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Atoms : Periodic Table, Isotopes, Atomic Structure, Electronic Structure, Ions

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Atoms

The idea of an atom was first put forward by the Greek Philosophers in 400 BC. From their observations on the continuous chances that were taking place in the universe, the Greeks felt that there must be atoms that make up the universe. Unfortunately, this proposal was not backed by experimental data. As such, Scientists did not accept the proposal. The situation remained for nearly 2000 years when it was revived by another Scientist by the name of John Dalton. Atoms, are the smallest units into which matter can be divided without the release of electrically charged particles. It also is the smallest unit of matter that has the characteristic properties of a chemical element.

An element contains only one type of atom and cannot be broken down into anything simpler through chemical reactions. Everything around us is made up of atoms. Atoms are made up of three types of subatomic particles called protons, neutrons, and electrons. The central part of an atom contains most of the mass of the atom. The nucleus is made of protons and neutrons. Electrons surround the nucleus and are arranged in orbits or rings.

Protons have a positive charge and a relative mass of 1.Nuetrons have no charge and a relative mass of 1.Electron have a negative charge and a very small relative mass of 0.000543.

In an atom, the number of protons and the number of electrons are the same. For this reason atoms have no overall electrical charge. The number of protons in an atom is known as the proton number (or atomic number). Sodium contains 11 protons, 12 neutrons and 11 electrons.The number of protons plus the number of neutrons is known as the nucleon number (or mass number).

All atoms of an element have the same proton number (the same number of protons and electrons). The chemical properties and chemical reactions of an element depend on the number of electrons present. Therefore all atoms of an element react in the same way because they have the same number of electrons.

The Periodic Table

Periodic Table

Periodic Table

The Periodic Table groups elements with similar chemical and physical properties together. Initially, the Russian chemist, Dmitri Mendeleev grouped elements in order of increasing atomic weights but this arrangement was later changed to one which grouped the elements in order of increasing atomic numbers. It was observed that if elements were arranged in order of increasing atomic numbers, elements with similar chemical properties recurred periodically and at regular intervals. This so called periodic law was used in the construction of the periodic table whose widely used version is shown on the inside of the back cover. In that table, there are eighteen (18) vertical columns which are called groups and about seven (7) rows, called periods. All the members of a group have the same outer shell electron configuration, apart from the ''n'' value which is different.


Some Representative Elements of Importance

Group 1 Elements: Li, Na, K. These elements have one outer shell electron. Group 1 elements are called alkali elements. These elements are so reactive that they are rarely found in elemental. They are soft, malleable and have a silver luster and being metals, they are good conductors of heat and electricity. They have low melting points which decrease down the group.

Alkali elements all have similar chemical properties. They are strong reducing agents as shown by their violent reaction with water, exemplified by sodium:

Na + H2O →Na+(aq) + OH-(aq) + H2(g)

The hydroxides produced are high melting solids which are soluble in water. The great reactivity of these elements is reflected in their atomic structure and their large atomic radius. The atoms lose the outer electron to acquire the electronic configuration of the nearest inert gas. Since their radii increase down the group.

Group 2 Elements: Be, Mg, Ca. These elements have two outer shell electrons. Group 2 elements are called alkaline earth elements. These elements are also rarely found in the elemental state. They are hard, dense, with high melting and boiling points and being metals, are good conductors of heat and electricity.

Compared with group 1a elements, group 2 elements have larger nuclear charges hence; their atomic sizes are smaller than those of group 1 elements.

Alkaline earths react with water only at high temperatures (compare, group 1) to form hydroxide:

Mg + 2H2O → Mg(OH)2 + H2(g)

These elements are, like group 1 elements, very reactive. Their reactivities also increase down the group for the same reason as the alkali elements. Unfortunately, elements like magnesium, readily react with oxygen in the air and form a thin layer of metal oxide on the surface of the metal. This greatly reduces their reactivity. Magnesium is therefore, much less reactive with say, cold water than calcium is, unless the water is heated up. At high temperatures, the elements in their metallic form can reduce other compounds as in the following example:

Na2O +Mg →MgO + 2Na

The resulting oxides are strong bases which react with acids and to some extent, water. In the latter solvent, the oxides form hydroxides (e.g., magnesium hydroxide, also known as milk of magnesia Mg(OH)2).

