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Structure of the Atom and the Periodic Table
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Structure of the Atom and the Periodic Table

Structure of the Atom and the Periodic Table
The Structure of the Atom
Dalton thought that atoms were solid, indivisible particles. But, as a result of work done mainly by Lord Rutherford, the idea has been greatly changed in recent years. According to Rutherford, the atom consists of 3 kinds of particles - protons, neutrons, and electrons. These are called sub-atomic particles. The center of the atom is called the nucleus. The nucleus contains a cluster of two sorts of particles, protons, and neutrons. The nucleus is very small, occupying only about 1% of the volume of an atom. The rest of the atom is mostly empty space, with electrons spread out in it.
The center of the atom is called the nucleus. The nucleus contains a cluster of two sorts of particles, protons, and neutrons. The nucleus is very small, occupying only about 1% of the volume of an atom. The rest of the atom is mostly empty space, with electrons spread out in it.
Electrons move around the nucleus in special paths called electron shells (orbits/or orbitals or energy levels). Protons and electrons have electric charges. Neutrons have no charges. All the particles in an atom are very light. Their masses are measured in atomic mass units rather than grams. The proton is a positively charged particle. Its mass is about equal to that of a hydrogen atom. The neutron has no charge, it is neutral. Its mass is about equal to that of a hydrogen atom. The electron is negatively charged. Its charge is equal but opposite to the charge on the proton. It has a very small mass, about 1⁄1840 of the mass of the proton.
The Properties of each Particle in an Atom
The properties of these particles are summarized in the table below:
A single atom is electrically neutral (it has no electrical charge). This means that in an atom there must be equal numbers of protons and electrons. In this way, the total positive charge on the protons is balanced by the total negative charge on the electrons orbiting the nucleus. So, the charges must cancel.
Names and symbols of the first twenty elements of the periodic table
Simple structure of the atom; protons, electrons and neutrons; electron energy levels in atoms
Atomic characteristics
Definitions of atomic number, mass number, isotopes and relative atomic mass(reference C-12); examples of isotopes
Relationship between Atomic Number and Number of Protons
All atoms of one element have the same number of protons. This is called the atomic number (or proton number) of that element. It is given by the symbol Z.No two elements can have the same atomic number. Sodium atoms have 11 protons. This is what makes them different from all other atoms. Only sodium atoms have 11 protons, and an atom with 11 protons must be a sodium atom.In the same way, an atom with 6 protons must be a carbon atom. Also, any atom with 7 protons must be a nitrogen atom. So, you identify an atom by the number of protons in it. There are 109elements altogether. Of these, hydrogen has the smallest atoms, with only 1 proton each. Helium atoms have 2 protons each. Lithium atoms have 3 protons each, and so on up to meitnerium atoms, which have 109 protons each. Table 5.3 shows the first20 elements arranged according to the number of protons they have.Every atom has an equal number of protons and electrons, so the atomic number also tells us the number of electrons in that atom. In any given atom of an element, the number of neutrons has no effect on the identity and properties of that particular element. It is the number of protons and electrons that determine the identity and properties of any given element. The number of neutrons only affects the mass, since each one of them has the same mass as that of a proton.
Mass Number of an Atom from Numbers of Protons and Neutrons
Protons alone do not makeup all the mass of an atom. The neutrons in the nucleus also contribute to the total mass. The mass of the electrons can be regarded as so small that it can be ignored. As a proton and a neutron have the same mass, the mass of a particular atom depends on the total number of protons and neutrons present. This is called the mass number (or nucleon number). The mass number of an atom is found by adding together the number of protons and neutrons. It is given by the symbol A. Table 5.3 shows the mass number of the first 20 elements, arranged in order of increasing atomic mass (mass number).If the mass number and atomic number for any given atom are known, then its sub-atomic composition can be worked out.The mass number = number of protons + neutrons in an atom. The sodium atom has 11 protons and 12 neutrons, so the mass number of sodium is 23. Since the atomic number is the number of protons only, then:Mass number – atomic number = number of neutrons. So, for sodium atom, the number of neutrons = (23-11) =12. You can also take into account the fact that, because the number of protons is always equal to the number of electrons, then the number of electrons in the sodium atom is simply 11. The same rule can be applied to work out the sub-atomic composition of any element.
