metals and non-metals

definition of metals

metals are elements which form positive ions by losing electrons. we can say that they are electropositive in nature

definition of non-metals

non-metals are elements which form negative ions by gaining electrons. hence they are electronegative in nature

abundance of metals and non-metals

metals are more abundant than non-metals. however, the crust is mostly made up of non-metals (oxygen is 50% of the crust, aluminium is 7% of the crust)

physical properties of metals

they are malleable and ductile; that is, they can be hammered into thin sheets or drawn into thin wires (especially Au, Ag)

they are good conductors of heat (especially Ag, Cu, Al; except Pb)

they are good conductors of electricity (especially Ag, Cu, Al)

they are lustrous/ shiny

they are generally hard, strong, with high density (except K, Na)

they are solid at room temperature (except Hg, which is a liquid)

they have high MP and BP (except K, Na, gallium, cesium)

they are sonorous

physical properties of non-metals

they are neither malleable nor ductile

they do not conduct heat and electricity (except carbon in the form of graphite)

they are not lustrous (except iodine)

they are soft (except diamond)

they may be solid, liquid or gas

they are not strong, have low densities and are brittle (except alkali metals)

they have low MP and BP

reaction of metals with oxygen

metal + oxygen → metal oxide

• room temp:

  • Na + O2 → Na2O

  • K + O2 → K2O

  • Ca + O2 → CaO

• on heating:

  • Mg + O2 → MgO

  • Al + O2 → Al2O3

note: metal oxides are basic in nature. some of them dissolve in water to form metal hydroxide

amphoteric oxides

they are metal oxides which exhibit both basic and acidic properties. examples: ZnO and Al2O3

Al2O3 + NaOH → NaAlO2 + H2O

ZnO + NaOH → Na2ZnO2 + H2O

reaction of metals with water

metal + water → metal hydroxide + hydrogen

metal + steam → metal oxide + hydrogen

• cold water

  • K + H2O → KOH + H2 + heat

  • Na + H2O → NaOH + H2 + heat

  • Ca + H2O → Ca(OH)2 + H2

• hot water

  • Mg + H2O → Mg(OH)2 + H2

• steam

  • Mg + H2O → MgO + H2

  • Al, Zn

  • Fe + H2O → Fe3O4 + H2

reason why metals displace hydrogen from water

water slightly ionises to give H+ and OH- ions. a reactive metals makes hydrogen accept electrons in order to form H2 gas, and then forms an ionic bond with the OH- ions

why potassium and sodium are stored under kerosene oil

its to prevent their reaction with oxygen in air, as they are very reactive

reaction of metals with acids

metal (above hydrogen) + acid → salt + water

why nitric acid does not evolve hydrogen gas after reacting with every metal except Mg and Mn

HNO3 is a strong oxidising agent, and oxidises the hydrogen formed into water. an exception is when very dilute HNO3 reacts with Mg/Mn

Mg/Mn + HNO3 → Mg/Mn(NO3)2 + H2

aqua-regia

it is a mixture of 1 part concentrated HNO3 and 3 parts concentrated HCl (1:3). it is capable of dissolving all metals, even gold and platinum

why some metals are more reactive than others

a metal that loses electrons easily to form positive ions is more reactive than a metal that loses electrons less readily to form positive ions

reactivity series table

K Na Ca Mg Al Zn Fe Sn Pb [H] Cu Hg Ag Au

reactivity series table

K Na Ca Mg Al Zn Fe Sn Pb [H] Cu Hg Ag Au

reaction of metals with chlorine

metal + chlorine → ionic chloride

Na + Cl2 → NaCl

Fe + Cl2 → FeCl3

Cu + Cl2 → CuCl2

reaction of metals with hydrogen

K/Na/Ca/Mg + hydrogen → metal hydride

this occurs because the very reactive metals force hydrogen atoms to accept electrons and form H- ions

