Metals and Non-metals

Metals in their pure state are lustrous (shiny), hard, malleable (can be beaten into thin sheets; gold and silver are the most malleable), ductile (can be drawn into wires; one gram of gold gives a wire about 2 km long), good conductors of heat and electricity, and sonorous (they ring when struck). Silver and copper are the best conductors of heat; lead and mercury are comparatively poor. Most metals have high melting points, but gallium and caesium melt on the palm of the hand, and mercury is the only metal that is liquid at room temperature. Sodium, potassium and lithium are so soft that they can be cut with a knife. Iodine is a non-metal that is lustrous.

Non-metals are mostly gases (hydrogen, oxygen, nitrogen, chlorine) or solids (carbon, sulphur, phosphorus, iodine); bromine is the only liquid non-metal. They are neither malleable nor ductile, are brittle when solid, and are poor conductors of heat and electricity. Graphite is the exception: a non-metal that conducts electricity. Diamond, a form of carbon, is the hardest natural substance and has a very high melting point.

The examiner uses these properties in two ways. First, as direct one-mark questions: name a liquid metal and a liquid non-metal (mercury, bromine); name a non-metal that conducts (graphite); name the best conductor of heat (silver). Second, as reasoning: why copper and aluminium are used in electric wires (good conductors, ductile), why school bells are made of metal (sonorous), why cooking vessels are metal (good conductors of heat). Also remember that the NCERT activities with a hammer, a battery-bulb circuit and a burner are used to test malleability, conductivity and heat conduction; the exercise question on the hammer-battery-bulb set-up is answered from here.

One caution: because of exceptions like mercury, graphite and iodine, physical properties alone cannot classify an element. The chapter therefore moves to chemical properties, which are the real basis of classification.

Almost all metals combine with oxygen to form metal oxides: Metal + O₂ → metal oxide. Copper heated in air turns black because a layer of copper(II) oxide forms: 2Cu + O₂ → 2CuO. Aluminium forms aluminium oxide: 4Al + 3O₂ → 2Al₂O₃. Magnesium burns with a dazzling white flame: 2Mg + O₂ → 2MgO. Sodium burns vigorously: 4Na + O₂ → 2Na₂O.

Metal oxides are basic in nature. A few, however, react with both acids and bases to form salt and water; these are amphoteric oxides, and the two you must know are aluminium oxide and zinc oxide.
Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O
Al₂O₃ + 2NaOH → 2NaAlO₂ (sodium aluminate) + H₂O
ZnO + 2HCl → ZnCl₂ + H₂O
ZnO + 2NaOH → Na₂ZnO₂ (sodium zincate) + H₂O

Most metal oxides are insoluble in water, but sodium oxide and potassium oxide dissolve to give alkalis: Na₂O + H₂O → 2NaOH; K₂O + H₂O → 2KOH.

The vigour of reaction with oxygen is the first clue to reactivity. Potassium and sodium react so violently that they catch fire if left in the open, which is why they are stored under kerosene oil. Magnesium, aluminium, zinc and lead react slowly at room temperature and form a thin oxide layer that stops further oxidation; this protective layer is the reason aluminium, though reactive, can be used for utensils. Iron does not burn on heating but iron filings burn vigorously when sprinkled into a flame. Copper does not burn but forms the black oxide coating. Silver and gold do not react with oxygen even at high temperature.

For a 3-mark question the examiner usually wants: the general word equation, two balanced examples, the definition of amphoteric oxide with two examples and equations, and the note on sodium and potassium oxides dissolving in water. Write the equations with state symbols where you can.

Metals react with water to produce a metal oxide and hydrogen gas. If the oxide is soluble, it dissolves further to form the metal hydroxide. How vigorous the reaction is depends on the metal, and this gives a neat ladder of reactivity.

Potassium and sodium react violently with cold water. The reaction is so exothermic that the hydrogen evolved catches fire immediately:
2K + 2H₂O → 2KOH + H₂ + heat
2Na + 2H₂O → 2NaOH + H₂ + heat

Calcium reacts with cold water less violently. The heat evolved is not enough to ignite the hydrogen. Instead the calcium starts floating, because bubbles of hydrogen stick to its surface: Ca + 2H₂O → Ca(OH)₂ + H₂.