These two relationships are useful:
  • Number of electrons = number of protons = atomic number
  • Number of neutrons = mass number (A) – atomic number (Z).
Calculation of relative atomic mass from relative abundance of isotopes of an element
The Concept of Isotope
Atoms of the same element may have different numbers of neutrons. In a normal situation, atoms of the same element will have the same number of neutrons. However, many cases occur in which two atoms of the same element contain the same number of protons but different numbers of neutrons. Having an equal number of protons, these atoms must also have equal numbers of electrons. However, the differing numbers of neutrons cause the atoms to have different mass numbers. An element showing such properties is said to show isotopy and the varieties of the atom are called isotopes of the element.Therefore, isotopy can be defined as the tendency of atoms of one element to possess the same atomic number but different mass numbers (atomic masses). Isotopes can be defined as atoms of the same element with the same number of protons but different numbers of neutrons, or as „atoms of the same element with the same atomic number but different atomic masses‟.The isotopes of an element have the same chemical properties because they contain the same number of electrons. It is the number of electrons in an atom that decides the way in which it forms bonds and reacts with other atoms. However, some physical properties of the isotopes are different. The masses of the atoms differ, and therefore other properties, such as density and rate of diffusion, also vary.Many isotopes (like tritium) are unstable. The extra neutrons in their nuclei cause them to be unstable so that nuclei break spontaneously (that is, without any extra energy being supplied), emitting certain types of radiation. They are known as radioisotopes.
Notation for isotopes
In order to distinguish between different isotopes of the same element in writing symbols and formulae, a simple system is adopted. The isotope of an element, say X will have the symbol X, AZ, where A is the mass number of the isotope and Z is the atomic number of an atom of X. Thus, for all isotopes of one element, Z is constant, and A varies because there are different numbers of neutrons in the different isotopes of the element. For example, the three isotopes of carbon are expressed as 12C6, 13C6, and 14C6. Chlorine has two isotopes: 35Cl17 and 37Cl17. Since A represents the total number of neutrons and protons in the nucleus of an atom (mass number/atomic mass), and because Z is the number of protons (atomic number), then the number of neutrons in the nucleus of a given isotope is given by: Number of neutrons in the nucleus = A – Z
Relative atomic masses
As we have seen, most elements exist naturally as isotopes. Therefore, the value we use for the atomic mass of an element is an average mass. This takes into account the proportions (abundance) of all the naturally occurring isotopes. If a particular isotope is present in a high proportion, it will make a large contribution to the average.
Example
A sample of chlorine gas contains 75% of the isotope 35Cl17 and 25% of the other isotope 37Cl17. What is the relative atomic mass of chlorine?
Solution
To work out this problem, simply multiply the mass number of each isotope with the abundance and sum up the products thus:
This average value for the masses of atoms of an element is known as the relative atomic mass (Ar). Therefore, the relative atomic mass of chlorine is 35.5 (i.e., Ar =35.5).
The periodic table
Consider the electronic configuration of the first twenty elements of the periodic table shown in the table below.
You will notice that elements in the same vertical columns (groups) have the same number of electrons in the outermost shells of their atoms. Because the outer electrons determine the chemical properties of an element, then the elements in each period tend to resemble each other closely in chemical behaviour. For instance, the noble gases, He, Ne, and Ar show a chemical inertness that is characterized by the stable outer electron octet or duplet. Due to this reason, the compounds of the noble gases with other elements have not been found.
Attempts to classify elements by arranging them in order of increasing atomic weights show that the properties of elements were periodic. This means elements with similar or comparable properties appear after a certain specific interval in a given arrangement. The occurrence of successive groups of elements showing strong chemical similarity in this way is called periodicity.
Therefore, periodicity is the repetition of similar chemical properties of elements after a certain specific interval in a given arrangement. The repetition in properties is due to repetition of the similar electronic configuration of outermost shells of elements after certain intervals.
