Introduction To Matter: 1.1 Atoms & Molecules 1.1.1 Classification of Matter

© Lau Kah Pew STPM 2006

CHAPTER 1 INTRODUCTION TO MATTER 1.1 Atoms & Molecules 1.1.1 Classification of Matter 1.1.2 Introduction to Atoms 1.1.3 Mass Spectrometer 1.2 The Mole Concept 1.3 Matter & Measurements 1.3.1 Basics 1.3.2 Concentration Units 1.4 Chemical Formulae 1.4.1 Chemical Nomenclature 1.4.2 Empirical & Molecular Formula 1.5 Stoichiometry 1.5.1 Limiting Reagents 1.5.2 Reaction Yield 1.5.3 Solution Stoichiometry

1.1 Atoms & Molecules 1.1.1 Classification of Matter Chemistry is the study of matter and the changes it undergoes 1. Matter is anything that occupies space and has mass. 2. A substance is a form of matter that has a definite composition and distinct properties. e.g. water, ammonia, sucrose, gold, oxygen

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A mixture is a combination of two or more substances in which the substances retain their distinct identities. Homogenous mixture

– composition of the mixture is the same throughout. e.g. soft drink, milk, solder

Heterogeneous mixture – composition is not uniform throughout. E.g. cement, iron filings in sand Physical means can be used to separate a mixture into its pure components. An element is a substance that cannot be separated into simpler substances by chemical means. A compound is a substance composed of atoms of two or more elements chemically united in fixed proportions. Compounds can only be separated into their pure components (elements) by chemical means. Zn(s) + 2 HCl (aq)  ZnCl2(aq) + 2 H2(g) 2 H2O(l)  2 H2(g) + O2(g) Three states of matter: SOLID, LIQUID, GAS A physical change does not alter the composition or identity of a substance… 2 H2O(l)  H2O(g) A chemical change alters the composition or identity of the substance(s) involved… 2 H2O(g)  2 H2(g) + O2(g) MATTER – anything that occupies space and has mass. MASS

– measure of the quantity of matter. SI unit of mass is the kilogram (kg)

WEIGHT – force that gravity exerts on an object 2



1.1.2 Introduction to Atoms

Rutherford’s Model of Atom

nucleus electron cloud

Bohr’s Model of Atom

Modern Model of the Atom: Quantum Mechanics

Mass (g)

Particle

Charge (Coulombs)

Electron (e-) 9.1 x 10-28

Charge (units)

-1.6 x 10-19

-1

Proton (p+)

1.67 x 10-24 +1.6 x 10-19

+1

Neutron (n)

1.67 x 10-24

0

0

Proton number (Z) = No. of protons in nucleus Nucleon number (A) = No. of protons + No. of neutrons = Proton number (Z) + No. of neutrons

Nucleon Number Proton Number

A Z

X 3

Element Symbol



Exercise 1:

A Z

X

c

1. ______ is the proton number. 2. It is the total number of _________ in an atomic nucleus. 3. A is the _________________ of the nuclide X. 4. Nucleon number is defined as the total number of _________ and _________ in an atomic nucleus.

84 36

Kr

216 O 8

59 3+ Co 27

5. The nucleon number of Kr = _____ 6. The proton number of Co3+ = _____ 7. The number of neutrons in Kr = _____ 8. ____ contains 10 electrons. 9. Co3+ consists of ____ protons, ____ electrons and _____ neutrons.

Isotopes are atoms of the same element with different numbers of neutrons in their nuclei

Hydrogen atom isotopes:

1 1 protium

H

2 1 deuterium

3 1 tritium

Proton Number:

1

1

1

Nucleon Number:

1

2

3

Number of Neutrons:

0

1

2

4


H

H


Some isotopes are unstable. These unstable isotopes may “break-up” to form smaller atoms and subatomic particles. This process is known as radioactive decay. This process involves emmission of radioactive rays. For example… α-decay of uranium-235 235 92