Magnesium does not react with cold water. It reacts with hot water to form magnesium hydroxide and hydrogen, and it also starts floating for the same reason: Mg + 2H₂O → Mg(OH)₂ + H₂.

Aluminium, iron and zinc do not react with cold or hot water, but they react with steam to form the metal oxide and hydrogen:
2Al + 3H₂O(steam) → Al₂O₃ + 3H₂
3Fe + 4H₂O(steam) → Fe₃O₄ + 4H₂
Zn + H₂O(steam) → ZnO + H₂

Lead, copper, silver and gold do not react with water at all.

Two marking details. First, iron with steam gives Fe₃O₄ (the mixed oxide), not Fe₂O₃ and not Fe(OH)₃; the coefficient set 3, 4, 1, 4 is asked often. Second, when the question says 'name the metal that starts floating', the answer is calcium for cold water and magnesium for hot water, and the reason is 'bubbles of hydrogen gas formed stick to its surface'. A common 2-mark question gives you a table of metals A, B, C, D with observations (reacts with cold water, hot water, steam, none) and asks you to arrange them in decreasing order of reactivity; use exactly this ladder.

Metals react with dilute acids to form a salt and hydrogen gas: Metal + dilute acid → salt + H₂.
Mg + 2HCl → MgCl₂ + H₂
2Al + 6HCl → 2AlCl₃ + 3H₂
Zn + 2HCl → ZnCl₂ + H₂
Fe + 2HCl → FeCl₂ + H₂
Note that iron gives iron(II) chloride, FeCl₂, with dilute HCl, not FeCl₃.

In the NCERT activity with dilute HCl, the rate of bubble formation is fastest for magnesium, then aluminium, zinc and iron; the temperature rise is also highest for magnesium. Copper does not react with dilute HCl or dilute H₂SO₄ at all, because it lies below hydrogen in the reactivity series. The hydrogen is tested by bringing a burning matchstick near the mouth of the tube: it burns with a pop sound.

Nitric acid is the exception. Hydrogen gas is not evolved when a metal reacts with nitric acid, because HNO₃ is a strong oxidising agent. It oxidises the hydrogen produced to water and is itself reduced to one of the oxides of nitrogen (N₂O, NO or NO₂). Only magnesium and manganese react with very dilute HNO₃ to give hydrogen. This point is a favourite Assertion–Reason item.

Aqua regia (Latin for 'royal water') is a freshly prepared mixture of concentrated hydrochloric acid and concentrated nitric acid in the ratio 3 : 1. It is highly corrosive and fuming, and it is one of the few reagents that can dissolve gold and platinum, even though neither acid alone can. The NCERT exercise story about the fake goldsmith who dipped a gold bangle in a solution and reduced its weight is answered with aqua regia.

For a 3-mark question, write the general equation, two or three balanced examples, the order of reactivity observed, and the nitric acid exception with the reason. If asked to name two metals that displace hydrogen from dilute acids and two that do not, answer Mg and Zn (or Fe, Al) versus Cu and Ag (or Au, Hg).

A more reactive metal displaces a less reactive metal from a solution of its salt. This is a displacement reaction: Metal A + salt solution of B → salt solution of A + metal B, provided A is more reactive than B.

Iron nail in copper sulphate solution: the blue colour of the solution fades to pale green (iron(II) sulphate) and a reddish-brown deposit of copper forms on the nail: Fe + CuSO₄ → FeSO₄ + Cu. Zinc in copper sulphate: the blue solution turns colourless and copper deposits: Zn + CuSO₄ → ZnSO₄ + Cu. Copper in silver nitrate: the solution slowly turns blue (copper nitrate) and a shining grey deposit of silver forms on the copper: Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag. In the reverse direction, copper placed in iron(II) sulphate or zinc sulphate does nothing, and silver in copper sulphate does nothing, because the metal added is less reactive than the one in the salt.

The NCERT activity with iron, copper, zinc and aluminium strips in solutions of their sulphates lets you rank the four: Al > Zn > Fe > Cu. In the exam the usual question gives a set of metals and salt solutions and asks (a) in which cases a reaction occurs, (b) what you will observe, and (c) the equation. Answer all three parts for each case; observations are marked separately from equations, and the colour change (blue to green, or blue to colourless) is worth its own half mark.