The Change in Properties of Elements Across the Periods
Atomic and ionic size
The sizes of atoms and ions may be given in terms of atomic radius and ionic radius units respectively. The number of shells an atom or ion posses and the nuclear charge determines the size of an atom or ion. This is how the two properties vary along the period and down the group:
Atomic size
Along the period: Considering the normal elements only, the size of the atoms decreases from left to right across the period. This is because as atomic number increases across the period, the nuclear charge (due to increasing protons) increases, and electrons in shells are pulled closer to the nucleus.
Ionic size
  1. Positive ions (cations): Across the period; The ionic size does not change, i.e. remains the same, as you move across the period from either direction.
  2. Negative ions (anions): A negative ion is larger compared to the corresponding neutral atom because on forming an ion, one or more electrons are added to the atom. The added electron(s) is/are repelled by the electron(s) already present in the outermost shell, hence leading to an increase in the size of an atom, even though no new shell is formed. Down the group and along the period: Ionic size increases down the group, and along the period, i.e. from left to right.
Atomic radii (singular: radius)
During the period: In the period, atomic radii decrease from left to right with an increase in the atomic number.
Electronegativity
Electronegativity is the tendency of an atom to attract the shared pair of electrons towards itself in a molecule. The electronegativity values of elements in group 0 (inert gases) are zero.Along the period: Electronegativity increases while moving across the period from left to right in the periodic table.
Metallic character (or electropositivity)
Electropositivity is the tendency of an element to lose the valency electron(s) and donate the same to other elements (usually non-metallic elements). This process occurs during the formation of new substances e.g. molecules and compounds. Literally, such reactions occur between metals and non-metals whereby metals donate electrons and non-metals receive these electrons. So, metals are electropositive elements while non- metals are electronegative elements.
Along the period: Generally, metallic character decreases along the period from left to right.The gradation in metallic properties across the period is as follows: Metals → poor metals → metalloids → non-metals → noble gases
Chemical reactivity
Reactivity is the tendency of an element to lose or gain electrons in a chemical reaction.
Along the period: For metals, the reactivity decreases from left to right in a period while it increases for non-metals.
Ionization Energy or Ionization Potential (I.E or I.P)
This refers to the minimum amount of energy required to remove the most loosely bound electron from an isolated atom or ion in its gaseous state. The smaller the value of ionization energy, the easier it is to remove the electron from the atom.M(g) →M+(g) + e-
Along the period: It increases along the period from left to right with the increase in atomic number.
Electron affinity (Ea):
This is just opposite to I.E. It is defined as the amount of energy released when an extra electron is added to an isolated neutral atom in its gaseous state.
Along the period: The value increases along the period from left to right.
Density and melting point
Build up of the periodic table for the 1st twenty elements on the basis of energy levels(rows=periods, columns=groups)
Ion formation
Formation of simple ions and cations
Qualitative treatment of the ionisation energy and electron affinity
Writing electron arrangement of ions formed from atoms:lithium, Na,Fl,Mg,Al, Cl,S
Definition of valency and oxidation numbers
Valency is the capacity of an atom to combine with one or more atoms to form a molecule or compound. Valency also refers to the number of electrons that an atom can gain, lose or share in forming a chemical bond with another atom. The valency (or combining power) depends on the number of electrons in the outermost orbit (or valency shell) involved in the formation of a chemical bond. The number of electrons in the valency shell is never greater than 7. The outermost electronic configuration is responsible for the variability of the valency.
Some elements exhibit more than one valency, i.e., they have variable valencies. Examples of elements with variable valencies are iron (2 and 3), tin (2 and 4) and copper (1 and 2). The other elements with variable valencies are as shown in table 7.1.
Valency and Oxidation States
There is a strong correlation between valency and oxidation state. The oxidation state of an element equals its valency or charge carried by its ion when an element ionizes in solution. An example of this relation is iron (II) whose oxidation state (or oxidation number) is 2 and its valency is 2. The same case applies to iron (III). Other elements with variable valencies such as copper (I) and copper (II) have oxidation state equal to 1 and 2 respectively. The list continues. You will learn more about oxidation states later.
