Theories and models

of the Atom

Theories and models of the Atom (I.I) 1

What do you know about chemistry? Have you ever wondered what everything around you is made of – stone, metal, air and water, even living creatures? People started asking that question thousands of years ago. Nowadays we can give an answer: everything is built up from atoms and in our world we find less than a hundred different kinds of atom. Atoms of any one kind form a chemical element and when these elements are put together, in countless different ways, they give us all known forms of matter.

Read the article and then answer the questions that follow

Answer these questions:

Who thought that atoms were small indestructible particles?

What model was developed to explain the existence of electrons?

What sub-atomic particle is thought to be in the nucleus according to the article?

What sub-atomic particle believed now to be in the nucleus is not mentioned in the article?

HISTORY OF THE ATOM

The name “atom” was first coined by the ancient Greek scientist Democritus. Much later in 1803, an English school teacher named John Dalton proposed the idea that all matter is made of atoms. Dalton thought that atoms were small, indestructible particles. Then in 1897, electrons were discovered. This meant that atoms were made of even smaller sub-atomic particles. In 1904, J.J. Thomson developed the “plum-pudding” model of the atom which suggested that the atom was a sphere with electrons embedded in its surface.

Seven years later, a New Zealand physicist named Ernest Rutherford found through experimentation that the atom was mostly empty space with a very, small but dense centre. His proposal was that the atom contained a nucleus of protons surrounded by a large area containing electrons. But there should be another kind of particles in the nucleus to account for the extra mass in it….

In 1923, the Japanese scientist, Nagaoka, suggested that electrons were moving in orbits around the nucleus in a way similar to that of the planets around the sun.

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1. Models and particles of the atom

People began wondering about matter more than 2500 years ago. Some of the early philosophers thought that matter was composed of tiny particles. They named these particles “atoms”, a term that means “cannot divided”.

The early philosophers didn’t try to prove their theories by doing experiments as scientists now do. Their theories were a result of reasoning, debating, and discussion-not evidence or proof. Today, scientist will not accept a theory that is supported by experimental evidence. But even if these philosophers had experimented, they could not have proven the existence of atoms. People had not yet discovered much about what is now called chemistry, the study of matter. The kind of equipment needed to study matter was a long way from being invented. Evan as recently as 500 years ago, atoms were still a mystery.

A long period passed before the theories about the atom were developed further. Finally during the eighteenth century, scientists in laboratories began debating the existence of atoms once more. Chemist were learning about matter and how it changes. They were putting substances together to form new substances and taking substances apart to find out what they were made of. They found that certain substances couldn’t be broken down into simpler substances. Scientists came to realize that all matter is made up of elements. An element is matter made of atoms o only one kind. For example, iron is an element made of iron atoms. Silver, another element, is made of silver atoms. Carbon, gold and oxygen are other examples of elements.

1. Dalton’s model

John Dalton, an English schoolteacher in the early nineteenth century, combined the idea of elements with the earlier theory of the atom. He proposed the following ideas about matter:

Matter is made up of atoms Atoms cannot be divided into smaller pieces All the atoms of an element are exactly alike: same mass and same properties Different elements are made of different kinds of atoms Combination of atoms of different elements form compounds

Dalton pictured an atom as a hard sphere that was the same throughout, something like a tiny marble.

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1.1. The electron and Thomson’s model

Dalton’s theory of the atom was tested in the second half of the nineteenth century.

In 1897, J.J. Thomson showed that there were negatively charged particles smaller than an atom. They were later called electrons.

J. J. Thomson constructed a glass tube which was partially evacuated i.e. much of the air was pumped out of the tube. Then he applied a high electrical voltage between two electrodes at either end of the tube. He detected that a stream of particle (ray) was coming out from the negatively charged electrode (cathode) to positively charged electrode (anode). This ray is called cathode ray and the whole construction is called cathode ray tube.

Some of the questions posed by scientists were answered in light of Thomson’s experiments. However, the answers inspired new questions. If atoms contain one or more negatively charged particles, then all matter, which is made of atoms, should be negatively charged as well. But all matter isn’t negatively charged. Could it be that atoms also contain some positive charge? The negatively charged electros and the unknown positive charge would then neutralize each other in the atom. Thomson came to this conclusion and included positive charge in his model of the atom.

