9 - Periodic Table

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Jan 9, 20227 min read

9.1 Development Of The Periodic Table

-The Periodic Table was devised in 1869 by the Russian Dmitri Mendeleev, who was the Professor of Chemistry at St Petersburg University

-His Periodic Table was based on the chemical and physical properties of the 63 elements that had been discovered at that time

-Mendeleev arranged the 63 known elements in order of increasing atomic weights but in such a way that elements with similar properties were in the same vertical column

-He called the vertical columns groups and the horizontal rows periods and these are numbered 1-7 going down the periodic table

-Between Groups II and III is the block of elements known as the transition elements

-These elements with similar chemical properties are found in the same columns or groups

-There are 8 groups of elements

Group I : The alkali metals

Group II : The alkaline earth metals

Group VII : The halogens

Group 0 : Inert gases or noble gases

- The periodic table can be divided into 2 as shown by the bold line that starts beneath boron

-The elements to the left of this line are metals (fewer than 3-quarters) and those on the right are non-metals (fewer than one-quarter)

-The elements which are on the right are non-metals

-The elements which lie on this dividing line are known as metalloids

-The metalloids behave in some ways as metals and in others as non-metals

-If you look at the properties of the elements across a period of the Periodic Table, you will notice certain trends

-For example, there is:

- A gradual change from metal to non-metal

-An increase in the number of electrons in the a change in the structure of the element, from giant metallic in the case of metals e.g. magnesium, through giant to simple molecular

9.2 Electronic Configuration and the Periodic Table

-The number of valence electrons corresponds with the group number of the element

-The number of occupied shells corresponds with the period number

9.3 Group I - The alkali metals

-Group I consists of the five metals lithium, sodium, potassium, rubidium and caesium, and the radioactive element francium

-Lithium, sodium and potassium are all very reactive metals and they are stored under oil to prevent them coming to contact with water or air

-These 3 metals have the following properties:

They are good conductors of electricity and heat

-They are soft metals. Lithium is the hardest and potassium the softest

-They are metals with low densities. For example, lithium has a density of 0.53 g/cm^3 and potassium has a density of 0.86 g/cm^3

-They have shiny surfaces when freshly cut with a knife

-They have low melting points. For example, lithium has a melting point of 181°C and potassium has a melting point of 64°C

-They burn in oxygen or air, with characteristic flame colours, to form white solid oxides. For example, lithium reacts with oxygen in air to form white lithium oxide, according to the following equation:

- Lithium + Oxygen -> Lithium oxide

- Li(s) + O2(g) -> 2Li2O(s)

-These Group I oxides all react with water to form alkaline solutions of the metal hydroxide

Lithium oxide + water -> Lithium hydroxide

Li20(s) + H2O(s) -> 2LiOH(aq)

-They react vigorously with water to give an alkaline solution of the metal hydroxide as well as producing hydrogen gas. For example:

- potassium + water -> potassium hydroxide + hydrogen gas

2K(s) + 2H2O(l) -> 2KOH(aq) + H2(g)

-Of the first 3 metals in Group I, potassium is the most reactive towards water followed by sodium and then lithium

-Such gradual changes we call trends. Trends are useful to as they allow predictions to be made about elements we have not observed in action

-They react vigorously with halogens, such as chlorine, to form metal halides, for example, sodium chloride

- Sodium + Chlorine -> Sodium Chloride

2Na (s) + Cl2(g) -> 2NaCl(s)

-Considering the group as a whole, the further down the group you go, the more reactive the metals become.

-Francium is, therefore, the most reactive Group I metal

Why do you think potassium is more reactive than lithium or sodium?

-Potassium is more reactive because less energy is required to remove the outer electron from its atom than for lithium or sodium

-This is because as you go down the group, the size of the atom increases and the outer electron gets further away from the nucleus and because easier to remove

Group II - The alkaline earth metals

-Group II consists of the 5 metals beryllium, magnesium, calcium, strontium and barium and the radioactive element radium

-These metals have the following properties :-

-They are harder than those in Group I

-They are silvery-grey in colour when pure and clean. They tarnish quickly, however, when left in air due to the formation of metal oxide on their surfaces

-They are good conductors of heat and electricity

-They burn in oxygen or air with characteristic flame colours to form solid white oxides For example:

Magnesium + Oxygen -> Magnesium oxide

2Mg(S) + 02(g) -> 2Mg0(s)

-They react with water, but they do so vigorously than the elements in Group I. For example:

calcium + water -> calcium hydroxide + hydrogen gas

Ca(s) 2H2O(l) -> Ca(OH)2 (aq) + H2(g)

-Considering the group as a whole, the further down the group you go, the more reactive the elements become

9.4 Group VII - the Halogens

Group VII consists of the four elements flourine, chlorine, bromine and iodine, and the radioactive element astatine

