In chemistry and physics, atomic theory is a scientific theory of the nature of matter, which states that matter is composed of discrete units called atoms.
The word “atom” from the ancient Greek adjective atomos, means indivisible.
The idea that all matter is made up of tiny particles called atoms has been in existence at least since the time of Greek philosopher Democritus (460-370 BC). However, little has been done to use ideas about atoms to explain the behavior of chemicals until John Dalton (1766-1844) developed his atomic theory.
In 1803 Dalton proposed that;
All elements are made up of small indivisible particles called atoms.
Atoms cannot be created or destroyed
The atoms of each individual element have the same size, and the same weight (atomic weight)
When elements combine together, their atoms join together in fixed proportions.
Dalton’s atomic theory was based on experimentation and chemical laws known at that time, assigned weights and combining capacities to the postulated atoms. The theory in its broad outline is still valid, however some of the particulars have been modified in the light of modern discovery, but the principal aspects as outlined are still useful in the study of chemistry.
Dalton’s atomic theory was a landmark event in the history of chemistry, but it had a crucial flaw. His procedure requires knowing the formulas of the simple compounds resulting from the combination of the elements.
In the following years, several leading chemists adopted essential elements of Dalton’s theory, but many objected to the hypothetical elements described.
In 1808, the French chemists Joseph-Louis Gay-Lussac discovered that when gasses combine chemically, they do so in small integral multiples by volume. Three years later the Italian physicist Amedeo Avogadro argued that this fact suggested that equal volumes of gases contain equal numbers of constituent particles (Avogadro’s law), physical conditions being the same. This idea provided a physical method of determining certain molecular formulas. However, many chemists rejected Avogadro’s hypotheses.
The greatest of the early atomists was the Swede Jons Jacob Berzelius, who accepted parts of Avogadros ideas and developed an elaborate of chemical atomism by 1826. It was Berzelius who in 1813 had proposed the alphabetic system for denoting elements, atoms and molecular formulas, and the use of formulas as an aid for studying chemical composition and reactions. He was one of the first to determine the atomic weights of elements.
THE MODERN ATOMIC THEORY
The modified Dalton’s atomic theory is referred to as the modern atomic theory and consists of the following;
All matter is composed of small particles, some of which are electrically neutral, some are positively charged while some are negatively charged.
Atoms of the same element are not alike, but may have different masses
An element may have atoms with different masses. These atoms of the same element with different masses are called Isotopes.
Atoms of different elements can combine together to form molecules
The molecules of a compound have definite compositions and structures.
The study of the structure of atom was investigated by so many scientists such as J.J Thompson (1897) who discovered the electron through his work on cathode rays, Henry Moseley (1913) discovered the atomic numbers of atoms i.e number of protons in the nucleus, which determine the order of elements in the periodic table.
James Chadwick (1932) discovered the neutron while R.A Millikan (1913) determined the charge of electron using the oil drop experiment.
Ernest Rutherford discovered the proton using the gold foil experiment. The work of Sir Rutherford holds lots of evidence because of his ability to modify some of the proposed theory of John Dalton, showing that atoms are not the smallest indivisible particles of an element but are composed of smaller particles such as the protons, electrons and neutrons.
Cathode Rays
It was Sir J.J Thomson who, in 1897 was credited with identifying cathode rays as the things we call electrons. Thomson used an apparatus which consists of a glass tube containing a gas at a very low pressure. He placed the cathode in front of a short metal cylinder, which forms the anode. At the far end of the tube was a fluorescent screen. In the middle of the tube is a parallel metal plate. A source of very high voltage is connected to the anode and cathode. When this happens, the fluorescence screen begins to glow. At that time, Thomson was able to prove that the glow was due to some of the rays which travelled in a straight line from the cathode. He called these rays’ cathode rays.