Group 4 Elements: C, Si, Ge, Sn. Group 4 elements have four outer shell electrons . Carbon is a non-metal, silicon and germanium are metalloid (such elements with properties between non-metals and metals are also termed semi-metals) and tin and lead are metals. Large amounts of carbon are found in living systems (plants and animals).

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There are two forms (different forms of the same element are also called allotropes) of carbon: graphite, which is widely used for making pencil leads, is an extended multilayered structure. In this structure, carbon atoms form six-membered rings which in turn, are joined to each other to firm a planar structure. Several of these layers are then stuck together by Van der Waals forces . Graphite conducts electricity in plane parallel to the plane but not perpendicular to it. The other allotrope of carbon is diamond in which each carbon is covalently bonded to four other carbons to form a three dimensional network which is hard. Carbon is somehow unique in group 4 in that it is able to form compounds in which several carbon atoms are bonded to each other to form chains or rings. This property of carbon is called catenation and is responsible for the variety of organic compounds. Silicon occurs in silica (SiO2) and silicate minerals. Silicon is widely used in the manufacture of computer chips and transistors. Lead occurs mostly in the metallic form. It is used for making electrodes in car batteries and also as an oil additive in the form of tetraethyllead (CH3CH2)Pb. Group 4 elements can form are able to form four covalent bonds in their compounds. Examples include the hydrides (e.g., CH4, SiH4, PbH4), chlorides (e.g., CCl4, SnCl4) and oxides (e.g., SnO2, PbO2, GeO2).

Group 5 Elements: N, P, As. Elements in this group have five outer shell electrons. phosphorus are non-metals, arsenic and antimony are metalloids while bismuth is a metal. Nitrogen occurs in abundance in the atmosphere as a diatomic molecule, N2. Phosphorus exists in rock phosphates (Ca3(PO4)2). Arsenic and antimony occur mostly as sulphite minerals while bismuth exists in pure form as a metal.There are three allotropes of phosphorus: the reactive white phosphorus has a tetrahedral structure (P4), red phosphorus exists as a polymer and the third form, black phosphorus is a layer structure much like graphite. Arsenic and antimony also have tetrahedral structure As4 and Sb4 respectively.Group 5a elements form various compounds including hydrides (e.g., NH3, PH3, AsH3, SbH3, BiH3), oxides (e.g., NO, NO2, P2O5, As4O6, Bi2O3) and halides (e.g., NF3, PCl3). The hydrides are basic compounds and this basicity decreases as we go down the group.Group 7 Elements: F, Cl, Br, I. These elements have seven outer shell electrons. Group 7a elements are called the halogens and unlike groups 1 and 2 elements which are metals, these are non-metals. Fluorine is the most electronegative of the elements and usually has an oxidation number of -1.


The other elements are also electronegative but unlike fluorine, they exhibit both negative as well as positive oxidation numbers. Halogens are very reactive and they act as oxidants in most of their reactions. Their strength as oxidants decreases as the atomic number increases. Unlike group 1 and 2 elements, halogens gain electrons when they react. This means that the smaller the volume of the element, the easier it will grab the electrons. Consequently, reactivity in the halogens increases as we go up the group, with fluorine being more the most reactive. The reaction of potassium with fluorine is more vigorous than that of potassium with chlorine. Halogens are colored, poisonous molecules with unpleasant smell; fluorine is a yellow gas, Chroline is green/yellow gas, bromine is red and iodine is a purple solid.Group 7a elements react with metals to form metal halides. Large metals, with low oxidation numbers (e.g., +1 or +2), form ionic halides whereas small metals with high oxidation numbers result in semi metallic halides. Halogens also react with each other as shown by the oxidation of bromide (Br-) by chlorine- a reaction that is used to recover bromine from sea water.

Cl2 + 2Br- →Br2 + 2Cl–

Group 8 Elements: He, Ne, Ar, Kr. Group 8 elements have eight outer shell electrons. Group 8 elements are called rare, inert or noble gases. Since they have a complete octet, their structures are very stable, hence the name inert. The word inert is however, a misnomer since compounds of some of these elements have been detected such as xenon difluoride (XeF2) and xenon oxyfluoride (XeOF2), among others. It is perhaps important to note that it is only this group of elements that has atoms which exist in nature as single-atomed molecules. This property is a reflection of the stable electronic configurations of these elements.


