Matter and Introduction to the atom
The traditional scientific definition of mass is that it is anything that has mass and takes up space. In everyday language, Matter is stuff → the chair you’re sitting on, your phone, your desk, the air around you.
All matter, that is all stuff is made up of 118 elements. Water for instance is made up of the elements Hydrogen and Oxygen. The elements are classified according to their characteristics into the Periodic Table:
A wall is built up from bricks. Similarly, matter is made up of atoms. Atoms are the building block of matter. They are the smallest particle that retains the chemical properties of the element. Each element is comprised of atoms that are unique to that element; hydrogen atoms are different from gold atoms, which are different from neon atoms.
The history of our understanding of the atom
Here are some significant milestones in our understanding of the atom.
1400 BC – Democritus (Greek Philosopher)
The story of the atom begins about 1 400 years before Jesus Christ in the days of the Greek philosopher Democrotus (470-380 BCE) who proposed that all matter was made up of smaller particles he dubbed “atomos”.
The 1800s – John Dalton’s Atomic Model
In 1808 John Dalton, an English school teacher created the first model atomic model. His atomic theory consisted of this points:
- All matter is made up of tiny, indivisible particles he called atoms. He imagined them as hard with a distinct mass and movable. They could neither be created nor destroyed.
- All atoms of the same element are identical. Identical in terms of mass and also properties.
- Atoms of different elements are different.
- Compounds are formed by combining atoms of different elements in fixed whole numbered ratios.
- Chemical reactions occur when atoms are rearranged.
Of course we now know that the statement that atoms are indivisible does not hold true.
The 1890s – J.J Thompson
J.J Thomson was a physicist who is credited for discovering the electron. This showed that the atoms were not indivisible as Dalton had suggested.
J.J. Thomson’s experiments with cathode ray tubes showed that all atoms contain tiny negatively charged subatomic particles or electrons.
J.J Thompson’s Plum Pudding Model
It was known that atoms had an overall neutral charge. Therefore, Thomson reasoned that there must be a source of positive charge within the atom to counterbalance the negative charge on the electrons.
This led Thomson to propose that atoms could be described as negative particles floating within a soup of diffuse positive charge. This model is often called the plum pudding model of the atom, due to the fact that its description is very similar to plum pudding, a popular English dessert.
The 1910s – Ernest Rutherford
In 1911, Rutherford and his colleagues Hans Geiger and Ernest Marsden initiated a series of groundbreaking experiments that would completely change our understanding of the atom. They bombarded very thin sheets of gold foil with fast moving alpha particles. This experiment is known as the gold foil experiment.
1. They fired alpha particles (positively charged) at a gold foil.
2. They measured the deflection as the particles came out the other side.
3. Most of the particles did not deflect at all. Every now and then a particle would deflect all the way back.
4. He said that there must be a positive centre of the foil. He called this centre the nucleus.
Since the majority of the alpha particles had passed through the gold, Rutherford reasoned that most of the atom was empty space. In contrast, the particles that were highly deflected must have experienced a very powerful force within the atom. He concluded that all of the positive charge and the majority of the mass of the atom must be concentrated in a very small space in the atom’s interior, which he called the nucleus.
Rutherford’s Atomic model
- The nucleus of the atom is a dense mass of positively
- The electrons orbit the nucleus, much like the planets orbit the sun.
The 1910s – Niels Bohr
The problem with Rutherford’s model was that it could not explain why the negatively charged electrons were not attracted to the positively charged protons in the nucleus. Bohr explained this, and because his model is a modification of Rutherford’s model, it is sometimes referred to as the Rutherford-Bohr model.
Bohr postulated that:
- Electrons orbit the nucleus in orbits that have a set size and energy.
- The lower the energy of the electron, the lower the orbit → this means that electrons that orbit near the centre of the atom have lower energy than those that orbit further away.
- As electrons fill up the orbitals, they will fill the lower energy level first.
- If that energy level is fill (or at capacity), a new energy level will begin.
- Radiation is when an electron moves from one energy level to another.
The Bohr model and all modern models describe the properties of atomic electrons in terms of a set of allowed (possible) values. Atoms absorb or emit radiation only when the electrons abruptly jump between allowed, or stationary, states. By limiting the orbiting electrons to a series of circular orbits having discrete radii, Bohr could account for the series of discrete wavelengths in the emission spectrum of hydrogen.
Bohr proposed that light radiated from hydrogen atoms only when an electron made a transition from an outer orbit to one closer to the nucleus. The energy lost by the electron in moving from a high to lower energy state is precisely the same as the energy of the quantum of emitted light. According to the first law of conservation, energy has to be conserved. So, the energy lost from the electron moving orbitals is emitted as a photon.
Rutherford predicted that another kind of particle must be present in the nucleus along with the proton. He predicted this because if there were only positively charged protons in the nucleus, then it should break up because of the repulsive forces between the like-charged protons. To make sure that the atom stays electrically neutral, this particle would have to be neutral itself. In 1932 James Chadwick discovered the neutron and measured its mass.
Schrödinger’s atomic model
Erwin Schrödinger used Heisenberg’s uncertainty principle to come up with the atomic model that we still use today.
Schrödinger’s model describes the electrons belonging to an atom in terms of four quantum numbers from which the wave function for each electron can be calculated. Instead of electrons orbiting a nucleus in an assigned place at a given time, the wave mechanical model provides a probability of where an electron may be found. For example, the probability distribution curve to the right is for hydrogen in its ground energy state. It indicates that while the electron is found on average at the Bohr radius (r1), its position may be further from or closer to the nucleus at any given time.
To sum it up:
- An electron does not travel in an exact orbit
- We can predict where it will probably be
- We cannot say for certain where it is, but only where it ought to be.
- The type of probability orbit is dependent on the energy level described by Bohr