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ACIDS, BASES AND SALTS

ACIDS
Acid-is a Latin word for “sour”. The sour taste usually shows that a substance contains an acid.
There two types: Organic & Mineral acids. Mineral acids react rapidly than organic acids and are
described as strong acids and organic acids as weak acids.

An acid- is a substances that gives/donates hydrogen ions (H+), as the only positive ions when
dissolved in water. They are also defined as proton donors.

When an acid is dissolved in water it split into ions and this is called ionization.

e.g.

HCl(aq) H+(aq) + Cl-(aq)

H2SO4(aq) 2H+(aq) + SO42-(aq)

HNO3(aq) H+(aq) + NO3-(aq)

NB: All acids are covalent compounds but only show ionic and acidic properties in aqueous
solution. e.g. Turns blue litmus paper red.

Acids are classified as:

(a) STRONG ACIDS-are acids which ionize completely in solution giving H+ ions i.e All
molecules separate into ions and these are the mineral/inorganic acids.

e.g.

HCl(aq) H+(aq) + Cl-(aq)

i.e. HCl is a strong acid because all molecules present will split into ions resulting in high
concentration of hydrogen ions (H+).

Examples

Hydrochloric acid HCl(aq)

Sulphuric acid H2SO4(aq)

Nitric acid HNO3(aq)

Phosphoric acid H3PO4(aq)

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(b) WEAK ACIDS-are acids which dissociated/ionize partially in solution.i.e. They don’t
ionize completely, some acid molecules remain unseparated and result in low concentration of
hydrogen ions (H+). Weak acids are indicated by a reversible arrow ( ) in aqueous solution
and are organic acids.

e.g.

H2CO3(aq) 2H+(aq) + CO32-(aq)

CH3COOH(aq) H+(aq) + CH3COO-(aq)

Examples

Name Where found Formula

Carbonic acid In fizzy drinks H2CO3
Ethanoic/Acetic acid In vinegar CH3COOH/ CH3CO2H
Tartaric acid In porridge C4H6O6
Ascorbic acid In tablets (vitamin C) C6H8O6
Citric acid In fruits (e.g. lemon) C6H8O7

Properties of acids

 They all dissolve to give H+ ions.

 They react with most metals to give salt & hydrogen gas except nitric acid which
produces oxides of nitrogen.
e.g. Mg(s) + H2SO4(aq) MgSO4(aq) + H2(g)

Test for Hydrogen gas: Use a burning/lighted splint.

Test results: produces a “pop” sound.

 They react with metal carbonates giving salt, carbon dioxide gas and water.
e.g. CaCO3(s) + 2HCl(aq) CaCl2(caq) + CO2(g) + H2O(l)

Test for Carbon dioxide gas: bubble the gas in limewater.

Test results: white ppt of calcium carbonate is formed.

 They react with bases giving salt and water only. This is called neutralisation
reaction.ie. It turns universal indicator paper green.
e.g. (a) NaOH + HCl NaCl + H2O
(b) CuO + H2SO4 CuSO4 + H2O

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Uses of acids

 Strong acids are used to remove rust from metals

 Weak acids are used to preserve food and other things.e.g. Vinegar,citric acid and benzoic acid
 Manufacture of ; paints, fertilizers and detergents
 Sulphuric acid is used as an electrolyte in car batteries.
 Medicine eg ascorbic acid or vitamin C tablets.
 Flavouring food (vinegar)
 Nitric acid is used to make fertilizers and an explosive called Trinitrotoluene (TNT)

BASES

Base-is an oxide or hydroxide of a metal or ammonium hydroxide. A base reacts with an

acid to give salt and water only.

Water soluble bases are also called alkalis. They accept H+ ions from acids. Alkalis form OH- ions in
solution.