Isotopes

Atoms of a particular element all have the same atomic number and therefore, the same number of protons. However, most of the elements do not have all their atoms with the same number of neutrons. Such atoms belonging to the same element but having different numbers of neutrons in their nuclei are called isotopes. It is clear therefore that isotopes have different masses .

Some atoms of certain elements can have a different nucleon number because they have a different number of neutrons, but the same proton number as they have the same number of protons and electrons. Atoms of the same element that have different numbers of neutrons are known as isotopes for example, carbon has two isotopes.

The proportion of different isotopes in an element does not change the chemical properties of the element. However, the proportion of different isotopes in an element can affect the physical properties of an element such as the melting point, boiling point, and density. Some isotopes are radioactive. This means that the nucleus is unstable and gives off radiation. Uranium is an example of a radiactive isotope. Radioactive isotopes can be used medically to treat cancer, find tumour, and follow the flow of fluids around the body. In industry, radioactive isotopes can be used to locate leaks in pipelines, measure the flow of substances through pipelines, and measure the thickness of meta.


Relative Atomic Mass (RAM)

We have stated that an atom has some components. The electrons are negatively charged and each has a charge of -1.60206 × 10^-19 coulombs and a mass of 9.109 × 10^-28g. The electron charge is abbreviated 'e' (-1.60206 ×10^-19 = e) so that one electron has a charge of -e or just -1. It should be noted here that the minus (-) sign does not mean that the electron is short of electric charge but that its electrical charge is of the type called negative. The protons are positively charged and each has a charge of +1.60206 × 10^-19 coulombs and a mass of 1.672 × 10^-24g. As for the electron, the charge on a proton is simplified to +1. All the atoms except hydrogen also have neutrons which are neutral (have a charge of zero) but have a mass of 1.675 × 10^-24g. The other components of an atom include neutrinos, positrons, μ-mesons and pi-mesons among others but the detailed nature of these is beyond the scope of this topic

The absolute masses of atoms cannot be determined as the atoms are very light. Only the relative masses can be obtained. To get these relative masses, we need to assign a weight to one atom (the standard) and relate the masses of the others to it. What we get then is the Relative Atomic Mass (RAM) of the atom. Initially, hydrogen which had been assigned a mass of one, was chosen as the standard. However, for some practical reasons a carbon atom with six protons and six neutrons (carbon-12) was later chosen as the new standard. The carbon-12 has a mass of exactly 12. The atomic masses of all the atoms were obtained with reference to a carbon-12 as the standard. However, with technological revolution in modern times, mass spectrometers are used to determine the relative masses of atoms directly. Note that the relative atomic masses are rations hence have no units.

Since there are more than one atom for a particular element (isotopes, Section 1.5) with differing relative masses we need to calculate a weighted mean of their masses taking into account their relative abundances. This weighted mean for the mass of a particular element is called the atomic weight of that element. Therefore, we define the atomic weight of an element as the weighted mean of the relative masses of the isotopes of that particular element. For example, chlorine has two isotopes with the relative masses and abundances of 35 (75.5%) and 37 (24.5%).


Relative Atomic Weight Solved Examples

1. Nitrogen is made up of two isotopes, 14–N and 15–N. Given the nitrogen’s atomic weight of 14.007, what is the percent abundance of each isotope?

Solution

Let assign abundance of 14–N “x” and “1 − x” for 15–N this is where trick is. Bear in mind the abundance always add up to 1.

Now let’s solve

14.00(x) +15(1 – x) =14.007

14x +15 – 15x =14.007

−x=14.007 −15 giving x =0.993 then 1− x is now same as 1−0.993 giving 0.007. Now let’s multiply by 100% in each 14–N= 0.993 x 100% giving 99.3% abundance and for 15–N = 0.007 x 100% giving 0.7% abundance

2. Rubidium (Rb) has two isotopes, one is 84.912–Rb and the other one is 86..909–Rb. The atomic weight of Rb is 85.47. Calculate the relative abundances of the two isotopes?

Solution

Let assign abundance of 85–Rb “x” and “1 − x” for 87–Rb.

84.912(x) +86.909(1 – x) =85.47

84.912x +86.909 – 86.909x =85.47

−1.997x =85.47 – 86.909

−1.997x = −1.439

X = 0.720580871 multiplying by 100% we get 72.05808713% relative abundance of 84.912–Rb

Relative abundance of 84.912–Rb + relative abundance of 86.909–Rb = 1

Relative abundance of 86.909–Rb = 1 − 0.720580871

Relative abundance of 86.909–Rb = 0.279419128 multiplying by 100% we get 27.94191287%


Mass Spectrometry

Mass spectrometers can be used to determine all the isotopes present in a sample of an element and their percentage or relative abundance.