Displacement reactions are also the reason for the NCERT exercise MCQ: of the pairs NaCl + Cu, MgCl₂ + Al, FeSO₄ + Ag and AgNO₃ + Cu, only AgNO₃ + Cu reacts, because copper is above silver in the series while in the other three the added metal is below the metal in the salt.

The reactivity series (activity series) is a list of metals arranged in the decreasing order of their reactivity, prepared from displacement experiments. From most to least reactive: K > Na > Ca > Mg > Al > Zn > Fe > Pb > [H] > Cu > Hg > Ag > Au. Hydrogen is not a metal, but it is placed in the series because it lets you predict which metals will displace hydrogen from dilute acids: metals above hydrogen do, those below it (copper, mercury, silver, gold) do not.

A good memory line is 'Please Send Cats, Monkeys And Zebras In Large Cages, Make Sure Properly Guarded' (K, Na, Ca, Mg, Al, Zn, Fe, Pb, [H], Cu, Hg, Ag, Au... adjust the last letters as you like), but check the order once against the textbook.

Applications the examiner expects when asked 'state two uses of the reactivity series':

• Predicting whether a displacement reaction will occur (a metal displaces any metal below it from its salt solution).
• Predicting the reaction of metals with water and dilute acids (metals above hydrogen liberate H₂; the higher the metal, the more vigorous the reaction).
• Choosing the method of extraction: metals at the top (K, Na, Ca, Mg, Al) need electrolysis, those in the middle (Zn, Fe, Pb, Cu) are reduced with carbon, those at the bottom (Hg, Cu, Ag, Au) are obtained by heating alone or occur free.
• Choosing a reducing agent for displacement (thermit) reactions, for instance aluminium reducing Fe₂O₃.
• Explaining why some metals occur free in nature (Au, Ag, Pt) while the reactive ones occur only as compounds.

Another standard 3-mark question describes a metal A that displaces iron from FeSO₄ but not aluminium from Al₂(SO₄)₃, and asks you to place it; A lies between aluminium and iron, so zinc is the natural example. Questions of the form 'arrange in decreasing order of reactivity' from a set of observations with water, acids or salt solutions are all answered by matching the observations to positions in this list.

The reactivity of an element is the tendency to attain a completely filled valence shell (the stable noble-gas configuration). Metals have 1, 2 or 3 valence electrons and lose them to form positive ions (cations); non-metals have 5, 6 or 7 valence electrons and gain electrons to form negative ions (anions). The electrostatic attraction between the oppositely charged ions is an ionic (electrovalent) bond, and compounds formed this way are ionic compounds.

Sodium chloride. Sodium (2,8,1) loses one electron to get the neon configuration (2,8) and becomes Na⁺. Chlorine (2,8,7) gains that electron to get the argon configuration (2,8,8) and becomes Cl⁻.
Na → Na⁺ + e⁻
Cl + e⁻ → Cl⁻
Na⁺ + Cl⁻ → NaCl (written as Na⁺Cl⁻). Sodium chloride does not exist as molecules but as a lattice of oppositely charged ions.

Magnesium chloride. Magnesium (2,8,2) loses two electrons to become Mg²⁺; each of two chlorine atoms gains one electron. Mg → Mg²⁺ + 2e⁻; 2Cl + 2e⁻ → 2Cl⁻; Mg²⁺ + 2Cl⁻ → MgCl₂.

Calcium oxide. Calcium (2,8,8,2) loses two electrons to become Ca²⁺; oxygen (2,6) gains two electrons to become O²⁻: CaO. Magnesium oxide (Mg²⁺O²⁻) and sodium oxide (2Na⁺ O²⁻) are drawn the same way.

In the exam, draw the electron-dot (Lewis) structures: write the symbol, show only the valence electrons as dots or crosses, use an arrow to show the transfer, and then draw the ions in square brackets with their charges. Show the electronic configuration of each atom before and after transfer; the marking scheme gives a mark for the configurations, a mark for the transfer diagram and a mark for the formula of the compound. For a 5-mark question on NaCl or MgCl₂, this diagram is followed by the properties of ionic compounds (next section), so treat the two sections as one answer.

Because ionic compounds are made of ions held in a rigid lattice by strong electrostatic forces, their properties follow directly from the bonding, and every property should be answered with its reason.