Thomson’s model of the atom
Thomson’s atomic model was proposed in 1904 and was called the Plum Pudding model. It was introduced right after Thomson’s discovery of the electron. Thomson said that the atom is a sphere of positively charged matter in which negatively charged electrons are embedded. Thomson’s model of the atom resembles a pudding studded with currants.
Diagram of Thomson’s model of the atom
Thomson’s model was proved wrong by the Rutherford’s gold foil experiment.
Millikan Oil Drop Experiment
Robert .A. Millikan, 1909, from his experiment determined the charge on the electron. The method used by R.A. Millikan was known as Millikan’s oil drop experiment. The idea was to allow tiny droplets of oil to fall through the air between two plates. X-rays passed into the apparatus caused molecules in the air to ionize. From time to time the ions would stick to the oil drops. The metal plates were given an electric charge, and, as the electric field between the plates was increased, it was possible to make some of the drops travel upwards at the same speed as they were previously falling. By measuring the speed, and knowing the strength of the field and density of the oil, Millikan was able to calculate the magnitude of the charge on the oil drops.
Moseley and Atomic Number
The evidence that there is a positive charge in the nucleus was of fundamental importance, which was provided by H.G.J Moseley. In the experiments he performed in 1913, Moseley bombarded a number of elements with cathode rays (electrons). The energy provided by the cathode rays caused the elements to give off x-rays. He used x-rays tubes to determine the charges on the nuclei of most atoms. He wrote that the atomic number of an element is equal to the number of protons in the nucleus. In 1914, Moseley was able to determine the order of elements in the Periodic table using the atomic number, which we recognize as the number of protons in the nucleus. His work was used to reorganize the periodic table based on atomic number instead of atomic mass.
Rutherford’s Model of Atomic Structure
The work of Sir Ernest Rutherford holds lots of evidence because of his ability to modify some of the proposed theory of John Dalton, showing that atoms are not the smallest indivisible particle of an element but composed of smaller particles such as the protons, electrons and neutrons. He said, the electrons moved in orbits around the nucleus and were held in their orbits by electrostatic attraction to the positively charged nucleus.
From his explanation of the model of the atom he showed that the structure of an atom can be view as the solar system which consisted of the sun and the nine planets revolving round it. In his view the sun was seen as the nucleus of the atom where all the atomic energy is packed up; while the planets where seen as electrons revolving round the nucleus hence, his theory showed that that the movement of the electrons could be found anywhere round about the nucleus. From the results of the experiments, Rutherford also had information on the size, charge and mass of the nucleus. However, he had no information on the distribution or position of electrons.
Diagram of Rutherford’s model of atom
From the results of his experiments, Rutherford concluded that the plum pudding model was wrong; he also made the following conclusions:
Atoms have a nucleus, very small and dense, containing the positive charge
The atom consists of mostly empty space
The electrons are attracted to the nucleus.
The theory proposed by Rutherford was later modified by Neils Bohr.
Bohr’s Model of the Atom
Neils Bohr in 1913 used the ideas of the quantum theory by Planck and Einstein to explain the spectrum of hydrogen. In his evidence he showed that the electrons depicted by Rutherford was not in an orbital movement because every electron has the ability of gaining some amount of energy which is quantized as such the electron will continue to rotate until it disappears into the nucleus. Bohr’s evidence showed that the electrons rotate in their orbit as such ground state electrons can gain energy and move to higher energy level which proves how reactions can take place.
Although Bohr’s model of the atom was successful in explaining the observed spectral lines of hydrogen with great accuracy, but could not predict the frequencies of the spectral lines for more complex atoms. Bohr’s model has now been replaced by the wave mechanic model.
Diagram of Bohr’s model of the atom
He made the following assumptions:
That in the atom there are various circular orbits called shells
That the spectral lines are caused by the electrons
Each circular orbit has a definite amount of energy
When an electron emits energy in the form of radiation when it moves from a higher to a lower orbit. This produces a line in atomic emission spectrum.
