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Unit 1-1

Unit 1Section 1.1(1)

Atoms, Molecules and StoichiometryThe Atomic Structure

Protons, neutrons and electrons as constituents of the atom

Atoms and molecules The first chemist to use the name atom was John Dalton (1766-1844). Dalton used the word atom to mean the smallest particle of an element. He then went on to explain how atoms could react together to form molecules : An atom is the smallest part of an element which can ever exist. A molecule is the smallest part of an element or a compound which can exist alone under ordinary conditions. Electrons In 1897, J. J. Thomson investigated the conductivity of electricity by gases at very low pressure. At ordinary pressures gases are electrical insulators, but when they are subjected to very high voltages at very low pressures (below 0.01 atm) they break down and conduct electricity. When Thomson applied 15000 volts across the electrodes of a tube containing a trace of gas, a bright green glow appeared on the glass. The green glow results from the bombardment of the glass by rays travelling in straight lines from the cathode. Thomson called these rays cathode rays. He also showed that when the rays were deflected by an electric field across a pair of charged plates, the rays moved away from the negative plate towards the positive plate. This suggested that the rays were negative. Thomson studied the blending of a thin beam of cathode rays by magnetic and electric fields and concluded that they consisted of electrons tiny negatively charged particles.

The same results were obtained with the cathode rays using different gases in the tube and with tubes and electrodes of different materials. This suggested that electrons were present in the atoms of all substances.

Unit 1-2

Rutherfords model of atomic structure In 1911, Ernest Rutherford had the idea of probing inside the atom using alpha-particles. He used alpha-particles from radioactive substances as nuclear bullets. Alpha-particles are helium ions, He2+, with positive charge. Rutherford expected that most of the very fast-moving alpha-particles would pass straight through the thin metal foil or be deviated a little. Most of the alpha-particles did pass straight through the foil, but when the detecting screen and microscope were rotated from the straight-on position flashes could still be seen. Clearly, some of the alpha-particles were deflected by the foil, but to every ones surprise, one particle in every 10000 appeared to rebound from the foil.

Rutherford suggested that deflections and reflections could only be caused by the particles coming close to a concentrated region of positive charge.

Rutherford concluded that atoms in the metal foil consisted of a central positive nucleus composed of protons, where the mass of the atom was concentrated. This nucleus was surrounded by a much larger volume in which the electrons move. From the angles through which alpha-particles are deflected, Rutherford calculated that the nucleus of an atom would have a radius of about 10-14 m. This is about one ten-thousandth of the size of the whole atom which has a radius of about 10-10 m.

Unit 1-3

Neutrons In spite of the success of Rutherford in explaining atomic structure, one major problem remained unsolved. If the hydrogen atom contains one proton and the helium atom contains two protons, then the relative atomic mass of helium should be twice that of hydrogen. Unfortunately, the relative atomic mass of helium is four and not two. In 1932, Chadwick, one of Rutherfords collaborators, was able to show where the extra mass in helium atoms came from. Chadwick bombarded a thin sheet of beryllium with alpha-particles. The alpha-particles can be traced by electric counter which detects charged particles. When the beryllium is in place, the counter registers nothing, showing that the alpha-particles are being stopped by the beryllium. However, if a piece of paraffin wax is placed between the beryllium and the counter, charged particles are detected again.

Chadwick provided an explanation. He suggested that the alpha-particles striking the beryllium foil displaced uncharged particles called neutrons from the nuclei of beryllium atoms. These uncharged neutrons could not affect the charged-particle counter, but they could displace positively charged protons from the paraffin wax which would affect the counter.

Further experiments showed that neutrons had almost the same mass as protons and Chadwick was able to explain the difficulty concerning the relative atomic masses of hydrogen and helium. Hydrogen atoms have one proton, no neutrons and one electron. Since the mass of the electron is negligible compared to the masses of the proton and neutron, a hydrogen atom has a relative mass of one unit. Helium atoms have two protons, two neutrons and two electrons, so the relative mass of a helium atom is four units. This means that a helium atom is four times as heavy as a hydrogen atom.