The relative atomic mass is a weighted average of all the isotopes of an element

The data collected from a mass spectrum can help to calculate a relative atomic mass.

Chlorine has 2 isotopes: Cl-35 (75%) and Cl-37 (25%) so the RAM is calculated to be 35.5


Uses For Mass Spectrometer

1. Drug testing

Drug testing in sports to identify chemicals in the blood

2. Quality control

Quality control in the pharmaceutical industry

3. Age determinants

Radioactive Carbon-13 Dating to determine ages of fossils or human remains

4. Rocks testing

Testing rocks on different planets


Structure of The Atom

In Science, people are always curious. Having accepted that an atom exists, the Scientists then wanted to know how this atom looks like. As such, they carried out several studies.From these studies, it was found that the atom has even smaller particles which are called subatomic particles. For example it was found that an atom consists of a positively charged nucleus surrounded by light (not heavy) negatively charged particles called which were named electrons. In 1909, Ernest Rutherford studied the deflection patterns of alpha particles (He2+ ions) by thin metal foils and an analysis of the results led him to the conclusion that an atom has a tiny central nucleus that carries practically all the mass of the atom and that the lighter particles (the electrons) are around this nucleus. The number of positive charges on the nucleus is called theatomic number (abbrev. Z). however, since the atom was found to be neutral, it was clear that the number of positively charged particles must equal the number of negatively charged particles (electrons). Therefore, the number of electrons surrounding the nucleus in a neutral atom, must also be equal to the atomic number. The positively charged particles are called protons. The hydrogen atom for example, has one proton in its nucleus and carries one electron. Furthermore, it was discovered that the electrons are arranged around the nucleus in ''shells''. Imagine this on a magnified scale. You put the first group of electrons on a small ''circle'' with the nucleus inside it. Then put the second group of electrons on a bigger circle with the smaller one inside it and so on.


Electronic Structure

At this point, you need to know how the electrons occupy the various orbitals in an atom. Electrons are added to each atom, one at a time across a row. The lowest energy shell fills first then the next and so on. Within a shell, The lowest energy orbitals fill first and once those are full, the next lower energy ones fill and so on. Structures which show the n-values and their orbitals that are occupied, together with the number of electrons in them are called electronic structures. The question we can ask at this point is: ''how many electrons can a single orbital accommodate?''. The answer to this question was arrived at by Wolfgang Pauli, who introduced a restriction on the number of electrons which can go into a single orbital. The restriction called the Pauli Exclusion Principle states that ''any orbital will not hold more than two electrons''. This means that any orbital can hold 0, 1 or 2 electrons but not more than 2. Therefore, the maximum number of electron allowed for each n value will be equal to 2n²


Electron Configuration In Atoms

The way electrons are arranged around the nucleus is known as electron configuration. Electrons are arranged around the nucleus in energy levels or shells. The first energy level or shell can only contain 2 electrons. The shells after this contain 8 electrons, as explained in the below. For example, the electron configuration for the element oxygen is written as 2,6 as oxygen has 8 electrons in total: 2 electrons in the first shell and 6 in the second shell. The electron configuration for potassium is 2,8,8,1 as potassium has 19 electrons in total: 2 electrons in the first shell, 8 in the second and third shell, and one in the last shell.

Electron Configuration

Electron Configuration

Group numbers are always written in Roman numeral.The number of electrons found in the outer shell is known as the number of valency electrons. The valency number refers to how many electrons are available for an element to bind to another element. All elements with the same number of valency electrons are found in the same group in the periodic table. Group numbers are always written in Roman numerals.

Elements with a full outer electron shell are the most stable elements. All elements in this group are gases and are known as the noble gases.


Compounds

A compound consists of atoms of two or more different elements chemically bonded together. For example, water is a compound made from the elements hydrogen and oxygen.The properties of the compound formed are different to the elements that make it up. For example,the properties of water are different to the properties of both hydrogen and oxygen. In order to separate a compound into the elements it contains, a chemical reaction must take place


Mixture

A mixture contains two or more substances mixed together but not chemically bonded together. The properties of the mixture are the same as the substances that make it up. As the parts of a mixture are not chemically bonded together they can be separated easily using physical methods such as distillation.