The difference in energy ∆E between the circular orbits is related to the frequency of radiation as expressed by the Planck’s equation
∆E = E2 – E1 = hv
Where h = Planck’s constant
V = frequency of the radiation emitted
QUANTUM NUMBERS
Quantum number can be defined as a number used when describing the energy levels available to atoms and molecules. In an atom, the electrons are distributed around the nucleus in groups called electron shells or quantum shells.
Bohr explained spectral lines on the basis of electronic configuration and assigned each shell a principal quantum number n.
An electron in atom or ion has four quantum numbers to describe its state.
There are four quantum numbers which are:
Principal quantum number (n)
The subsidiary or azimuthal quantum number (i)
The magnetic quantum number (mi)
The spin quantum number (ms)
Principal quantum number (n): n= 1, 2, 3, 4…………..∞. where n is an integer and the value of n begins from 1 to ∞.It specifies the energy of an electron and the size of the orbital. The orbit closest to the nucleus has the lowest energy. As the value of n increase the energy of the orbit increases. These orbits are called shells and are designated as K, L, M, N etc. The maximum number of electrons in a shell is given as 2n2.
Value of n Name of electron shell No of electrons
1 K 2
2 L 8
3 M 18
4 N 32
The subsidiary or Azimuthal quantum number (i): this describes the possibility of locating an electron within a region in a shell with a certain amount of energy. This has an integral values ranging from 0 to (n – 1) i.e i= 0, 1, 2, 3, -------n-1. The electrons with subsidiary quantum numbers 0, 1, 2 and 3 are usually to as the s, p, d and f electrons respectively.
The magnetic quantum number: this is the third quantum. It describes the orbital of the subshell. The magnetic quantum number is represented by m and the values ranges from -1 to 0 to 1 i.e m= -1-----0--+1, showing the number of orbital present in each energy subshell.
The spin quantum number: it is represented by ms. This describes the spin of an electron on its own axis and may have a value of either +1/2 or -1/2 i.e only two directions of spin are possible.
SIGNIFICANCE OF QUANTUM NUMBERS
It reflects the fact that for two electrons to occupy the same orbital, they must have opposite spin states, and therefore not have the same set of 4 quantum numbers
It tells about the orientation of atomic orbitals which is actually obtained as a resolution due to the application of magnetic or electric field
It gives us the energy state of electron which further helps to know whether electron is closer or farther from nucleus.
ISOTOPES
Isotopes are elements which have the same atomic number but different atomic mass.
Isotopes, from Greek word mean ‘same place’, because all of these forms fit into the same place on the periodic table. After neutrons were discovered, it was realized that isotopes are substances that have the same number of protons in their nuclei, but different numbers of neutrons. These atoms are called Isotopes and the phenomenon is called Isotopy.
Isotopy is not limited to radioactive elements only. Lighter elements always occur naturally as a mixture of isotopes. For example, hydrogen can occur in three isotopes;
Hydrogen – 1 (protium) has one proton and no neutrons in its nucleus
Hydrogen – 2 (deuterium) has one proton and one neutron in its nucleus
Hydrogen – 3 (tritium) is a radioactive isotope with one proton and two neutrons per nucleus. Tritium is present in extremely low concentrations in natural hydrogen, but can be produced in large quantities by nuclear weapons.
Other examples of isotopes are:
16O8 , 17O8 , and 18O8
35Cl17 and 37Cl17
10B5 and 11B5 .
Isotopes of an element, therefore, have the same chemical properties but different physical properties.
Example:
Chlorine exists in two isotopic forms, 35Cl and 37Cl. The relative abundance of 35Cl is 75% and 37Cl is 25%. Calculate the relative atomic mass of chlorine.
Solution
Mass number of the atom with 75% = 35
Mass number of the atom with 25% = 25
R.A.M =75/100 × 35 + 25/100 × 37
= 26.25 + 9.25
= 35.5
Uses of Isotopes
Archeological dating: used to determine the age of older materials from carbon -14 dating and other dating. It only works for materials that have carbon in them like organic items.