1. Physical nature: ionic compounds are hard solids because of the strong force of attraction between the positive and negative ions. They are generally brittle and break into pieces when pressure is applied.
2. Melting and boiling points: they have high melting and boiling points because a large amount of energy is needed to break the strong inter-ionic attraction. NCERT melting points (in kelvin): LiCl 887, NaCl 1074, KCl 1045, CaCl₂ 1045, MgCl₂ 981, CaO 2850, MgO 3073. Boiling points: NaCl 1686 K, MgO 3873 K.
3. Solubility: generally soluble in water and insoluble in solvents such as kerosene and petrol.
4. Conduction of electricity: they do not conduct electricity in the solid state because the ions are held in fixed positions and cannot move. They conduct in the molten state and in aqueous solution, because heat or water overcomes the attraction and the ions become free to move towards the electrodes.

Two experimental facts from the NCERT activity are worth quoting: the samples (NaCl, KI, BaCl₂) do not melt easily on a burner, and a bulb in the circuit glows only when the electrodes dip into a solution of the salt, not when they touch the dry crystals.

The standard HOTS question: 'A metal X loses two electrons and a non-metal Y gains one electron. Give the formula, draw the electron-dot structure, and state whether the compound has a high or low melting point and whether it conducts in the solid state.' The formula is XY₂; it is ionic; high melting point; conducts only when molten or dissolved; soluble in water. Case-based questions also give the melting point table and ask why MgO melts higher than NaCl: the ions in MgO carry double charges (Mg²⁺, O²⁻), so the attraction is stronger. That reason is beyond strict NCERT but is accepted.

The earth's crust is the major source of metals; sea water contains soluble salts such as sodium chloride and magnesium chloride. The elements or compounds that occur naturally in the crust are minerals. A mineral from which a metal can be extracted profitably and conveniently is an ore. The unwanted sand and rocky material mixed with an ore is called gangue. These three definitions are a regular 2- or 3-mark question.

Where a metal occurs depends on its reactivity. Metals at the bottom of the series (gold, silver, platinum, and some copper) are so unreactive that they are found in the free (native) state. Metals in the middle (Zn, Fe, Pb, Cu) are moderately reactive and occur as oxides, sulphides or carbonates. Metals at the top (K, Na, Ca, Mg, Al) are so reactive that they never occur free; they are found as chlorides, oxides or carbonates. Ores of many metals are oxides, because oxygen is very reactive and abundant.

Extraction of a metal from its ore is metallurgy, and it has three broad stages: (1) concentration or enrichment of the ore, which removes the gangue by physical or chemical methods chosen according to the properties of the ore and the gangue (the specific methods are not in the current syllabus); (2) extraction of the crude metal from the concentrated ore; (3) refining of the crude metal to obtain the pure metal. The method used in stage 2 depends entirely on the position of the metal in the reactivity series:

• Low in the series (Hg, Cu, Ag): the oxide is reduced by heating alone.
• Middle of the series (Zn, Fe, Pb, Cu): the ore is first converted to the oxide (roasting or calcination) and then reduced by carbon or by a more reactive metal.
• Top of the series (K, Na, Ca, Mg, Al): the molten chloride or oxide is reduced by electrolysis, because carbon cannot reduce these oxides.

When a question asks 'name the method used to extract sodium and give the reason', the full answer is: electrolytic reduction of molten NaCl, because sodium is highly reactive and its oxide cannot be reduced by carbon; in fact sodium has a greater affinity for oxygen than carbon has.

Metals near the bottom of the reactivity series are very unreactive, so their oxides can be reduced to the metal simply by heating. The two examples the board asks are mercury from cinnabar and copper from copper glance.

Mercury from cinnabar (HgS). Cinnabar is heated in air, which converts the sulphide to the oxide; the oxide then decomposes on further heating to give mercury:
2HgS(s) + 3O₂(g) → (heat) → 2HgO(s) + 2SO₂(g)
2HgO(s) → (heat) → 2Hg(l) + O₂(g)

Copper from copper glance (Cu₂S). Part of the sulphide is roasted to the oxide, and the oxide then reacts with the remaining sulphide to give copper:
2Cu₂S(s) + 3O₂(g) → (heat) → 2Cu₂O(s) + 2SO₂(g)
2Cu₂O + Cu₂S → (heat) → 6Cu(s) + SO₂(g)

The second copper equation is the one students drop. Notice that here the sulphide itself acts as the reducing agent for the oxide, so no carbon is needed. Both examples illustrate the general point that sulphide ores are first converted to oxides because a metal is more easily obtained from its oxide than from its sulphide.