Unit 1-4


The relative masses and charges of a proton, neutron and electron

Nowadays scientists believe that all atoms are composed of three important sub-atomic particles : protons, neutrons and electrons. Sub-atomic particle Proton Neutron Electron Mass / Kg 1.6726 x 10-27 1.6750 x 10-27 9.1095 x 10-31 Relative mass (to that of a proton) 1 1 1 1836 Charge /C + 1.6022 x 10-19 0 - 1.6022 x 10-19 Relative charge (to that on a proton) +1 0 -1


The atomic nucleus

The nucleus of an atom is composed of protons and neutrons. Because protons and neutrons occupy the nucleus, they are sometimes collectively called nucleons. Virtually all the mass of the atom is concentrated in the nucleus, which occupies only a small fraction of the total volume of the atom. The neutron has no charge, whereas the proton carries one positive charge. Electrons with one negative charge occupy the space outside the nucleus. The mass of an electron is 1836 times less than that of a proton. The atomic number or the proton number of an element is the most important feature of an elements individuality because it represents (i) the number of protons in the nucleus, (ii) the number of electrons in the neutral atom, (iii) the position in which the element appears in the periodic table. The number of protons + the number of neutrons in an atom is called the mass number or the nucleon number. The word nuclide is used to describe any atomic species of which the atomic number and the mass number are specified. The symbol AZ X is used to represent the nuclide X with atomic number Z and mass number A. Example : Particle/Atom Symbol


Proton 1 1H or 1p

Neutron 1 0n

Electron 0 -1e

Hydrogen 1 1H

Helium 4 2He

Atoms of the same element with different masses are called isotopes. All the isotopes of one particular element have the same atomic number because they have the same number of protons, but they have different mass numbers because they have different numbers of neutrons. Isotopes have the same number of electrons and hence the same chemical properties, because chemical properties depend upon the transfer and redistribution of electrons. As isotopes have different number of neutrons, they have different masses and hence different physical properties. For example, pure 3717Cl2 has a higher density, higher melting point and higher boiling point than pure 3517Cl2. Example : Hydrogen have three isotopes : hydrogen-1, hydrogen-2 and hydrogen-3. Write their symbols.

Example : The nucleus of a fluorine atom has a diameter of about 1.0 x 10-12 cm and a mass of 3.1 x 10-23 g, calculate the density of the fluorine nucleus.

Unit 1-5

Section 1.2(1)

Relative isotopic, atomic and molecular masses

Relative isotopic and atomic masses

Carbon-12 scale Chemists use a relative atomic mass scale to compare the masses of different atoms. In 1961 carbon-12 (12C) was chosen as standard against which the masses of other atoms were compared Carbon-12, an isotope of carbon, has been assigned a relative atomic mass of exactly 12. This scale is called the carbon-12 scale. The isotope of carbon was chosen because carbon is a very common element. Being a solid, it is easy to store and transport. Relative isotopic mass Relative isotopic mass of a particular isotope of an element is the relative mass of one atom of that isotope on the carbon-12 scale. Mass of one atom of an isotope of an element Relative isotopic mass = 1 Mass of one atom of carbon - 12 12 Atomic mass unit

1 of 12 the mass of the carbon-12 atom. Thus one atom of C-12 weighs 12.000 a.m.u. and its relative isotopic mass is 12.000. Example : The mass of a carbon-12 atom is 1.9926 x 10-26 kg. What is the mass of 1 a.m.u. in kg ? Relative isotopic mass is a ratio, it has no unit. An atomic mass unit (a.m.u.) is taken as

On the carbon-12 scale, the mass of the proton (1.0074 a.m.u.) is almost the same as that of the neutron (1.0089 a.m.u.), and the mass of the electron is very small in comparison (0.0005 a.m.u.). Now since the relative masses of the proton and neutron are very close to one and the electron has a negligible mass, it follows that all relative isotopic masses will be very close to whole numbers. In fact, the relative isotopic mass of an isotope will be very close to its mass number and the two are assumed to be almost identical in all but the most accurate work. Example : Symbol Relative isotopic mass Mass / a.m.u.12 16 17 18 35 37

C 12.000

O 15.995

O 16.999

O 17.999

Cl 34.969

Cl 36.966

Relative atomic mass Naturally occurring elements often consist of a mixture of isotopes. Relative atomic mass ( Ar ) of an element is the weighted average of the relative isotopic masses of the natural isotopes of an element on t


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