The Atom

Table of Contents
1. Historical development and models of the atom
2. Atomic mass, units and size
3. Rutherford's alpha-particle scattering experiment (detail)
4. Structure of the atom: particles and properties
5. Atomic number and mass number
6. Relative atomic mass and isotopes
7. Electronic structure and electron configuration
8. Importance of electron configuration
9. Informal experiment: flame tests (class demonstration)
10. Exercises and practice questions
11. Group activity: build a model of an atom
12. Summary
13. End-of-chapter exercises
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The atom

Introduction

We have investigated many types of matter and how materials are classified. To understand why different materials behave differently we must examine their basic building blocks - the atom. Atoms form all structures and organisms in the universe: planets, stars, plants, air and living organisms are composed of different combinations of atoms.

Our modern understanding of the atom developed over a long period of time, through the work of many scientists. The following sections reconstruct the historical development, define the parts of an atom, explain atomic masses and sizes, and present electron arrangement and its importance for chemical behaviour.

Historical development and models of the atom

Scientific models of the atom evolved as new evidence and experiments became available. A model is a simplified representation of a real system that helps us understand its properties. Atomic models are representations based on experimental observations; they are not literal pictures but tools for understanding and prediction.

Early ideas

The idea that matter is made of indivisible particles was proposed by Greek philosophers such as Leucippus and Democritus in the fifth century BC. The Greek word atomon means "indivisible". Early ideas were qualitative; later scientists developed experimental evidence and quantitative models.

Dalton's model

John Dalton proposed that all matter is composed of small indivisible particles called atoms. Dalton's model treated atoms as solid, indivisible spheres and helped explain conservation of mass and definite proportions in chemical reactions. At the time electrons and the nucleus were not known.

Thomson's plum-pudding model

After the discovery of the electron by J. J. Thomson (1897), the plum-pudding model (1904) pictured the atom as a diffuse positive "soup" with negatively charged electrons embedded in it, analogous to plums in a pudding. This model accounted for the existence of electrons but did not explain how positive charge was distributed or how electrons were arranged.

Rutherford's nuclear model

Ernest Rutherford (1911) proposed that positive charge and most of the mass are concentrated in a tiny central nucleus, with electrons around it. He compared the atom to a miniature solar system: a small dense nucleus with electrons orbiting at relatively large distances. Rutherford's model followed from scattering experiments (see detailed section on the alpha-particle scattering experiment below).

Bohr's model

Niels Bohr (1913) refined the planetary view by introducing quantised electron orbits. Electrons can occupy only certain discrete energy levels; transitions between levels explain why atoms emit or absorb light at particular wavelengths. Bohr's model explained atomic emission spectra of hydrogen but had limitations for multi-electron atoms.

Discovery of the neutron

Rutherford predicted a neutral particle in the nucleus to account for mass without extra positive charge. James Chadwick discovered the neutron in 1932 and measured its mass. The nucleus is therefore made of protons (positively charged) and neutrons (neutral), collectively called nucleons.

Quantum models and modern view

Further development led to quantum mechanics. Scientists such as Erwin Schrödinger, Werner Heisenberg, Max Born and others developed the quantum-mechanical model where electrons occupy orbitals - regions of space with specific energies and probability distributions. This modern view replaces classical orbits with probability clouds and explains atomic structure and chemical behaviour more accurately.

Atomic mass, units and size

Because atomic masses are extremely small, special units are used and comparisons are more convenient than absolute SI masses.

Atomic mass unit (amu or u)

The atomic mass unit (u) is defined using the carbon standard: one atom of carbon-12 is assigned a mass of exactly 12.00 u. In SI units, 1 u ≈ 1.6605 × 10^-27 kg (≈ 1.6605 × 10^-24 g). Using atomic mass units simplifies comparison of masses of atoms.

Example atomic masses (approximate, in u):

How heavy and how large is an atom?

Individual atoms have masses around 10^-27 to 10^-26 kg. For example, a hydrogen atom has mass ≈ 1.67 × 10^-27 kg and a carbon atom ≈ 1.99 × 10^-26 kg. Atoms are extremely small; the nucleus occupies only a tiny fraction of the atom's volume. A common analogy: if an atom were the size of a football stadium, the nucleus would be roughly the size of a pea at the centre.

Rutherford's alpha-particle scattering experiment (detail)

Rutherford investigated how alpha (α) particles (positively charged helium nuclei) scatter when directed at thin metal foils (gold foil experiment). He placed a zinc sulphide screen around the foil to detect where alpha particles struck. Predictions from the plum-pudding model would have most α particles deflected only slightly. The experiment showed that most α particles passed through the foil, a small fraction were deflected at large angles and a very few rebounded. Rutherford concluded that:

• Most of the atom is empty space.
• Positive charge and most of the mass are concentrated in a very small central nucleus.
• Electrons occupy much larger regions around the nucleus.