The valencies of the common transition elements should be memorized. Valencies of the normal elements may be deduced from the group number they occupy in the Periodic Table. The valencies of elements of group I to IV are equal to the group numbers they occupy in the periodic table. The valency of an element in group V to VIII is equal to eight minus the group number. For example, the valency of chlorine which is in group VII is 1, i.e. (8 -7) =1. The valency of oxygen in group VI is 2, i.e. (8-6) =2. Elements in group 0 (or VIII) have zero valency i.e. (8 - 8) = 0.
Table 7.1. Valencies of common metals and non-metals
Derive valency and oxydation number of an element from atoms(it's position in the periodic table)
The Concept of Valence
Valency is the capacity of an atom to combine with one or more atoms to form a molecule or compound. Valency also refers to the number of electrons that an atom can gain, lose, or share in forming a chemical bond with another atom. The valency (or combining power) depends on the number of electrons in the outermost orbit (or valency shell) involved in the formation of a chemical bond. The number of electrons in the valency shell is never greater than 7. The outermost electronic configuration is responsible for the variability of the valency.
Some elements exhibit more than one valency, i.e., they have variable valencies. Examples of elements with variable valencies are iron (2 and 3), tin (2 and 4), and copper (1 and 2). The other elements with variable valencies are as shown in table 7.1.
Valency and Oxidation States
There is a strong correlation between valency and oxidation state. The oxidation state of an element equals its valency or charge carried by its ion when an element ionizes in solution. An example of this relation is iron (II) whose oxidation state (or oxidation number) is 2 and its valency is 2. The same case applies to iron (III). Other elements with variable valencies such as copper (I) and copper (II) have an oxidation state equal to 1 and 2 respectively. The list continues. You will learn more about oxidation states later.
The valencies of the common transition elements should be memorized. Valencies of the normal elements may be deduced from the group number they occupy in the Periodic Table. The valencies of elements of group I to IV are equal to the group numbers they occupy in the periodic table. The valency of an element in group V to VIII is equal to eight minus the group number. For example, the valency of chlorine which is in group VII is 1, i.e. (8 -7) =1. The valency of oxygen in group VI is 2, i.e. (8-6) =2. Elements in group 0 (or VIII) have zero valency i.e. (8 - 8) = 0.
Simple Formulae of Binary Compounds
The Chemical formula is a method of representing the molecule of a compound by using chemical symbols. It is a way of expressing information about the atoms that constitute a particular chemical compound. The chemical formula identifies each constituent element by its chemical symbol and indicates the number of atoms of each element found in each single molecule of that compound.
The symbol for hydrogen atom is H. When two hydrogen atoms join together they form a molecule,H2. The number 2 to the right and below the symbol shows the number of atoms a molecule contains. P4 and S8 represents 4 atoms of phosphorus and 8 atoms of sulphur contained in one molecule of phosphorus and one molecule of sulphur respectively.
While the formula for chlorine molecule isCl2, it cannot be expressed as 2Cl. This is because 2Cl means two atoms of chlorine and not a molecule of chlorine.H 2 O stands for a molecule of water which consists of two hydrogen atoms and one oxygen atom.
H2SO4stands for a molecule of sulphuric acid containing 2atoms of hydrogen, 1 atom of sulphur and 4 atoms of oxygen.CaCO3 stands for a molecule of calcium carbonate containing 1atom of calcium, 1 atom of carbon and 3 atoms of oxygen.Where it deems necessary to show the number of molecules a compound contains, this is achieved by writing the appropriate number before the formula of the compound. A few examples are shown below:
  • 2H 2O means two molecules of water
  • 3H2 SO4 means three molecules of sulphuric acid
  • 5CaCO3 means five molecules of calcium carbonate
It is important to note that the figure appearing before the formula multiplies the whole of it. For example, 3H2SO4 stands for 3 molecules of sulphuric acid containing six atoms of hydrogen, three atoms of sulphur and twelve atoms of oxygen.It is a big mistake to think that the number 3 before the molecule multiplies only the symbol which immediately follows it (that is,H2). This is quite wrong. The 3 multiplies the whole of the formula
Formulae of Binary Compounds
A binary compound is a compound made of only two types of the reacting species, for example, sodium chloride (NaCl),which is made of only sodium and chlorine, is a binary compound. Look at table 7.1. The size of the charge on an ion isa measure of its valency or combining power. You will notice that, ignoring the signs for the charge on ions, the value of the charge on ion is equal to the valency of the atom. You will need to memorise the valencies of these metals as much as possible so as to be able to write the formulae of their compounds correctly.