Periodicity

Periodicity is the trend of chemical and physical properties across a period or down a group of elements.

Atomic Radius decreases across a period as there are more protons so the nuclear charge is greater and pulls the electrons in more.

Atomic Radius increases down a group as more shells and increased shielding weaken the nuclear attraction on the valence electrons.

First Ionisation Energy tends to increase across a period due to an increase in protons providing a stronger nuclear attractive force on the electrons.

First Ionisation Energy tends to decrease down a group because the increased number of shells results in shielding weakening the nuclear attractive force.

Melting Points are based on bonding and intermolecular forces within the elements.

For Na, Mg and Al – These elements have metallic bonding so the increase in outer electrons and a smaller ion results in a larger melting point.

For Si – Silicon is a macromolecular substance so requires a lot of energy to break the multiple covalent bonds in the structure.

For Cl2, S8, P4 and Ar – These elements exist as simple molecular substances that have weak Van der Waals Forces between them so less energy is required to break these Intermolecular Forces. The more electrons they have in a molecule, the higher the m.p.


Isotopes and Ions

An Isotope is an element whose atomic number is the same but the mass number(number of neutrons) is different.

Isotopes only have different numbers of neutrons. The number of protons and electrons is the same

Ions are charged atoms which have a different number of electrons. Positive Ions have lost electrons while negative Ions have gained electrons

Protons determine the element. Changing the number of protons means a new element is formed


Ionization Energies

The first ionisation energy is the energy required when one mole of gaseous atoms forms one mole of gaseous ions with a 1+ charge.

X(g) –> X+(g) + e-

The second ionisation energy is the energy required when one mole of gaseous ions with a 1+ charge forms one mole of gaseous ions with a 2+ charge.

Ionization energies are affected by

  1. The attraction of the nucleus (the more protons, the greater the attraction).
  2. The distance of the valence electrons from the nucleus (the further from the nucleus, the weaker the attraction).
  3. The shielding provided by the electrons (the more electrons infront of the valence electrons, the weaker the attraction).

Helium has the highest first ionisation energy because its nuclear attraction is strongest as the atomic radius is shorter than other atoms and there is no shielding.

Successive ionisation energies are always larger because there is a greater positive:negative charge ratio as there are fewer electrons and the ionic radius is smaller


Ionization Energies and Electronic Structures

The more electrons removed from an atom, the higher the ionisation energy as there are more protons to electrons so the nuclear attraction is greater for each subsequent electron removed.

A larger jump in Successive Ionisation energy can be seen when an electron is removed from a shell closer to the nucleus as there is less shielding and the distance is smaller

Nobel Gases have higher first ionisation energies as they have the most protons with the least amounts of shielding.

A large drop in first ionisation energies occurs between the Nobel Gases and Group 1 metals as shielding increases weakening the nuclear attraction and causing electrons to be more easily lost.

Small drops in first ionisation energies can be seen between Group 2 and 3elements (e.g. Be and B) as the electrons are removed from a different subshell.

Small drops in first ionisation energies can be seen between Group 5 and 6elements (e.g. N and O) as the electrons start to pair up in a subshell

Graph of First Ionization Energies

Graph of First Ionization Energies

Ionic Bonding

Ionic bonding is the bond of attraction between METALS and NON- METALS.

In ionic bonding atoms either lose or gain electrons – this causing them to become charged (ions).

When they become charged they then are strongly attracted to each other (because of their opposite charges) – however the overall charge of any ionic compound is zero (because the negative and positive charge cancel each other out). Group 1 & 2 and 6 & 7 are most likely to form ions

Metals (not counting transition metals) are keen to get rid of their electrons – this is so they can have a full shell left. Non – metals, however, are keen to gain electrons to have a full outer shell and become stable.

When they either gain (non –metals) or lose (metals) electrons then they become ions, as they either have more protons because they lost electrons (metals) so are positively charged, or they have more electrons because they gained them (non-metals) and are negatively charged


Metallic Bond Strength

  1. Increases across a period as more electrons become delocalize.
  2. Decreases down a group as the atomic radius increases.
  3. Magnesium is an example of a metallic crystal with the above properties. Each magnesium atom loses their 2 outer electrons to become Mg2+ ions.



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