-These elements are coloured and become darker going down the group

-They exist as diatomic molecules, for example, Cl2, Br2 and I2

-At room temperature and pressure, they show a gradual change from a gas (Cl2), through a liquid (Br2), to a solid (I2) as the density increases

-They form molecular compounds with other non-metallic elements, for example, HCL

-They react with hydrogen to produce the hydrogen halides, which dissolve in water to form acidic solutions

Hydrogen + Chlorine -> Hydrogen chloride

H2(g) + Cl2 (G) -> 2HCl (g)

H2(g) + Cl2 (g) -> 2HCl(g)

Hydrogen chloride + water -> hydrochloric acid

HCl (g) + H2O -> HCl (aq) -> H+(aq) + Cl- (aq)

-They react with metals to produce ionic metal halides, for example, chlorine and iron produce iron (III) chloride

iron + chlorine -> iron (III) chloride

2Fe(S) + 3Cl2(g) -> 2FeCl3c(s)

Displacement reactions

-If chlorine is bubbled into a solution of potassium iodide, the less reactive halogen, iodine, is displaced by the more reactive halogen, chlorine

- Potassium iodide + chlorine -> potassium chloride+ iodine

2Kl(aq) + Cl2(g) -> 2KCl (aq) + I2 (aq)

-The observed order of reactivity of the halogens, confirmed by similar displacement reactions, is:

The electronic configuration for chlorine and bromine :-

-In each case, the outer electron shell contains seven electrons. When these elements react, they gain one electron per atom to gain the stable electronic configuration of a noble gas

Chlorine is more reactive than bromine because the incoming electron is gained more easily by the smaller chlorine atom than in the larger bromine atom

-It is gained more easily because there is a stronger attraction between the negative charge of the incoming electron and the positive charge of the nucleus

In the larger bromine atom, there are more occupied electron shells surrounding the nucleus which lessen the attraction of the nucleus, and the electrons in these shells repel the incoming electron

-This makes it harder for the bromine atom to gain the extra electron it needs to gain a stable electronic configuration

-This is the reason the reactivity of the halogens decreases going down the group

-In the reaction of chlorine with potassium iodide, both Cl atoms in Cl2 gain an electron from an iodide ion, I-, thus forming 2 chloride ions, Cl-

-The iodide atoms formed by the loss of an electron combine to give an iodine molecule, I2

-The iodide ion has been oxidised because it has lost electrons

-The oxidation number has increased

-Chlorine has been reduced because it has gained electrons

-The oxidation number has decreased

Going further

Fluorine is used in the form of fluorides in drinking water and drinking water and toothpaste because it reduces tooth decay by hardening the enamel on teeth

-Chlorine is used to make PVC plastic as well as household bleaches. It is also used to kill bacteria and viruses in drinking water

-Bromine is used to make disinfectants, medicines and fire retardants

-Iodine is used in medicines and disinfectants, ad also as a photographic chemical

Practical Skills

Halogen displacement reactions

Safety:

- Eye protection must be worn

Experiment: Displacement reactions can be used to determine the reactivity of the Group VII elements, the halogens. In this experiment, a student uses these reactions to determine an order of reactivity for iodine, chlorine and bromine, and predicts the reactivity of the other two halogens, fluorine and astatine

If a more reactive halogen reacts with a compound of a less reactive halogen, the less reactive halogen will be displaced and it will form the halogen molecule, while the more reactive halogen becomes a halide ion

To observe the presence of the different halogens in solution, the student used an organic solvent as the halogens produce more vivid colours in this solvent compared to with water

To show the colour of the halogens in the organic solvent solutions, the student separately placed 3 halogens in test tubes and added a small amount of the organic solvent

The tubes were then fitted with a rubber bung and shaken

The colours of the halogens in water can be used to identify if a reaction has occurred when a solution of a halogen, for example, chlorine water, is mixed with a solution of a sodium halide, for example, sodium iodide solution.

9.5 Group VIII - The Noble Gases

-Helium, neon, argon, krypton and the radioactive element radon make up the most unusual group of non-metals, called the noble gases

-They were all discovered after Mendeleev had published his Periodic Table

-They were discovered between 1864 and 1900, mainly through the work of the British scientists Sir William Ramsay and Lord John William Strutt Rayleigh

-They are colourless gases

-They are monatomic gases- they exist as individual atoms, for example, He, Ne and Ar

-They are very unreactive

-No compounds of helium, neon or argon have ever been found

-However, more recently a number of compounds of xenon and krypton with fluorine and oxygen have been produced, for example, XeF4

-These gases are chemically unreactive because they have electronic configurations which are stable and very difficult to change

-They are so stable that other elements that other elements attempt to attain these electronic configurations during chemical reactions

9.6 Transition elements

-They are less reactive metals

-They form a range of brightly coloured compounds

-They are harder and stronger than the metals in Groups I and II

-They have much higher densities than the metals in Groups I and II

-They have high melting points (except for mercury, which is a liquid at room temperature)