Typical 2-mark question: 'Name the ore of mercury. How is mercury extracted from it? Write equations.' Give the name cinnabar, its formula HgS, both equations with 'heat' written above the arrows, and one line saying 'mercury is low in the reactivity series, so heating alone reduces the oxide'. A 3-mark question may instead give you 'copper is obtained from Cu₂S by heating in air' and ask for the two equations and the name of the process (roasting followed by self-reduction, though NCERT only calls it heating in air). For the case-based passage on this section, the fourth sub-question often asks why the same method will not work for zinc: zinc oxide is stable and needs carbon as a reducing agent.

Zinc, iron, lead and copper are moderately reactive. Their ores are usually sulphides or carbonates, and it is easier to obtain a metal from its oxide than from its sulphide or carbonate, so the ore is first converted to the oxide.

Roasting is heating a sulphide ore strongly in the presence of excess air to convert it into the oxide: 2ZnS(s) + 3O₂(g) → (heat) → 2ZnO(s) + 2SO₂(g).
Calcination is heating a carbonate ore strongly in a limited supply of air to convert it into the oxide: ZnCO₃(s) → (heat) → ZnO(s) + CO₂(g).
The difference the examiner looks for: roasting is for sulphide ores and uses excess air; calcination is for carbonate ores and uses limited air (or absence of air). Roasting gives SO₂; calcination gives CO₂.

Reduction. The oxide is then reduced by carbon, a reducing agent, when heated together: ZnO(s) + C(s) → Zn(s) + CO(g). Similar reductions give lead from PbO and iron from Fe₂O₃ in the blast furnace.

Displacement (thermit) reactions. Highly reactive metals such as sodium, calcium and aluminium can also be used as reducing agents, because they displace metals of lower reactivity from their compounds. These reactions are highly exothermic and the metal is produced in the molten state:
3MnO₂(s) + 4Al(s) → 3Mn(l) + 2Al₂O₃(s) + heat
Fe₂O₃(s) + 2Al(s) → 2Fe(l) + Al₂O₃(s) + heat
The reaction of iron(III) oxide with aluminium is the thermit reaction, used to join railway tracks and to repair cracked machine parts, because the molten iron produced flows into the joint. In MCQs, the metal used is aluminium and the oxide reduced is Fe₂O₃; aluminium is the reducing agent and Fe₂O₃ is the oxidising agent.

The 5-mark 'extraction of zinc' answer is: name the ores (zinc blende ZnS, calamine ZnCO₃), define roasting and calcination with equations, write the carbon reduction, and end with refining. Add the thermit reaction if the question says 'name a reaction in which a metal is used as reducing agent'.

Potassium, sodium, calcium, magnesium and aluminium are very reactive and have a very high affinity for oxygen. Their oxides cannot be reduced by carbon, because these metals hold on to oxygen more strongly than carbon does. Carbon is therefore useless here, and so is any other cheap chemical reducing agent. These metals are obtained by electrolytic reduction: the metal chloride or oxide is melted and electrolysed, and the metal is deposited at the cathode (the negative electrode), while chlorine or oxygen is liberated at the anode (the positive electrode).

For sodium from molten sodium chloride:
At cathode: Na⁺ + e⁻ → Na
At anode: 2Cl⁻ → Cl₂ + 2e⁻
Aluminium is obtained by the electrolytic reduction of molten aluminium oxide (from the ore bauxite); magnesium and calcium from their molten chlorides.

Two reasons are asked repeatedly and must be stated exactly. (1) 'Why is carbon not used to reduce the oxides of sodium, magnesium, calcium and aluminium?' Because these metals have a greater affinity for oxygen than carbon; they are stronger reducing agents than carbon, so carbon cannot take the oxygen away from them. (2) 'Why is the chloride electrolysed in the molten state and not in aqueous solution?' Because in solution the water would be reduced at the cathode to hydrogen instead of the metal being deposited (and sodium would react with water). The second reason goes slightly beyond NCERT but is a common follow-up.