K + H2 → KH

Mg + H2 → MgH2

reaction of non-metals with oxygen

non-metal + oxygen → non-metal oxide

C + O2 → CO2 [CO2 dissolved in water gives H2CO3]

S + O2 → SO2 [SO2 dissolved in water gives H2SO3]

P4 + O2 → P2O5

non-metal oxides are acidic or neutral in nature, they are also called acid anhydrides

reaction of non-metals with water and also acids

they do not react with water, as they cannot give electrons to reduce hydrogen ions into hydrogen gas

reaction of non-metals with salt

a more reactive non-metal displaces a less reactive non-metal from its salt solution

NaBr + Cl2 → NaCl + Br2

reaction of non-metals with chlorine

non-metal + chlorine → covalent chloride

H2 + Cl2 → HCl

P4 + Cl2 → PCl3

reaction of non-metals with hydrogen

non-metal + hydrogen → covalent hydride

H2 + S → H2S

N2 + H2 → NH3

non-metals form covalent hydrides because they cannot give electrons to hydrogen atoms to form hydride ion

inertness of noble gases

as noble gases have 8 (or 2) electrons in their outermost shell, they have a stable configuration and do not take part in the sharing or exchange of electrons

chemical bonds

• atoms combine with one another to attain the inert gas electron configuration and become stable

• the force of attraction between them is called a chemical bond

• they are of two types

  • ionic bonds

  • covalent bonds

ions

a neutral atom has an equal number of electrons and protons. when it loses or gains an electrons, it results in an unequal amount of electrons and protons, which gives the atom a charge. this atom is called an ion

protons > electrons : positively charged ions called cations

electrons > protons : negatively charged ions called anions

ionic bond

• an ionic bond is formed by the transfer of electrons from one atom to another

• it occurs between metals and non-metals

• the strong force of attraction between oppositely charged ions is called the ionic bond

• it is made of ions and not molecules

covalent bond

• a covalent bond is the chemical bond formed by the sharing of electrons between two atoms

• it is formed between non-metals only, as the reacting atoms both need electrons to achieve stability

• they are of 3 types

  • single covalent bond: sharing of one pair of electrons

  • double covalent bond: sharing of two pairs of electrons

  • triple covalent bond: sharing of three pairs of electrons

properties of ionic compounds

• they are crystalline solids

• they have high MP and BP

• they are soluble in water but insoluble in organic solvents

• they conduct electricity

properties of covalent compounds

• they are usually liquids or gases, due to the weak force of attraction

• they have low MP and BP

• they are insoluble in water but soluble in organic solvents

• they do not conduct electricity

occurrence of metals

as most metals are very reactive, they do not occur as free elements in nature but as compounds of oxides, carbonates etc. only less reactive metals like Au and Pt occur in free state. Cu and Ag can occur both in free state and as a compound

minerals and ores

• the natural materials in which metals/ metal compounds are found in the earth are called minerals

• the minerals from which the metal can be extracted conveniently and profitably are called ores

• all ores are minerals, but not all minerals are ores

rock salt

sodium chloride; NaCl

bauxite

aluminium oxide; Al2O3.2H2O

pyrolusite

manganese dioxide; MnO2

calamine

zinc carbonate; ZnCO3

zinc blende

zinc sulphide; ZnS

haematite

iron (III) oxide; Fe2O3

cuprite

copper (I) oxide; Cu2O

copper glance

copper (I) sulphide; Cu2S

cinnabar

mercury (II) sulphide; HgS

extraction of metal

• the process of obtaining a metal from its ore is called extraction of metal

• there are 3 steps involved in this process

  • concentration/ enrichment of ore

  • conversion of concentrated ore

  • refining of impure metal

concentration/ enrichment of ore

• ores are impure compounds and contain a large amount of impurities like sand, rocky material, limestone etc. called gangue