Structure of the atom: particles and properties

An atom consists of a central nucleus (protons and neutrons) and electrons around the nucleus. The protons and neutrons are called nucleons.

Electrons are treated as elementary (point-like) particles in atomic models; they cannot be subdivided by current experimental evidence.

Atomic number and mass number

Atomic number (Z) is the number of protons in the nucleus and determines the chemical identity of the element. The atomic number appears on the periodic table and is an integer (1 to about 118).

Mass number (A) is the total number of nucleons (protons + neutrons) in the nucleus. The number of neutrons is N = A - Z.

Standard notation for nuclides is written with mass number and atomic number as superscript and subscript, respectively, before the chemical symbol:

^A_ZX where X is the chemical symbol, A is the mass number and Z is the atomic number. For example, the iron nucleus with 26 protons and 30 neutrons is written as ⁵⁶₂₆Fe (A = 56, Z = 26, N = 30).

For a neutral atom the number of electrons equals the number of protons. If electrons are gained or lost, the resulting charged atom is called an ion. For example, Na⁺ is a sodium ion that has lost one electron; Cl^- is a chlorine ion that has gained one electron.

Example 1: Standard notation

Relative atomic mass and isotopes

Relative atomic mass of an element is the average mass of all naturally occurring isotopes of that element, weighted by their natural abundances. Units are atomic mass units (u). The value shown on the periodic table is the relative atomic mass (often a decimal) and not necessarily an integer.

Isotopes

Isotopes are nuclides of the same element that have the same number of protons (same Z) but different numbers of neutrons (different N), and therefore different mass numbers A. Chemical properties of isotopes are essentially the same because chemistry depends mainly on the number of electrons/protons, but nuclear stability and some physical properties can differ.

Example: Chlorine has two common isotopes: Cl-35 (35 is the mass number: 17 protons + 18 neutrons) and Cl-37 (17 protons + 20 neutrons). If Cl-35 occurs with 75% natural abundance and Cl-37 with 25%, the average relative atomic mass is calculated by weighting:

Electronic structure and electron configuration

Electrons in an atom have quantised energies. Electrons with lowest energy are closest to the nucleus; higher-energy electrons are found further away. Electrons occupy energy levels (also called shells) labelled 1, 2, 3, ...

Energy levels and capacity

The first main energy level (n = 1) can hold up to 2 electrons. The second (n = 2) can hold up to 8 electrons (2 in an s subshell and 6 in p subshells). The third level starts to fill similarly (it contains s, p and d subshells).

Orbitals

An atomic orbital is a region of space around a nucleus where an electron is likely to be found. Subshells and orbitals have maximum capacities: an s orbital holds 2 electrons, a p subshell (three p orbitals) holds 6 electrons, etc.

Rules for filling orbitals

• Each orbital can hold a maximum of two electrons (with opposite spins).
• Electrons fill the lowest-energy orbitals available first (principle of minimum energy).
• Hund's rule: within a subshell, electrons occupy separate orbitals singly before pairing up.
• Pauli exclusion principle: no two electrons in an atom can have the same set of quantum numbers; in practice electrons in the same orbital have opposite spins.

Aufbau principle and notation

The Aufbau principle describes the order in which orbitals are filled as electron number increases (building up). Electron configurations are often written in spectroscopic notation, e.g. lithium: 1s² 2s¹. Noble-gas shorthand (condensed notation) uses the nearest preceding noble gas in square brackets followed by the outer electrons, e.g. magnesium: [Ne] 3s².

Examples of Aufbau diagrams and spectroscopic notation

Valence electrons are the electrons in the outermost energy level and determine most chemical behaviour. Core electrons are all the electrons that are not valence electrons.

Importance of electron configuration

Electron configuration explains periodic trends, bonding and reactivity. Atoms are most stable when their outer energy level is full; noble gases have full outer levels and are chemically inert. Other atoms tend to gain, lose or share electrons to achieve a more stable (often noble-gas-like) configuration. The octet rule (mainly for main-group elements) summarises this tendency: atoms often aim for eight electrons in their valence shell.

Informal experiment: flame tests (class demonstration)

Aim: To observe flame colours produced by different metal cations and relate these to electronic transitions.

