rn → ₈₄ ˣPo + ₂ ⁴H e + energy

ii. ₈₂ ʸPb → ₈₃ ²¹⁴Bi +− 1 0e + energy

iii. ₄ ⁹Be +₂ ⁴He → ₆ ᶻC + ₀ ¹n + energy

iv. ₉₂ ²³⁵U + ₀ ¹n→ ₅₇ ¹⁴⁸La + ₃₅ ⁸⁵Br + x ₀ ¹n + energy