Using his new findings, Thomson revised Dalton’s model of the atom. Instead of a solid ball that was the same throughout, Thomson pictures a sphere of positive charge. The negatively charged electrons were spread evenly among the positive charge.

Therefore, the atom is neutral. It was later discovered that not all atoms are neutral. The number of electrons within an element can vary. If there is more positive charge than negative electrons, the atom has an overall positive charge. I there are more negative electrons than positive charge, the atom hs an overall negative charge.

1.2. The proton and Rutherford’s model

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In his experiments, Rutherford showed that positive charge existed in a small region of the atom. He hypothesized that all the mass of the atom and all of tis positive charge are crammed into an incredibly small region of space at the center of the atom which he called the nucleus. Eventually, his prediction was proved true. In 1920 scientists identified the positive charges in the nucleus as protons. A proton is a positively charged particle present in the nucleus of all atoms. The rest of each atom is empty space occupied by the atom’s almost massless electrons.

Most of the mass of the atom and its positive charge is in the nucleus of the atom

Electrons are orbiting the nucleus

**Other theories say that it was E. Goldstein, a German physicist, in 1886, who discovered the proton in cells. He made this discovery by using a hydrogen gas-filled tube, which was similar to Thomson's tube. These positively charge particles were called protons, and their mass is 1.837 times the mass of the electron.

1.3. The neutron

Rutherford predicted (in 1920) 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 into bits 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, a British physicist, discovered the neutron and measured its mass: the neutron has the same mass as a proton and is electrically neutral.

Particle Symbol Electrical charge Mass

Electron e- -1,602.10-19 C 9,109.10-31 kg

Proton p+ +1,602.10-19 C 1,673.10-27 kg

Neutron n 0 1,673.10-27 kg

1.4. Bohr’s model

There were, however, some problems with Rutherford's model: for example it could not explain the very interesting observation that atoms only emit light at certain wavelengths or frequencies. Niels Bohr solved this problem by proposing that the electrons could only orbit the nucleus in certain special orbits at different energy levels around the nucleus.

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So the Bohr´s model for some elements it is:

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1.5. The quantum mechanical model

Although the most commonly used model of the atom is the Bohr model, scientists are still developing new and improved theories on what the atom looks like. One of the most important contributions to atomic theory (the field of science that looks at atoms) was the development of quantum theory. Schrodinger, Heisenberg, Born and many others have had a role in developing quantum theory.

The quantum mechanical model is based on quantum theory, which says matter also has properties associated with waves. According to quantum theory, it’s impossible to know the exact position and momentum of an electron at the same time. This is known as the Uncertainty Principle.

The quantum mechanical model of the atom uses complex shapes of orbitals (sometimes called electron clouds), volumes of space in which there is likely to be an electron. So, this model is based on probability rather than certainty.

Four numbers, called quantum numbers, were introduced to describe the characteristics of electrons and their orbitals: Principal quantum number, n; Angular momentum quantum number, l; Magnetic quantum number,m1; Spin quantum number, ms.

In the next two figures you can see:

The relative size of the s orbitals in the first three levels

The shapes of the s, p, and d orbitals.

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In each orbital there can be only two electrons.

There are 7 energy levels, and four types of orbitals: s, p ,d, f. They have different shape and orientation is space.

The distribution of electrons in the first 4 energy levels would be as follows:

Energy level (n) 1 2 3 4

Sublevel s s p s p d s p d f

Number of orbitals 1 1 3 1 3 5 1 3 5 7

Name 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f

Maximum number of electrons in each sublevel

2 2 6 2 6 10 2 6 10 14

Maximum number of electrons in each energy level

2 8 18 32

1.6. Atomic number and mass number

Atomic number, Z, is the number of protons in the nucleus of an atom of an element

The electrons in an atom have almost no mass. So the mass of an atom is nearly all due to its protons and neutrons. For this reason, the number of protons and neutrons in the nucleus of

an atom of an element is called the mass number, A, and A = Z + N

Example: A sodium atom has 11 protons and 12 neutrons, so the mass number of sodium is 23.