U

β-decay of carbon-14 14 6

4 2

0 -1

C

He e

+

231 90

Th

+

14 7

N

1.1.3 Mass Spectrometer The mass spectrometer is machine used: - to determine the relative atomic mass of an element - to determine the relative molecular mass of a compound - to determine the types of isotopes, their relative isotopic masses, and the abundance of the isotopes of an element - to recognize the structure of an unknown compound

D B

B A

E D

A C A

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A) Vaporisation Chamber Sample of the element is vaporised into gaseous atom. A pump maintains a vacuum inside the mass spectrometer as any air molecules inside would block the movement of the ions and to avoid the contamination of the sample. B) Ionisation Chamber A hot filament emits high-energy electrons. When the electrons collide with the gaseous sample (atom or molecule), positive ions are produced by dislodging an electron from each atom or molecule C) Acceleration Chamber The positive ions are accelerated by an electric field towards the two oppositely charged plates. The electric field is produced by a high voltage between the two plates. The emerging ions are of high and constant velocity. D) Magnetic Field The positive ions are separated and deflected into a circular path by a magnet according to its m/e ratio. E) Ion Detector The numbers of ions and types of isotopes are recorded as a mass spectrum. Example: Mass Spectrum Of Rubidium The mass spectrum of rubidium shows that naturally occurring rubidium consists of two isotopes (two peaks): 85Rb and 87Rb.

Relative Abundance

18 7 85

87

m/e

The height of each line is proportional to the abundance of each isotope.

In this example, Rb-85 is more abundant than Rb-87. 6



Calculation a) What is the relative atomic mass of Rb?  (mi x Qi) Ar Rb =

 Qi

= =

(85 x 18) + (87 x 7) (18 + 7) 85.56

b) What is the percentage abundance of each of the isotopes? 18 x 100 = 72 % % 85Rb = 25 % 87Rb = 100 – 72 = 28 %

1.2 The Mole Concept 1.2.1 Basics RELATIVE ATOMIC MASS is the mass of an atom measured in relative to the mass of 1/12 the mass of a 12C atom. Relative Atomic Mass, Ar, is written in atomic mass units (amu). But sometimes, no units are used.. Notice that relative atomic mass of elements are usually not in round numbers. This is because most elements contain isotopes. Relative atomic mass of elements (that contain isotopes) are determined by calculating the average mass of all its naturally existing isotopes. The mole (mol) is the amount of a substance that contains as many elementary entities as there are atoms in exactly 12.000 grams of 12C 1 mol = NA = 6.0221367 X 1023 Avogadro’s number (NA)

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Mole (symbol = mol): The amount of substance that contains as many elementary particles as there are atoms in exactly 12.000 g of 12C. 1 mol = 6.022 x 1023 particles (Avogadro’s constant = 6.022 x 1023 mol1 ) 1 12C atom = 12.00 amu but… 1 mole 12C atoms = 6.022 X 1023 atoms = 12.00 g For any element, Relative Atomic Mass (amu) = Molar Mass (grams) For all substances, the molar mass in grams per mole is numerically equal to the formula weight in atomic mass units. Molar Mass = Ar or Mr (in g mol1) Example: How many atoms are in 0.551 g of potassium (K)? 1 mol K = 39.10 g K = 6.022 x 1023 atoms K 23 6.022 x 10 atom Number of K atoms = 0.551 g X 39.10 g = 8.49 x 1021 atoms

Relative Molecular Mass (or molecular weight) is the sum of the atomic masses (in amu) in a molecule. Due to the fact that the quantity of gaseous materials depends on its pressure, temperature and/or volume, 1 mol of any substance in this state would occupy different volumes. Room Temp. & Pressure (R.T.P.)

T = 25oC (298 K)

Standard Room Temp. & Pressure (S.T.P.)