Assertion–Reason items from this section pair 'Sodium is extracted by electrolysis of molten NaCl' with 'Sodium is highly reactive and its oxide cannot be reduced by carbon': both true, and the reason explains the assertion (option a). The marking scheme for a 3-mark question wants the method named, the reason for it, and the two electrode half-reactions with the electrodes identified.

The metal obtained by reduction is impure. The most widely used method of purifying it is electrolytic refining, used for copper, zinc, tin, nickel, silver, gold and others. For copper the set-up is:

• Anode (positive): a thick block of impure copper.
• Cathode (negative): a thin strip of pure copper.
• Electrolyte: an aqueous solution of copper sulphate acidified with a little dilute sulphuric acid.

When current is passed, pure copper from the anode dissolves into the electrolyte as Cu²⁺ ions, and an equivalent amount of copper from the electrolyte is deposited on the cathode:
At anode: Cu → Cu²⁺ + 2e⁻ (oxidation; the anode gets thinner)
At cathode: Cu²⁺ + 2e⁻ → Cu (reduction; the cathode gets thicker)
The soluble impurities pass into the solution, and the insoluble impurities, including valuable metals such as silver and gold, settle at the bottom of the anode as anode mud.

The diagram is asked with the description, and it needs: a tank with the electrolyte labelled 'acidified CuSO₄ solution', the impure copper anode connected to the positive terminal of the battery, the pure copper cathode connected to the negative terminal, anode mud shown as a deposit below the anode, and the direction of the current or of the Cu²⁺ ion movement.

Board papers have asked why the anode becomes thinner and the cathode thicker, and case-based questions have asked what anode mud contains and why the impurities settle rather than dissolve. Answer both from the half reactions: only copper (and metals more reactive than copper) dissolves at the anode, while less reactive impurities like Ag and Au do not oxidise and fall as mud. If the question is about 'metal M', simply replace copper with M in the set-up and say M is in the middle or low part of the series.

Corrosion is the slow eating away of a metal by the action of air, moisture or chemicals on its surface. Three examples the board asks: silver articles turn black because silver reacts with sulphur compounds (H₂S) in the air to form a coating of silver sulphide, Ag₂S; copper turns green in moist air because it forms basic copper carbonate, CuCO₃·Cu(OH)₂, with CO₂ and moisture; iron acquires a reddish-brown flaky coating of rust, hydrated iron(III) oxide, Fe₂O₃·xH₂O. Rusting of iron is a serious problem because it wastes huge amounts of money every year.

Conditions for rusting (NCERT activity). Take three test tubes with clean iron nails. Tube A: tap water (air and water both present). Tube B: boiled distilled water covered with a layer of oil (water but no air). Tube C: some anhydrous calcium chloride, corked (air but no moisture). After a few days only the nails in A rust. Conclusion: both air (oxygen) and moisture are necessary for rusting.

Prevention. Rusting is prevented by keeping air and water away from the iron surface: painting, oiling, greasing, galvanising (coating steel or iron with a thin layer of zinc; the zinc layer protects even after it is scratched, because zinc is more reactive than iron and corrodes first), chrome plating, anodising and alloying. Anodising is the process of forming a thick protective oxide layer on aluminium: the aluminium article is made the anode and electrolysed with dilute sulphuric acid, so oxygen liberated at the anode thickens the natural Al₂O₃ coating, which can also be dyed.

Alloys. An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal, made by melting the primary metal and dissolving the other elements in it. Pure iron is very soft and stretches easily when hot; mixing in about 0.05 % carbon makes it hard and strong. Stainless steel (Fe + Ni + Cr) does not rust. Brass = Cu + Zn; bronze = Cu + Sn; solder = Pb + Sn (low melting point, used to weld electrical wires). Brass and bronze are poor conductors compared with pure copper, and solder melts lower than either metal: alloying generally lowers electrical conductivity and melting point. An amalgam is an alloy containing mercury. Pure gold (24 carat) is too soft for jewellery, so it is alloyed with copper or silver to make 22-carat gold (22 parts gold, 2 parts Cu or Ag). Wrought iron is the purest form of iron. The 'galvanisation versus alloying' question (2020) wants both definitions and the point that galvanisation is a coating on the surface while alloying changes the whole metal.