Apparatus:

• Watch glass
• Bunsen burner
• Methanol (or ethanol as a volatile solvent)
• Wooden toothpicks or skewers
• Samples of metal salts (for example: NaCl, CuCl₂, CaCl₂, KCl)
• Metal powders (optional: copper, magnesium, zinc, iron)

Safety: Tie back long hair, secure loose clothing, work in a ventilated space, and keep flammable materials away from the flame.

Method:

1. Dip a clean toothpick in methanol.
2. Dip the toothpick into a small amount of the metal salt or powder.
3. Wave the toothpick quickly through the Bunsen burner flame (do not hold it stationary in the hottest part of the flame).
4. Observe and record the flame colour for each metal.

Conclusion: Different metal cations produce characteristic flame colours because electrons in the metal ions absorb energy, move to higher energy levels, then fall back emitting light at characteristic wavelengths (line emission).

Exercises and practice questions

Match exercise

Match the information in column A with the key discoverer in column B.

Further exercises (cleaned and structured):

1. Explain the meaning of each of the following terms:

   1. nucleus
   2. electron
   3. atomic mass

2. Complete the following table (use atomic data from the periodic table):

   Element	Atomic mass units (A)	Atomic number (Z)	Number of protons	Number of electrons	Number of neutrons
   Mg	24	12
   O		8
   Cl	17
   Ni	28	40
   Zn		30
   C	12	6
   Al3+	13
   O2-	18

3. Use standard notation to represent the following elements:

   1. potassium
   2. copper
   3. chlorine

4. For the element ³⁵₁₇Cl, give the number of:

   1. protons
   2. neutrons
   3. electrons

5. Which of the following atoms has 7 electrons?

   1. ⁵²He
   2. ¹³C
   3. ⁷³Li
   4. ¹⁵N

6. Complete isotope and nuclide tables and solve percentage-abundance relative mass problems (examples were given earlier).

Further exercises on isotopes and electron configurations (selected):

1. State whether pairs of atoms are isotopes and give numbers of protons, neutrons and electrons.
2. Calculate relative atomic masses given isotopic abundances (as in the chlorine and magnesium examples).
3. Draw Aufbau diagrams and give spectroscopic electron configurations for elements and ions (examples: Mg, K, S, Ne, N, Ca²⁺, Cl^-).
4. Rank a set of elements by reactivity and justify using electron configuration and valence electrons.

Group activity: build a model of an atom

In small groups, build a physical model of an atom using craft materials. Before building, discuss:

• What parts make up the atom?
• Which materials best represent the nucleus and electrons?
• How will you show relative positions and motion (qualitatively)?

After building, evaluate accuracy and limitations of the model and suggest improvements. Use models to discuss why they are simplified and what they do and do not show about real atoms.

Summary

• Many scientists contributed to the development of atomic theory (examples include Dalton, J. J. Thomson, Rutherford, Bohr, Chadwick, Schrödinger and Heisenberg).
• Because atomic masses are tiny, atomic mass units (u) are used. 1 u ≈ 1.6605 × 10^-27 kg.
• Relative atomic mass is the weighted average mass of naturally occurring isotopes and is the value given on periodic tables.
• An atom consists of a small dense nucleus (protons and neutrons) surrounded by electrons; most of the atom's volume is empty space.
• Atomic number Z is the number of protons; mass number A is the number of protons plus neutrons. Standard nuclide notation is ^A_ZX.
• Isotopes have the same Z but different N and A. The chemical properties of isotopes are largely similar.
• Electrons occupy orbitals with quantised energies. Electron configuration determines valence electrons and chemical behaviour.
• Aufbau diagrams, Hund's rule and the Pauli exclusion principle are tools to determine electron configurations.
• Flame tests and emission spectra are direct consequences of electronic transitions between energy levels.

End-of-chapter exercises

1. Write down only the word/term for each of the following descriptions.

   1. The sum of the number of protons and neutrons in an atom
   2. The defined space around an atom's nucleus, where an electron is most likely to be found

2. For each of the following, say whether the statement is true or false. If it is false, re-write the statement correctly.

   1. ²⁰₁₀Ne and ²²₁₀Ne each have 10 protons, 12 electrons and 12 neutrons.
   2. The atomic mass of any atom of a particular element is always the same.
   3. It is safer to use helium gas rather than hydrogen gas in balloons.
   4. Group 1 elements readily form negative ions.

3. The three basic components of an atom are:

   1. protons, neutrons, and ions
   2. protons, neutrons, and electrons
   3. protons, neutrinos, and ions
   4. protium, deuterium, and tritium

4. The charge of an atom is:

   1. positive
   2. negative
   3. zero
   4. depends on the number of neutrons

Use the explanations, examples and rules in this chapter to answer the exercises. Where numerical values are needed, use the appropriate data from the periodic table or the sections above.