Eg. NaOH(aq) OH-(aq) + Na+(aq)

Bases and alkalis can be made of:

 Metal hydroxide (metal ion & OH- ion) e.g. NaOH, KOH, Ca(OH)2 etc.
 Metal oxides e.g. CuO, Fe2O3, ZnO etc.
 Metal carbonates (metal ion & CO32-) e.g. CaCO3, Na2CO3 etc.
 Metal hydrogen carbonate (Bicarbonate) e.g. Mg(HCO3)2, Ca(HCO3)2 etc.
 Ammonium hydroxide (NH4OH)
 Ammonium Carbonate ((NH4)2CO3)

Classification of Bases
(a) Soluble bases – they are oxides which dissolve in water giving hydroxide solutions called
alkalis.
e.g.
 Potassium hydroxide
 Calcium hydroxide
 Sodium hydroxide

Another unusual alkali is called aqueous ammonia or Ammonium hydroxide.

NB: All oxides of group 1 elements are soluble, calcium oxide is partially soluble.

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(b) Insoluble bases- they do not dissolve in water and they are mostly oxides of transition
elements
NB: “All Alkalis are Bases but NOT all bases are Alkalis.”

Properties of bases and alkalis:

 Bitter taste
 Strong alkalis are soapy feel to touch as they react with natural oil in the skin making soap.
 Have pH’s above 7
 Strong alkalis are corrosive eg. NaOH, KOH
 They turn red litmus paper blue
 All alkalis dissolve in water and give solutions which contain hydroxide ions
 Alkalis will react with most metal ions to form insoluble precipitates (ppt)

e.g. Cu(NO3)2 + NaOH Cu(OH)2 + NaNO3

copper(ii) hydroxide is a blue insoluble precipitate

 Bases neutralize acids(neutralization reaction)

e.g. NaOH + H2SO4 …………………+……………..

 Strong alkalis displace ammonia gas from ammonium salts

e.g. ammonium salt + alkali salt + ammonia gas + water

Like acids, alkalis' strength is determined by its ability to be ionized into metal and hydroxide OH- ions.
Completely ionized alkalis are the strongest and partially ionized alkalis are the weakest. Ammonium
hydroxide ( or aqueous Ammonia) is one of the weakest alkalis while strong alkalis include the
hydroxides of sodium, potassium and magnesium.

Uses of bases

 treatment of indigestion(milk of magnesia AKA calcium hydroxide and antiacid tablets)

 in toothpaste to neutralize acid in the mouth (Mg(OH)2)
 dissolve grease from surfaces
 manufacture of detergents (cleaning chemicals)
 treatment of insect stings.
 Controlling Soil pH:
If the pH of the soil goes below or above 7, it has to be neutralized using an acid or a base. If the
pH of the soil goes below 7, calcium carbonate (lime stone) is used to neutralize it. The pH of the
soil can be measured by taking a sample from the soil, crushing it, dissolving in water then
measuring the pH of the solution.

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STRENGTH & CONCENTRATION

CONCENTRATION

-it is the amount of solute in a given volume of a solvent.

-it is measured in mol/dm3.

STRENGTH

-it is the extent at which an acid or alkali separates/dissociates/ionize into ions in a solution.

RE: Strong acid/alkali dissociates completely in water while weak acid/alkali dissociates
partially in water.

Strength can be determined using a universal indicator or pH scale

How acidic or basic a substance is shown by its pH, therefore pH can be defined as the measure of the
degree of acidity or alkalinity of a solution.

pH Scale:

This is a scale that runs from 0 to 14. Substances with a pH below 7 are acidic. Substances with pH above
7 are basic. And those with pH 7 are neutral.

INDICATORS:

Indicators are substances that identify acidity or alkalinity of substances. They cannot be used in solid
form. They change from one colour when mixed with an acid or when mixed with an alkali.

Universal Indicator:
This is a substance that changes color when added to another substance depending on its pH. It is a
mixture of several indicators used to distinguish strong and weak acids as well as alkalis. The indicator
and the substance should be in aqueous form. The colours and strength on the universal indicator are as
follows:

Red Orange Yellow Green Blue Indigo Violet

Strong Acid Weak Acid Very Weak acid Neutral Very Weak Alkali Weak Alkali Strong Alkali

To remember the colours use the abbrieviations; ROY G. BIV, the name of a boy.