An isotope of hydrogen, deuterium is used to make heavy water D2O that is deuterium oxide which acts as a moderator in the nuclear power plants.
Isotopes are used in medicine; these isotopes are added into biological molecules like sucrose, sugar which are injected into humans to look for accumulation which may indicate cancer and other diseases.
Stable isotopes of oxygen, hydrogen, sulphur, nitrogen and carbon are used in environmental and ecological experiments. By using stable isotopes, geochemists can determine the age of the geological material they are studying.
It is also used in forensic science and nuclear energy.
VALENCY
Valency is defined as the combining power of an element. When atoms are combining they make use of what is called valency or combining power.
The more general definition of valence is the number of electrons with which a given atom generally bonds or numbers of bonds an atom forms.
The valency of an element is a property of its atomic structure. E.g sodium (Na)
Diagram of valency of sodium
Hence, the combining power of sodium in any reaction is 1.
Group 1 2 3 4 5 6 7 8
Atoms Na Mg B C N O F Ne
Valency 1 2 3 4 3 2 1 0
Generally, metals exhibit positive valences while non-metals tend to have negative valences. Some elements exhibit more than one valency.
Note:
The valence of an atom helps in determining its oxidation state
It also helps to determine the type of bond that will be formed
It helps to determine the type of chemical reaction that such compounds formed can take part in.
CHEMICAL BONDING
Atoms are the basic building blocks of all types of matter. Atoms link with other atoms through chemical bonds resulting from the strong attractive forces that exist between the atoms.
A chemical bond is an attraction between atoms that allows the formation of chemical substances that contain two or more atoms. The bond is caused by the electrostatic force of attraction between apposite charges. The electrons that participate in chemical bonds are the valence electrons, which are the electrons found in atom’s outermost shell. When two atoms approach each other these outer electrons interact.
The two main types of bonds formed between atoms are ionic and covalent bonds. An ionic bond is formed when one atom accepts or donates one or more of its valence electrons to another atom. A covalent bond is formed when atoms share valence electron.
The atoms do not always share the electrons equally, so a polar covalent bond maybe the result. When electrons are shared by two metallic atoms a metallic bond maybe formed. In a covalent bond, electros are shared between two atoms. The electrons that participate in metallic bonds may be shared between any of the metal atoms in the region.
Note:
When atoms combine it is the electrons on the outermost shell of the atoms that react by being exchanged or shared.
Types of Chemical Bonding
There are two main types of chemical bonding; ionic or electrovalent and covalent bond. Electrovalent and covalent bonds are the most important in bond types of compounds. Other types of chemical combinations include the co-ordinate covalency (dative) bond, hydrogen bond, metallic bond and Van der Waal’s force.
Electrovalent (Ionic) Bonding
Electrovalent bond involves transfer of electrons from metallic atoms to non-metallic atoms during a chemical reaction, i.e donor-acceptor principle.
The metallic atoms after donating their valence electron become positively charged, while the non-metallic atoms become negatively charged after acquiring extra electrons. Examples of electrovalent bonds are Na+Cl-, formation of calcium oxide (Ca2+O2-), magnesium oxide and magnesium chloride.
Properties of Electrovalent (ionic) Compound
They are solids at room temperature and do not vaporize easily
They have high melting and boiling points because of the strong electrovalent bonds between the ions
Ionic compounds readily dissolve in water and other polar solvents
They are good conductors of electricity
They do not dissolve in organic solvents such as toluene, ether, benzene e.t.c.
Covalent Bond
In covalent bonding, electrons are not transferred but are shared. A covalent bond consists of a pair of electrons shared by two atoms. The shared electrons are each contributed by the reacting atoms.
Sharing of electrons occurs between atoms of comparable electronegativities and atoms of the same element. In covalent bonding, molecules and not ions are formed, because the shared electrons may be regarded as revolving in orbits controlled by both nuclei. Examples are formation of methane, nitrogen molecule, H-Cl, H-H.
