Problem #5: Describe the Atom [Solved]

Daniel Iyamuremye -

Describe the atom, giving a brief historical background development that led to the present description.

Answer:

As said in the previous problem, an atom is the smallest part of an element that can exist chemically. But the atom is not the ultimate particle of matter; it is itself made of smaller particles called sub-atomic particles and it has a structure.

The atom has been an object of discussions and studies since antiquity. Greek Philosophers asked the question: What are the ultimate constituents of matter? Among the ideas proposed at that time is Democritus’s suggestion that matter discontinuous, i.e. matter is composed of atoms (in Greek, atomos means indivisible).

But it is at the beginning of 19th Century that real scientific research about the nature and structure of the atom started. From that time up to now, many atomic models have been proposed:

I. Dalton’s Atomic Model

In 1803, the British chemist and physicist advanced a first scientific proposition of an atomic model or atomic theory. In his theory:

  1. Elements consist of indivisible small particles called “atoms”
  2. All atoms of the element are identical. Different elements have different types of atoms
  3. An atom can neither be created nor destroyed
  4. Compounds are formed when atoms of different elements join in simple ratios to form molecules
Source: Analist Chemistry Blog

Although this model constituted the cornerstone in the study of matter, it was discovered later on that some of the statements were right (4), others half-truth (2),  others wrong (1, 3).

II. Thomson’s Atomic Model

In 1897, the British physicist, J.J. Thomson, in his famous Cathode ray experiment (Fig. 1a), discovered the existence of negatively charged particles, he called them “electrons”. He also found that the electron is found in any material; hence the electron is a universal particle. This proved wrong the Dalton’s first statement. Thomson’s experiment showed that the atom has two parts, a part made of a positive mass in which negative electrons are embedded. This new atomic model was called the “Plum pudding model” (Fig. 1.b)


Figure 1a: Cathode Ray Experiment
Source: Study.com


Figure 1b: JJ Thomson Atomic Model (Plum pudding atomic model)
Source: classnotes.org.in

III. Rutherford’s Atomic Model

In 1911, Ernest Rutherford, New Zealand-born British physicist, a former student of J.J. Thomson, after his gold foil experiment (Fig.2), proposed a new atomic model.

Figure 2: Gold Foil Experiment
Source: padakshep.org

The model consisted in a very dense positively charged nucleus, with electrons orbiting the nucleus similar to how planets turn around the sun. This model was then called the “Planetary atomic model” (Fig.3).

In that model, almost the total mass of the atom is concentrated in the nucleus, surrounded by an empty space occupied by the tiny electrons revolving around the nucleus.


Figure 3: Rutherford's Atomic Model (Planetary Atomic Model)
Source: sutori.com

But this model couldn’t explain why the electrons, negatively charged, wouldn’t be attracted by the positively charged nucleus and spiral into the nucleus. According to James C. Maxwell, “An electron that is accelerating radiates energy. As it loses energy, it spirals in to the nucleus”. Hence the atom proposed by Rutherford couldn’t be stable!

IV. Bohr’s Atomic Model

In 1913, in order to solve the problem raised by Rutherford’s atomic model, Niels Bohr introduced the concept of quantization of atomic energy levels.

In his atomic model, Bohr proposed the existence of many different energy levels around the nucleus called “orbits”. Electrons can only turn/revolve on those allowed energy levels or orbits. Those energy level were given numbers: n1 = 1, n2= 2, n3 = 3...  n∞.

Only if an electron receives the appropriate energy corresponding to the difference between two energy levels, ΔE = En2 – En1= hν, then it can jump (excited state) from n1 level to n2level. But since the excited state is not stable, the electron will return back to the non-excited state or ground state by emitting the absorbed energy.

The Bohr’s atomic model helped in explaining the absorption and emission atomic spectra of hydrogen atom.


Formation of the Absorption and Emission Spectra of Hydrogen
Source: intl.siyavula.com

You have certainly observed the emission phenomenon when you drop willingly or accidentally some crystals of salt (NaCl) in a blue flame: a very brilliant yellow flame is observed, which is the emission flame of Sodium atom.

Emission Flame of Sodium
Source: thoughtco.com


Flame Test
Source: dornsife.usc.edu

Although Bohr’s atomic model helped in explaining some phenomena and behaviors of the atom, it has its own weaknesses:

  • This model applies well for Hydrogen atom, the simplest atom of the chemical elements, made of 1 proton and 1 electron; it couldn’t apply for multi-electron (more than 1 electron) species
  • It is against the Heisenberg’s Uncertainty Principle, since in this model, the electron can be localized at any point of the orbit. The uncertainty principle states that it is not possible to know with high accuracy both the position and momentum of a moving particle.

V. Quantum Mechanical Model of the Atom

This model was introduced by Erwin Schrodinger in 1926. Schrodinger’s model considers that an electron cannot be localized precisely on an orbit or point due to Heisenberg’s Uncertainty Principle. But only the probability of finding an electron in a certain region can be estimated.

Schrodinger introduced the concept of an “orbital”, a limited region around the nucleus where there is  high probability (>95%) of finding the electron.

In this model, an electron in an orbital is described by 4 quantum numbers:

  1. The principal quantum number: n = 1, 2, 3, 4..., gives the main or principal energy level. Traditionally those energy levels have been named by letters: K, L, M, N .
  2. The orbital quantum number, l: with values: (n – 1), (n – 2) ... 1, 0. It shows the angular moment of the electron. It gives the shape of the orbital. Traditionally the orbital quantum numbers have been given specific letters to identify them: s(l = 0), p(l = 1), d(l = 2), f(l = 3), etc.
  3. The magnetic quantum number, (ml): ml = + l, +(l – 1), ....1 , 0, -1, .... –(l -1), -l, which governs the energies of electrons in external magnetic fields. It gives the orientation of the orbital.
  4. The spin of electron: an electron can spin around its axis, clock or anti-clockwise; that quantum number, represented by the symbol s, indicates the spinning movement of the electron; it can take two values: s = +1/2, -1/2, sometime represented by the signs ↑, ↓.

The first 3 quantum numbers describe an orbital in terms of the principal quantum number, its shape and its orientation.


s, p, and d Orbitals
Source: chemsite.lsrhs.net

VI. Chadwick

In May 1932, James Chadwick announced that the atomic nucleus contains a new uncharged particle, which he named the ‘neutron”. This discovery helped to explain the existence of Isotopes.

Isotopes are atoms of a same element, i.e. same number of protons and electrons, but different number of neutrons. This proved wrong one of Dalton’s statement that all atoms of the same element are identical.

This discovery concluded more than 1 1/3 century of research on the composition and the structure of the atom, made of:

  • Nucleus: the center of the atoms where Protons(p) positively charged and Neutrons (n) with no charge are found.
  • Electrons(e), negatively charged that surround and move around the nucleus.
  • In an atom, the number of protons is equal to the number of electrons; that is why an atom is neutral.

Representation of an atom:

X = Chemical symbol of the element

A = Mass number, equal  to number of protons + number of neutrons

Z = Number of protons or atomic number, equal to the number of electrons