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Litmus Paper or Solution:

This indicator is present in two colors: red and blue. We use blue litmus if we want test a substance for
acidity. We use red litmus if we want to test a substance for alkalinity. Its results are:

 Acids: Turns blue litmus paper/ solution red,

 Bases: Turns red litmus paper/ solution blue,
 Neutral: if it is used as paper the color doesn’t change. If it is used as solution it turns purple.

Note: Use damp litmus paper if testing gases.

Phenolphthalein:
This is an indicator that is used to test for alkalinity because it is colorless if used with an acidic or neutral
substance and it is pink if it is used with a basic substance. It changes colour at around pH 9.5

Methyl Orange:
This indicator gives fire colors: Red with acids, yellow with neutrals and orange with bases.it changes
colour at around pH 4.5

BASICITY OF AN ACID

It is the number of hydrogen ions that an acid can be produced by one molecule of an acid.

(a) MONOBASIC ACID: produces one (1) H+ ion.

e.g. HNO3(aq) H+(aq) + NO3-(aq)
HCl(aq) H+(aq) + Cl- (aq)

(b) DIBASIC ACID: produces two (2) H+ ions.

e.g. H2 SO4(aq) 2H+(aq) + SO42-(aq)
H2 CO3(aq) 2H+(aq) + CO32-(aq)

(c) TRIBASIC ACID: produces three (3) H+ ions.

e.g. H3 PO4(aq) 3H+(aq) + PO43-(aq)

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OXIDES

These are compounds with oxygen atoms or oxide ions attached to other elements. They are formed when
elements are burned or react with oxygen. Rust is an example of a metal oxide(iron reacting with oxygen
in the presence of water) that’s why it is encouraged to keep iron metals coated with unreactive elements
(i.e galvanized) to prevent them from reacting with oxygen and water.

Types of Oxides:

Oxides

Metal oxides Non metal oxides

Basic oxides, Amphoteric oxides Acidic oxides, Neutral oxides

Acidic Oxides
Basic Oxides
Amphoteric Oxides -
metal oxides except
oxides non-metal
monoxides
Aluminum, Zinc &
acids forming a salt
Lead
and water
alkali to form salt
when reacting with
and water
an alkali & vice versa
in water except : metal
group 1 metal monoxides are
hydroxides are
oxides. Neutral oxides
amphoteric too
(they don’t react with
acid forming salt acids or alkalis)
and water when
and Examples: CO2,
reacting with an acid
water NO2, SO2 (acidic
or an alkali.
Examples: Na2O, oxides) & CO, NO,
CaO and CuO H2O (neutral
oxides)

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SALTS:

A salt is a neutral ionic compound. Salts are one of the products of a reaction between an acid and a base.
Salts are formed in reactions in which the H+ ion (acidic hydrogen of an acid) from the acid is replaced
by any other metal ion. Some salts are soluble in water and some are insoluble.

Insoluble Salts:
Soluble Salts:

(AgCl & PbCl2)

AgCl and PbCl2
lead sulphates (CaSO4,
All sulphates EXCEPT
BaSO4, PbSO4)
CaSO4, BaSO4, PbSO4
group 1 metals and
ammonium carbonates

PREPARATION OF SALTS

A. PREPARING SOLUBLE SALTS

For soluble salts, there are four methods that can be used and the general reaction scheme as well as the
procedure for the first three methods is as follows;
General reaction scheme:

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1. ACID + METAL → SALT + HYDROGEN

Note: this type of method is suitable to for making salts of moderately reactive metals because highly
reactive metals like K, Na and Ca will cause an explosion. This method is used with the MAZIT
(Magnesium, Aluminum, Zinc, Iron and Tin) metals only and is a displacement reaction method as an
acidic hydrogen of an acid is displaced by a metal.

Example: 2HCl(aq) + Mg(s) → MgCl2(aq) + H2(g)

NB: The procedure is as in the general reaction scheme and it will be as follows; (a)(i) then, (b), (c) and
(d).