Properties of Covalent Bond
They have low melting and boiling points
Covalent compounds consist of molecules which have definite shape
Covalent compounds readily dissolve in non-polar organic solvents
Covalent compounds do not conduct electricity i.e they are non-electrolytes.
Co-ordinate Bond
This type of bonding involves the combination of two elements such that only one of the elements donates the entire electron for the bonding.
A co-ordinate bond is formed when one of the reactants possesses a lone pair of electrons. The lone pair of electrons is donated to an atom needing them to complete an electron octet or duplet of great stability. Examples are the formation of ammonium ion and hydronium ion.
MOLE, AVOGADRO’S NUMBER (NA) AND MOLAR VOLUME
A mole of any substance is the amount of it which contains as many elementary particles as there are atoms in exactly 12 grams of carbon-12. The mole therefore is the unit used for measuring the amount of matter in terms of number of particles.
Understanding the mole concept helps to determine the mass quantity or % by mass combination of elements, particles, atoms or molecules in a given substance or chemical reaction.
The number of specified particles present in one mole of any substance is equal to 6.02 x 1023. This is known as the Avogadro’s number or constant. Avogadro’s number is the number of atoms of carbon in 12 grams of carbon 12 and is numerically equal to 6.02 x 1023. Avogadro’s number also represents the number of molecules of any gas that occupies 22.4 dm3 at s.t.p.
The mass of one mole of a substance is the sum of the relative atomic masses of all atoms in its molecular formula, expressed in grams, example:
The molar mass of hydrogen, H2 is:
H + H H2
1.008 + 1.008 = 2.016
1 mole of hydrogen gas weighs 2.016g/mol
Similarly, the molar mass of Na2SO4 = 142g/mol
The molar volume of any gas is the volume occupied by one mole of that gas at a standard temperature and pressure (s.t.p). This value holds anywhere in the world and it is equal to 22.4 dm3.
The molar volume of a gas is obtained by dividing its mass by its density at s.t.p. Example
dm3 of hydrogen gas weighs 0.09g at s.t.p. the molar mass is 2.016g, find its molar volume.
Molar volume = massdensity
= 2.016g0.09g/dm = 22.4dm3
Examples
Calculate the volume of (a). 3.0 moles of oxygen, (b) 50.0g of hydrogen chloride gas at s.t.p (molar volume of a gas is 22.4 dm3, molar mass of HCl = 36.5g)
KINETIC THEORY OF GASES
1. Gases consist of molecules in a constant state of random motion.
2. The pressure of a gas is due to the collisions of the molecules with the walls of the container. The collisions of the gas molecules are perfectly elastic.
3. The actual volume occupied by the gas molecules themselves is negligible relative to the volume of the container.
4. The cohesive forces between the gas molecules are negligible.
5. The temperature of the gas is a measure of the average kinetic energy of the gas particles i.e the kinetic energy depends on the temperature. Solids are determined by the type of bonds or forces holding the particles together. Examples are ice, stone e.t.c.
MATTER
What is matter?
Matter is anything that has weight and can occupy space.
It is of good importance to know that matter is built up of one or more of the following elementary particles atom, molecules and ions. Examples of matter include air, water, animals.
STATES OF MATTER
There are three states of matter namely:
Solid State
Liquid
gaseous
Solid state: They are forms of matter in which their particles are tightly packed together i.e they are not free to move. The particles are held by cohesive forces. The particles of solid can only vibrate and rotate about fixed position but they cannot translate i.e move from one place to another. Solids have definite shapes, definite volumes and are difficult to compress. The strength of any solids is determined by the type of bonds or forces holding the particles together. Examples are ice, stone e.t.c.
Liquid state: the particles in a liquid move freely because they have more kinetic energy than solid and they are not longer held in a fixed position. The forces of attraction are weaker than those in solid particles. A liquid does not have a fixed shape but it takes the shape of its container. Examples are liquid water, kerosene e.t.c.