The following are the rules for writing down the chemical formulae of chemical substances:
  1. Metals (or their positively charged ions) must start in the formula, followed by non-metals (or they're negatively charged ions).
  2. Where the formula is to include a radical, the radical must be treated as a single atom and must be bracketed if need be.
  3. The ammonium ion is to be treated as if it were a metal.
  4. Positive charges must be equal to negative charges for a neutral molecule or compound.
  5. Single elements; say Na, K, Si, Ag, etc. should not be bracketed.
  6. Valencies of metals (or positive ions) should be exchanged and written as subscripts.
  7. The valency of 1 is simply assumed and not written in the formula.
This is best shown by using some examples. The following procedure must be followed when writing the formulae of binary compounds:
  1. Write down correct symbols for atoms of elements or ions that make up the compound.
  2. Write down the valencies of the atoms of the elements.
  3. Interchange the valencies and write them as subscripts in the final formula of the compound. Remember that the valency of 1 is not expressed in the formula. At this final step, the radicals must be bracketed if necessary.
Study the following examples and make sure you understand how this works:
Nomenclature of Inorganic Compounds
The term "nomenclature" refers to the "system of naming". The system of naming in use is recommended by the IUPAC (International Union of Pure and Applied Chemistry). The modern system of naming reveals the type of elements present in a given compound. The old or trivial names have been dropped out.
Some common and important compounds have historical names that do not seem to fit in the system, for example water H O 2 ,ammonia 〈NH3〉, methane , CH4and mineral acids such as sulphuric (VI) acid〈H2SO4〉, nitric (V) acid〈HNO3〉and hydrochloric acid〈HCl〉. Also organic acids such as ethanoic acid (CH3COOH) are also included in this group. These names are trivial but they have been adopted in modern nomenclature.If these exceptions are omitted, there are basic generalizations that are useful:
  1. If there is a metal in the compound, it must be named first. In this case ammonium ion, NH4¹, is regarded as if it were a metal in the compounds it occurs such as NH4 NO3 , NH Cl 4 , etc.
  2. For elements with variable valencies such as iron and lead,Roman numerals are included in the name to indicate the valency of the metal or the ion which is present. For example,iron (III) chloride contains Fe3+ while iron (II) chloride contains Fe² . The same case applies to lead (II) and lead (IV)compounds and so on.
  3. Compounds containing two elements only (binary compounds) have names ending in …..ide; for example sodium chloride (NaCl), calcium bromide (CaBr2), magnesium nitrite( Mg3 N2) , etc. The important exception to this is hydroxides,which contains the hydroxide (OH) ion.
  4. Compounds containing a poly atomic ion (usually containing oxygen) have names that end with …ate; for example calcium carbonate (CaCO3) , potassium nitrate(KNO3) , magnesium sulphate(MgSO4) , sodium ethanoate (CH3 COONa) , etc.
  5. The names of some compounds use prefixes to tell you the number of that particular atom in the molecule. This is useful if two elements form more than one compound. For example:carbon monoxide (CO), carbon dioxide (CO2), nitrogen dioxide NO2 dinitrogen tetra oxide N2 O4 , sulphur dioxide SO2 sulphur trioxide SO3 , etc.
The following prefixes indicate the number of atoms in caseslike this: mono – one; di – two; tri – three; tetra – four; pent –five; hex – six; hept – seven; oct – eight; non – nine; and dec –ten.
Names and formulae of common radicals
Use of valencies in determining the chemical formulae of common compounds
Writing simple balanced equations
Chemical equations with state symbols
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