Observations of this type of reactions:

 Bubbles of colorless gas evolve (hydrogen). To test for Hydrogen gas: a lighted/burning splint is
placed at the mouth of the test tube with the gas. Results: Hydrogen burns with a pop sound
 The temperature rises (exothermic reaction)
 The metal disappears

You will know the reaction is over when:

 No more gas bubbles evolve

 No more magnesium can dissolve
 The temperature stops rising
 The solution becomes neutral or can no longer turns blue litmus paper red.

Note: The following reaction methods are used to make salts of metals below hydrogen in the
reactivity series. If the base is a metal oxide or metal hydroxide, the products will be salt and water
only. If the base is a metal carbonate, the products will be salt, water and carbon dioxide.

2. ACID + METAL OXIDE → SALT + WATER

Example: To obtain copper(ii) sulphate salt given copper(ii) oxide and sulphuric acid:

H2SO4(aq) + CuO(s) → CuSO4(aq) + H2O(l)

NB: The procedure is as in the general reaction scheme and it will be as follows; (a)(ii) then, (b), (c) and
(d).

Observations of this reaction:

 The amount of copper(ii) oxide decreases

 The solution changes color from colorless to blue
 The temperature rises

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You will know the reaction is over when

 No more copper(ii) oxide dissolves

 The temperature stops rising
 The solution becomes neutral or can no longer turns blue litmus paper red.

3. ACID + METAL CARBONATE → SALT + WATER + CARBON DIOXIDE

Example:To obtain copper(ii) sulfate salt given copper(ii) carbonate and sulfuric acid:

CuCO3 (s) + H2SO4(aq) → CuSO4(aq) + H2O(l) + CO2(g)

NB: The procedure is as in the general reaction scheme and it will be as follows; (a)(iii) then, (b), (c)
and (d) BUT the reaction takes place at room temperature.

Observations:

 Bubbles of colorless gas (carbon dioxide) evolve, Test; bubble the gas through limewater
(aqueous calcium hydroxide), Results: Limewater turns milky or a white ppt(precipitate) is seen.
 Green Copper(ii) carbonate starts to disappear
 The temperature rises
 The solution turns blue

You will know the reaction is finished when:

 No more bubbles are evolving

 The temperature stops rising
 No more copper(ii) carbonate can dissolve.
 The pH of the solution becomes neutral or can no longer turns blue litmus paper red

4. ACID + METAL HYDROXIDE → SALT + WATER

This method of salt preparation is known as a Titration method.

This is a method to make a neutralization reaction between a base and an acid producing a salt without
any excess. In this method, the experiment is preformed twice, the first time is to find the amounts of
reactants to use, and the second experiment is the actual one.

Example: To obtain sodium chloride crystals given sodium hydroxide and hydrochloric acid:

HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

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The procedure will be as follows:

1st Experiment:

1. Add 25 cm3 of sodium hydroxide using a pipette to be accurate to flask

2. Add 2-3 drops of methyl orange indicator to the sodium hydroxide. The solution turns yellow
indicating the presence of a base
3. Fill a burette to zero mark with hydrochloric acid
4. Add drops of the acid to conical flask with continuous swirling of the conical flask
5. When the solution turns light pink, stop adding the acid (ie. The End point: is the point at which
every base molecule is neutralized by an acid molecule)
6. Record the amount of hydrochloric acid used and repeat the experiment (ie 2nd experiment)
without using the indicator.

After the 2nd experiment, you will have a sodium chloride solution. Evaporate it till dryness to obtain
powdered sodium chloride or crystalize it as in the general reaction scheme mentioned before from (c)
to (d) to obtain sodium chloride crystals

Experimental set-ups;

in general reaction scheme.

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PREPARING INSOLUBLE SALTS:

Insoluble salts are prepared by mixing/reacting together solutions of two SOLUBLE SALTS, each
containing half of the required salt. This type of reaction is called precipitation reaction AKA double
decomposition. The products of a precipitation reaction are two other salts, one of them is soluble and
the other one is insoluble (precipitate) ie. a solid.

Example: To obtain barium sulphate, you have to select a soluble salt of barium eg. Barium chloride
and a soluble salt containing the sulphate ions eg. Sodium sulphate.