Shorthand of an atom: The sodium atom can be described in a short way, using:

The symbol for sodium (Na); its proton number (11); its mass number (23) 2311Na

X: symbol of the element; A: mass number; Z: atomic number

Since the atom is neutral, the number of electrons and protons is the same.

Isotopes

Although the number of protons changes form element to element, every atom of the same element has the same number of protons. However, the number of neutrons can vary even for one element. The different types of atoms are called isotopes, which are atoms of the same element that have different numbers of neutrons.

You can tell someone exactly which isotope you are referring to by using its mass number. An atom’s mass number is the number of protons plus the number of neutrons it contains. Hydrogen has three isotopes with mass numbers of 1,2 and 3. Each hydrogen atom always has one proton, but in each isotope the number of neutrons is different.

11H 2

1H 31H

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Activities: 1. The atomic number of an atom is Z=6. If its number of neutrons is 7, what is its mass

number? How many protons are in the atom? How many electrons?

2. Name each of these atoms and say how many protons, electrons and neutrons it has:

5626 Fe 1

1 H 2612 Mg 39

19 K 4018 Ar 16

8 O

3. Finish this sentence: Two atoms are isotopes if …………

4. Are these sentences true or false?

a) According to Thomson’s model of the atom, protons can be found all over the atom. b) Protons and electrons have fixed positions in the atom. c) The radius of the atom and the radius of the nucleus is the same. d) According to Rutherford’s model of the atom, electrons move around the nucleus of

the atom as planets do around the Sun. e) Neutrons have no mass f) According to Bohr’s model of the atom, electrons behave like waves.

2. Electronic structure or electron configuration

The number of electrons in the highest or outermost energy level determines an atom´s chemical properties.

Predicting the chemical properties of atoms of different elements using the number of electrons is very important. It gives us information, for example, about the type of bonds they will form.

To do so, we need to know the electron configuration.

The electron configuration is the way the electrons are distributed in the shell of an atom.

Rules for filling orbitals:

Electrons are arranged in different shells around the nucleus.

Atomic orbitals fill from least to most energy; the innermost shell - or lowest energy level - is filled first.

Each succeeding shell can only hold a certain number of electrons before it becomes full. Remember that each orbital can hold a maximum of 2 electrons.

We need to know the orther of the orbitals in terms of energy, as shown in the Moeller´s diagram figure.

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As orbitals with the same energy level fill up, the next ones begin filling. In the next picture each arrow represents an electron. See how there are two arrows, or electrons, for each full orbital.

Electrons occupy the maximum number of orbitals at any given energy level.

This is known as Hund´s rule of maximum multiplicity.

Remember that the number of electrons in a neutral atom equals the number of protons, so in order to know how many electrons are in an atom, the atomic number of the element will be used.

Let’s see more examples:

Li (Z=3) 2 e- in the first energy level and 1 e- in the second one. 1s2 2s1

C (Z=6) 2 e- in the first energy level and 4 e- in the second one 1s2 2s2 2p2

Examples

Z= 11 1s2 2s2 2p6 3s1

Z= 27 1s2 2s2 2p6 3s2 3p6 4s2 3d7

Z=54 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6

Activities:

5. Choose 5 elements from the periodic table, write their name and give their electronic structure.

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An element´s properties come from its electron configuration, but specifically from its valence electrons.

The electrons in the last occupied level are called the valence electrons.

For magnesium, whit electron configuration [Mg] = 1s2 2s2 2p6 3s2, the electons in the 3s orbital, the las tilled shell, are its valence electrons.

Is an atom that loses electrons and becomes positively charged.

Ex.- Mg2+

Is an atom that gains electrons and becomes negatively charged.

Ex.- Cl-

Activities:

6. Given the atomic number of germanium (Z=32):

a. Write the electron configuration of the neutral isotope of germanium

b. Specify the valence electrons.

c. Write the electron configuration of the cation Ge2+

7. Write the electron configuration of the K neutral atom and K+ cation.

8. Write the electron configuration of the neutral isotopes of fluorine (Z=9), chlorine (Z=17) and bromine (Z=35). What do you conclude?

9. Explain why the n=4 level may have up to 32 electrons.

10. Explain the number of electrons in an atom with the following electron configuration:

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4

Which element is it?

11. Explain the difference between an orbit and an orbital.