T = 25oC (298 K)

P = 1 atm (101 325 Pa)

P = 1 atm (101 325 Pa) 8

24.0 L mol-1 22.4 L mol-1



1.3 Matter & Measurements 1.3.1 Basics S.I. Base Units Base Quantity Length Mass Time Current Temperature Amount of substance Luminous intensity

Name of Unit meter kilogram second ampere kelvin mole candela

Symbol m kg s A K mol cd

VOLUME – S.I. derived unit for volume is cubic meter (m3). but in chemistry, dm3 (or litres) are usually used. 1 cm3 = (1 x 10-2 m)3 = 1 x 10-6 m3 1 dm3 = (1 x 10-1 m)3 = 1 x 10-3 m3 1L

=

1,000 mL

= 1,000 cm3

= 1 dm3

1 mL = 1 cm3 DENSITY – S.I. SI derived unit for density is kg/m3 mass density = volume Significant Figures Any digit that is not zero is significant 1.234 kg…… 4 significant figures Zeros between nonzero digits are significant 606 m …… 3 significant figures Zeros to the left of the first non-zero digit are not significant 0.08 L …… 1 significant figure 9



Significant Figures If a number is greater than 1, then all zeros to the right of the decimal point are significant 2.0 mg …… 2 significant figures If a number is less than 1, then only the zeros that are at the end and in the middle of the number are significant 0.00420 g …… 3 significant figures

ACCURACY – how close a measurement is to the true value PRECISION – how close a set of measurements are to each other

1.3.2 Concentration Units The concentration of solutions is the quantity of dissolved substance per unit quantity of solvent in a solution. Concentration is measured in various ways: • Molarity • weight/volume per cent • Molal concentration (or Molality)

• mole fraction • parts per million

• weight per cent

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A) Molarity The number of mole of dissolved solute divided by the volume of the mixture. Symbol: M (sometimes the symbol “c” is used) Unit:

mol L1 or mol dm3 or M

M =

n (mol) Vsolution (L or dm-3)

Example: A student prepared a solution by dissolving 0.586 g of sodium carbonate, Na2CO3 in 250.0 cm3 of water. Calculate its concentration.

B) Molal Concentration (Molality) The number of mole of dissolved solute divided by the mass (in kg) of the solvent. Unit: mol Kg1 or molal or m

Symbol: m or M

M =

n (mol) msolvent (Kg)

Example: Calculate the molal concentration of ethylene glycol (C2H6O2) solution containing 8.40 g of ethylene glycol in 200 g of water. The molar mass of ethylene glycol is 62 g/mol.

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C) Weight per Cent (Per Cent By Mass) Mass ratio between solute and solution. Symbol: % w/w Unit: % 10% w/w of NaOH means 10 g NaOH dissolved in 90 g of water (solvent). % w/w =

msolute msolution

Note: msolution = msolute + msolvent

Example: A sample of 0.892 g of potassium chloride, KCl is dissolved in 54.3 g of water. What is the per cent by mass of KCl in this solution?

D) Weight/Volume per Cent The ratio between the mass of the solute and the volume of the solution. Symbol: % w/v Unit: % g/mL 5% w/v of KCl means that 5 g of KCl is dissolved in 100 mL of KCl (aq) solution. % w/v =

msolute Vsolution

Example: What mass of NaCl is needed to prepare 250 mL of 0.9% w/v solution

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E) Mole Fraction The ratio between the numbers of mole of a component compared to the total number of moles of every component found in the solution. Symbol: xA =

xA (mole fraction for A)

Unit: none

nA = number of mol of component A in a mixture

nA nT

of two or more component nT = total number of moles of all the components (including A) in a mixture

Example: What is the mole fraction of CuCl2 in a solution prepared by dissolving 0.30 mol of CuCl2 in 40.0 mol of H2O.

F) Parts per Million Percentages (%) are parts per hundred. Parts per million (ppm) is quantity of component (in grams) in 106 g of the mixture. Symbol:

Cppm

Unit:

ppm @ mg/L @ g/g @ g/mL @ mg/kg

Cppm =

mass of solute (g) Volume of solution (g)

X 106

Example: The concentration of calcium ions in blood is 100.0 ppm. Calculate the mass of calcium ions in 500.0 g of blood.