Reaction Equation: BaCl2(aq) + Na2SO4(aq) → BaSO4(s) + 2NaCl(aq)

Ionic Equation: Ba2+(aq) + SO42-(aq) → BaSO4(s)

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Procedure:

1. Dissolve the two soluble salts in two different beakers.

2. Mix the two solutions and the reaction will readily take place.
3. When the reaction is over, filter and take the residue (ppt).
4. Wash the residue with distilled water and dry it in the oven/between filter papers.

Observations:

 Temperature increases
 An insoluble solid, precipitate (eg. Barium sulphate) forms

You know the reaction is over when:

 The temperature stops rising

 No more precipitate is being formed

Question 2

Name the reactants that can be used to prepare the following INSOLUBLE SALTS by PRECIPITATION

A. Lead(ii) iodide ………….. ….................... ………………………………

B. Calcium carbonate ………….. ….................... ………………………………

C. Silver chloride ………….. ….................... ………………………………

D. Barium sulphate ………….. ….................... ………………………………

E. Calcium sulphate ………….. ….................... ………………………………

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QUALITATIVE ANALYSIS (TEST FOR IONS)

It is the analysis of ions from samples using colours of their precipitates and types of gases produced in a
chemical reaction

Colors of Salts:

Salt Formula Solid In Solution

Hydrated copper sulphate CuSO4.5H2O Blue crystals Blue
Anhydrous copper sulphate CuSO4 White powder Blue
Copper nitrate Cu(NO3)2 Blue crystals Blue
Copper chloride CuCl2 Green Green
Copper carbonate CuCO3 Green Insoluble
Copper oxide CuO Black Insoluble
Iron(II) salts E.g.: FeSO4, Fe(NO3)2 Pale green crystals Pale green
Iron(III) salts E.g.: Fe(NO3)3 Reddish brown Reddish brown

Tests for Gases:

Gas Formula Tests

Ammonia NH3 Turns damp red litmus paper blue
Carbon dioxide CO2 Turns limewater milky
Oxygen O2 Relights a glowing splint
Hydrogen H2 ‘Pops’ with a lighted splint
Chlorine Cl2 Bleaches damp litmus paper
Nitrogen dioxide NO2 Turns damp blue litmus paper red
Sulphur dioxide SO2 Turns acidified aqueous potassium dichromate(VI) from orange to green

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Tests for Anions:

Anion Test Result
Effervescence,
Carbonate (CO32-) Add dilute acid carbon dioxide produced forms a white
ppt when bubbled through limewater
Chloride (Cl-) Acidify with dilute nitric acid, then add
White ppt.
(in solution) aqueous Silver nitrate or Lead (ii) nitrate
Iodide (I-) Acidify with dilute nitric acid, then add
Yellow ppt.
(in solution) aqueous Silver nitrate or Lead (ii) nitrate
Nitrate (NO3-) Add aqueous sodium hydroxide, then Ammonia is produced and it turns red
(in solution) aluminium foil; warm carefully litmus paper blue.
Sulfate (SO42-) Acidify, then add aqueous barium nitrate White ppt. of BaSO4 is formed

NOTE; specific solutions are used to test for specific anions

For the Nitrate (NO3-) ions, ammonia is produced in two steps:

1. Nitrate ion from the sample is reduced to ammonium ion by the nascent hydrogen produced by
aluminium metal and sodium hydroxide.
i.e. NO3-(aq) + 8[H+] NH4+(aq) + H2O(l) + 2OH-(aq)

2. The ammonium ion reacts with the hydroxide ion to produce ammonia.
i.e. NH4+(aq) + OH-(aq) NH4(g) + H2O(l)

NB: To test for the Cl-, I- & SO42- ions, an acid (nitric acid) is added to destroy any carbonate which may
interfere with results in the sample to form soluble nitrate, CO2 and water

Tests for Aqueous Cations:

Cation Effect of aqueous sodium hydroxide Effect of aqueous ammonia
White ppt., soluble in excess giving a
Aluminium (Al3+) White ppt., insoluble in excess
colourless solution
Ammonium (NH4+) Ammonia produced on warming –
Calcium (Ca2+) White ppt., insoluble in excess No ppt. or very slight white ppt.
Light blue ppt., soluble in excess,
Copper (Cu2+) Light blue ppt., insoluble in excess
giving a dark blue solution
Iron(II) (Fe2+) Green ppt., insoluble in excess Green ppt., insoluble in excess
Iron(III) (Fe3+) Red-brown ppt., insoluble in excess Red-brown ppt., insoluble in excess
White ppt., soluble in excess giving a
Lead ( Pb2+) White ppt., insoluble in excess
colourless solution
White ppt., soluble in excess, White ppt., soluble in excess,
Zinc (Zn2+)
giving a colourless solution giving a colourless solution

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NOTE; Al3+, Pb2+, Zn2+, Ca2+ are all white precipitates hence excess sodium hydroxide solution and
excess aqueous ammonia may be used to differentiate them BUT because the Al3+ and Pb2+ ions behave
the same, they can be differentiated by a solution containing the iodide ions and a yellow ppt will be
formed with Pb2+ ions.

NB: Excess solutions are added to test for the solubility of precipitates.

PRACTICE QUESTIONS

1.The following table shows the tests on substance W and the conclusions made from the observations.

Complete the table by describing these observations and suggest the test and observations which led to
the conclusion from test 4
Test Observations Conclusion
1. W was dissolved in ……………………………………………….. W is not a compound of a
water and the solution ……………………………………………….. transition metal
divided into three parts ………………………………………………[1]
for tests 2, 3 and 4.
2. (a). To the first part, ………………………………………………..
aqueous sodium ………………………………………………..
hydroxide was added ………………………………………………[1]
until a change was seen.

(b). An excess of
aqueous sodium ……………………………………………….. W may contain Al3+ or Zn2+ ions
hydroxide was added to ………………………………………………..
the mixture from (a). ………………………………………………[1]
3. (a). To the second part ………………………………………………..
aqueous ammonia was ………………………………………………..
added until a change ………………………………………………[1]
was seen.

(b).An excess of aqueous ……………………………………………….. The presence of Zn2+ ions

ammonia was added to ……………………………………………….. confirmed.
the mixture from (a). ………………………………………………[1]
4. ……………………………………. ………………………………………………..
……………………………………. ………………………………………………..
……………………………………. ……………………………………………….. W contains Cl- ions
……………………………………. ………………………………………………[1]
…………………………………[2]
[8]

Conclusion:
The name for substance W is:……………………………………………………………………………………………………….
The formula for substance W would be………………………………………………………………………………………… [2]

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2. The following table shows the tests a student carried on substance S and the conclusions made from
the observations.

Complete the table by describing the expected observations and conclusion for test 1, 2, and 3.

Test Observations Conclusion

1. (a). S was dissolved in …………………………………………………. a gas was produced.
dilute nitric acid and the ………………………………………………… S is compound of a transition
solution was divide into ……………………………………………….[1] metal.
two parts for test 2 & 3.

(b). the gas produced was …………………………………………………..

tested with…………. …………………………………………………..
…………………………………..[1] ……………………………………………….[1] S contains CO32- ions
2. (a). To the first part, ………………………………………………….
aqueous sodium ………………………………………………….
hydroxide was added ……………………………………………….[1]
until a change was seen.

(b). An excess of aqueous ………………………………………………….

sodium hydroxide was ………………………………………………….
added to the mixture ………………………………………………[1] S may contain Fe2+ ions
from (a).
3. (a). To the second part …………………………………………………
aqueous ammonia was ………………………………………………..
added until a change was ……………………………………………..[1]
seen.

(b).An excess of aqueous ………………………………………………. The presence of Fe2+ ions

ammonia was added to ………………………………………………. confirmed.
the mixture from (a). ……………………………………………..[1]
[7]

Conclusion

The name and chemical formula for substance S is:

Name…………………………………………………………………….

Formula.……………………………………………………… [2]

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