Gaseous state: the particles of gaseous state are free to move about in all directions at great speed. They constantly collide with themselves and the walls of their containers. Hence, they exhibit Brownian motion (haphazard movement). A gas has no definite shape. It occupies the whole volume of its container and it can be compressed. Examples are dust, smoke, fumes e.t.c.
PROPERTIES OF MATTER
The properties of matter can be characterized as physical or chemical properties.
When we say physical properties we refer to physical changes that can be detected by senses, such as color, odour, taste, boiling point, melting point, density, malleability, crystalline e.t.c.
While chemical properties refer to changes that will result in the formation of new substances such as the rusting of iron and the formation and decay of substances.
BOYLE’S LAW
Boyle’s law states that the pressure of a given mass of gas is inversely proportional to its volume, provided that the temperature remains constant.
p α1v Where P= Pressure
p = Kv V= Volume
Or PV = K. K = mathematical constant
As such P1V1=P2V2
This can be represented graphically as;
Examples;
200cm3 of a gas has a pressure of 510mm/Hg. What will be its volume if the pressure is increased to 780 mm/Hg, assuming there is no change in temperature?
A gas has a volume of 500cm3, when a pressure of 76mm/Hg is exerted on it. What will its volume be if the pressure on it is changed to 73mm/Hg, assuming the temperature remains constant?
CHARLES’ LAW
Charles’ law states that the volume of a given mass of gas is directly proportional to its temperature in Kelvin, provided that the pressure is kept constant.
V α T Where V = Volume
V = KT T = Temperature
Or VT = K K = mathematical constant
Hence, V1T1 = V2T2------------------------VnTn
It can be represented graphically as;
NB: -273OC is said to be the absolute zero temperature.
Examples:
A certain mass of gas occupies 300cm3 at 35oc. At what temperature will it have its volume reduced by half, assuming the pressure remains unchanged?
A fixed mass of gas occupies 40cm3 at 0oc. find the temperature at which its volume becomes 54.7cm3, assuming that its pressure remains constant.
Dalton’s Law of Partial Pressure
Dalton’s law of partial pressure states that, in a mixture of gases which do not react chemically together, the total pressure of a mixture of gases is equal to the sum of the partial pressure each gas would exert, if it alone occupies the whole volume of the container at the same temperature.
PTotal = PA + PB + PC + PD + ....................PN
Where the pressure each constituent gas exerts is called its partial pressure.
Graham’s law of diffusion of Gases
Graham’s law of diffusion states that at a constant temperature and pressure, the rate of diffusion of a gas is inversely proportional to the square root of its density d.
R α1/√d Where R = rate of diffusion
d = density of gas
R = K√d (where K is constant)
For the rate of diffusion of two gases;
R1R2=√d2d1
since the density of a gas, d, is proportional to its relative molecular mass, M. the relation can be written as:
R1R2=√M2M1
Rate is the reciprocal of time, i.e R = 1t
Therefore, substituting R1 = 1t1 and R2 = 1t2
R1R2= t2t1
Thus; t2t1= √m2m1
Example;
200cm3 of hydrogen diffused through a porous pot in 40 seconds. How long will it take 300cm3 of chlorine to diffuse through same pot?
GENERAL GAS EQUATION
Boyle’s and Charles ’laws are combined into a single expression known as the general gas equation and it states that the product of pressure and volume divided by the absolute temperature is constant for a fixed mass of gas, i.e
PVT = K
Therefore, P1V1T1= P2V2T2-------------PnVnTn .
This expression shows the relationship between the three variables, i.e, volume, temperature and pressure.
Examples:
The volume of a given mass of gas is 300cm3 at 27oc and 700mm/Hg. What will be its volume at 36oc and 750mm/Hg?
At S.T.P, at certain mass of gas occupies a volume of 800cm3. Find the temperature at which the gas occupies 940cm3 and has a pressure of 675mm/Hg.
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