• the removal of this gangue is called concentration of ore

conversion of concentrated ore into metal (extraction)

the extraction of metal from a concentrated ore is the process of reduction of the metal compound

there are different methods of extraction depending on the reactivity of the metal

highly reactive metals (K, Na, Ca, Mg, Al): electrolysis of their molten chloride/ oxide

moderately reactive metals (Zn, Fe, Mn): roasting/calcination, and then reduction of their oxide with C, Al

less reactive metals (Hg, Cu): roasting of sulphide ore, and then reduction of their oxide

calcination and roasting

calcination is the process in which carbonate ore is heated strongly in the absence of air to convert it into a metal oxide

roasting is the process in which sulphide ore is heated strongly in the presence of air to convert it into a metal oxide

electrolytic reduction of molten chlorides/oxides

the chlorides/oxides of highly reactive metals (K, Na, Ca, Mg, Al) are extracted through electrolytic reduction

the chlorine/ oxygen is formed at anode (positive electrode) and the metal is formed at cathode (negative electrode)

NaCl →(electrolysis) Na + Cl2

Al2O3 →(electrolysis) Al + O2

extraction of calamine and zinc blende

• step 1: conversion of ore into metal oxide

  • calamine: ZnCO3 → ZnO + CO2 (calcination)

  • zinc blende: ZnS + O2 → ZnO + SO2 (roasting)

• step 2: reduction of metal oxide using carbon

  • ZnO + C → Zn + CO

thermite reaction

• as manganese and iron cannot be reduced by carbon, its reduced by aluminium

• MnO2 + Al → Mn + Al2O3 + heat (HIGHLY exothermic)

• Fe2O3 + Al → Fe + Al2O3 + heat (HIGHLY exothermic)

this is used in thermite welding to join broken parts of heavy objects. as the reaction is highly exothermic, the metal is obtained in molten state

refining of metals

the process of purifying impure metals is called refining of metals. it is mostly done by the process of electrolytic refining

the impure metal is made the cathode, the same metal but pure is made the anode and a water soluble salt of the metal is taken as electrolyte

on passing electricity, the impure metal from the cathode dissolves into the salt solution and the pure metal from the salt solution is deposited on the cathode

corrosion

the eating up of metals by the action of air, moisture or chemical on their surface is called corrosio

rusting and its prevention

the corrosion of iron due to the presence of oxygen and water is called rusting

rust is hydrated iron oxide; Fe2O3.xH2O

it can be prevented in the following ways

by painting the surface to prevent air and moisture to come in contact with the metal

by applying grease or oil

by galvanisation, which is depositing a thin layer of zinc metal on it (itll form zinc oxide and protect the metal)

by tin-plating or chromium plating, as they are resistant to corrosion

corrosion of aluminium

aluminium loses its shine soon after use due to the formation of aluminium oxide on its surface. this makes it resistant from further corrosion. this is proof that corrosion is sometimes useful

sometimes, to make the oxide layer thicker, a process called anodising is used (the object is made the anode during the electrolysis of H2SO4)

corrosion of copper

the copper objects lose their shine over time due to the formation of a layer of copper oxide. after even more time, the oxide reacts with carbon dioxide and water to form green copper carbonate. this copper carbonate corrodes the metal

since copper is not very reactive, the corrosion is very slow. corroded copper can be cleaned with dilute acid

why gold and platinum are used for jewellery

as gold and platinum are not very reactive, they do not corrode and have a bright shiny surface. hence, they are used for making jewellery

alloys

the various properties of a metal like malleability, ductility, hardness, vulnerability to corrosion etc. can be improved by mixing metals. this mixture is called an alloy

it is prepared by mixing various metals in molten state and then cooling it to room temperature

they are stronger, harder, and have lower MP and electric conductivity as compared to the pure metals

brass

alloy of Cu & Zn (80% and 20%)

bronze

alloy of Cu & Sn (90% and 10%)

solder

alloy of Pb & Sn (50% and 50%)

alloys of gold

purity of gold is measured in ‘carats’. pure gold consists of 24 carats, however it is very soft and cannot be used for jewellery

therefore, 22 carat gold is used for jewellery, which is an alloy of gold and silver/copper