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1.4 Chemical Formulae A diatomic molecule contains only two atoms: O2, Br2, HCl, CO A polyatomic molecule contains more than two atoms: O3, H2O, NH3 An ion is an atom, or group of atoms, that has a net positive or negative charge CATION – ion with a positive charge. If a neutral atom loses one or more electrons it becomes a cation ANION – ion with a negative charge. If a neutral atom gains one or more electrons it becomes an anion. 27

How many protons and electrons are in 13

Al

3+

?

protons = 13 electrons = 13 – 3 = 10 A monatomic ion contains only one atom: Na+, Cl-, Ca2+, O2A polyatomic ion contains more than one atom: OH-, CN-, NH4+, NO3Ionic compounds consist of a combination of cations and anions. • the formula is always the same as the empirical formula • the sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero 2 x +3 = +6 3+

Al

3 x -2 = -6

Al2O3

O2-

1.4.1 Chemical Nomenclature Ionic Compounds • often a metal + nonmetal • anion (nonmetal), add suffix “ide” to element name Mg(OH)2 – magnesium hydroxide BaCl – barium chloride Pb(CN)2

K2O – potassium oxide 14

– lead cyanide © Lau Kah Pew STPM 2006


Transition metal ionic compounds • indicate charge on metal with Roman numerals

FeCl2 2 Cl─  -2, so Fe is +2

Iron(II) chloride

FeCl3 3 Cl─  -3, so Fe is +3

Iron(III) chloride

Cr2S3

6

3 S2-  -6, so Cr is — = +3 Chromium(III) sulfide 2

Note that the Roman numeral is written in parenthesis and is joined to the name of the metal. Molecular Compounds • nonmetals or nonmetals + metalloids • The name of the first element in the formula is the same as in the periodic table. But the last element ends in the suffix “ide”. HCl – hydrogen chloride HBr – hydrogen bromide SiC – silicon carbide • If there are more than one combination of compounds from the same elements, the prefix “mono”, “di”, “tri” etc to denote number of atoms in the molecule. NO – nitrogen monoxide NO2 – nitrogen dioxide N2O4 – dinitrogen tetroxide • Some compounds are more commonly known by their traditional names (which may be non-systematic), or for organic compounds, they have their own systematic nomenclature. H2O – water

CH4

NH3 – ammonia

C2H4 – ethene

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– methane



Acids Acids is usually named according to its physical state Acid

Molecular Form

Aqueous Form

HCl

hydrogen chloride

hydrochloric acid

HBr

hydrogen bromide

hydrobromic acid

HI

hydrogen iodide

hydroiodic acid

HCN

hydrogen cyanide

hydrocyanic acid

H2S

hydrogen sulphide

hydrosulphuric acid

Oxoacids: are acids that contain H, O & another element. • The formulas are normally written with the H first, and the O last. Acid

Name

HNO3

Nitric acid

H2CO3

Carbonic acid

H2SO4

Sulphuric acid

HClO3

Chloric acid

• Some oxoacids have the same central element but contain different number H & O. • Naming oxoacids… “more oxygen”

“per” + Name of element + “ic”

HClO4

perchloric acid

“normal oxoacid” e.g. HClO3

Name of element + “ic”

“less 1 oxygen”

Name of element + “ous”

HClO2

chlorous acid

“least oxygen”

“hypo” + Name of element + “ous”

HClO

hypochlorous acid

chloric acid

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Likewise, this system applies similarly to oxoanions… “more oxygen”

“per” + Name of anion + “ate”

NaClO4

sodium perchlorate

“normal oxoacid” e.g. NaClO3

Name of anion + “ate”

“less 1 oxygen”

Name of anion + “ite”

NaClO2

sodium chlorite

“least oxygen”

“hypo” + Name of anion + “ite”

NaClO

sodium hypochlorite

sodium chlorate

Therefore... H2SO4 – hydrogen sulphate

KNO3

– potassium nitrate

H2SO3 – hydrohen sulphite (less one O)

KNO2

– potassium nitrite (less one O)

Bases Bases contain hydroxide ion or yields hydroxide ions when dissolved in water. KOH

– potassium hydroxide

NaOH

– sodium hydroxide

Ba(OH)2

– barium hydroxide

1.4.2 Empirical & Molecular Formula Percent Composition PERCENT n X molar mass of element COMPOSITION X 100% = of an element in a molar mass of compound compound 17



Molecular Formula & Empirical Formula A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance. An empirical formula shows the simplest whole-number ratio of the atoms in a substance Compounds with different molecular formulae can have the same empirical formula, and such substances will have the same percentage composition We can use percent composition to determine the empirical formula of a compound Example Determine the empirical formula for ethanol. If the combustion of 11.5 g ethanol is found to produce 22.0 g CO2 and 13.5 g H2O.

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1.5 Stoichiometry Stoichiometry is the calculation of the quantities of reactants and products involved in a chemical reaction. A + 2B  C + 3D Based upon a balanced chemical equation, we know the mol ratio of reactants and products, thus we are able to determine the quantity of any one of these if the required amount of information is available. Example Methanol burns in air according to the equation 2 CH3OH + 3 O2  2 CO2 + 4 H2O If 209 g of methanol are used up in the combustion, what mass of water is produced?

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1.5.1 Limiting Reagents Normally the amount of reactants that are put together for a reaction is not in the exact proportion as stated in the equation. The limiting reagent is ENTIRELY CONSUMED and limits the amount of products that can be formed because of its INSUFFICIENT QUANTITY The excess reagent does not take part in reaction because it is the balance of what is left AFTER the reaction has consumed the other (limited) reactant(s).

All 6 used 6 balance (leftover) Example In one process, 124 g of Al are reacted with 601 g of Fe2O3 2 Al + Fe2O3  Al2O3 + 2 Fe Calculate the mass of Al2O3 formed.

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1.5.2 Reaction Yield But sometimes in a chemical reaction, we find ourselves in a situation whereby we DON’T GET what we are SUPPOSED TO GET… Theoretical Yield – is the amount of product that would result if all the limiting reagent reacted. Actual Yield

– is the amount of product actually obtained from a reaction.

% Yield =

Actual Yield x 100 Theoretical Yield

Example Cu + 2 AgNO3  Cu(NO3)2 + 2 Ag When 10.0 g of copper was reacted with excess silver nitrate solution, 30.0 g of silver was obtained. a) What is the maximum number of grams of Ag that could have been obtained? b) What was the actual yield of Ag in grams? c) Calculate the percentage yield for this reaction.

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1.5.3 Solution Stoichiometry Example 1 What mass of KI is required to make 500 mL of a 2.80 M KI solution?

Example 2 A particular analytical chemistry procedure requires 0.0500 M K2CrO4. What volume of 0.250 M K2CrO4 must be diluted with water to prepare 100 mL of 0.0500 M K2CrO4?

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In a titration, a solution of accurately known concentration is gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete. During titration, the point whereby the reaction is completed is called the equivalence point The INDICATOR is the substance that changes color at (or near) the equivalence point. It functions as a marker to indicate when the reaction is completed. Example 3 A 25.00 mL sample of HCl solution is titrated against Na2CO3 solution of 0.150 M. It requires 21.20 mL of Na2CO3 for complete neutralisation. Calculate the concentration of HCl solution.

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© Lau Kah Pew STPM 2006 EXERCISE 1. Copper exist naturally on Earth as 63Cu and 65Cu with an isotopic ratio of 2.333 respectively. Based on the carbon-12 scale, the relative isotopic mass of 63Cu = 62.9396 and 65Cu = 64.9278. Determine the relative atomic mass of copper. [63.54]

2. Natural lithium is: 7.42% 6Li (6.015 amu) and 92.58% 7Li (7.016 amu). Calculate the relative atomic mass of lithium metal. [6.941]

3. How many H atoms are in 72.5 g of C3H8O? [5.82 x 1024 atoms H]

4. A sample of iron, Fe, weighs 1.00 kg. What is the amount (mole) of Fe? [17.9 mol]

5. Calculate the number of atoms in 0.20 mol of magnesium. [1.2 X 1023 atoms]

6. How many moles are there in 6.5 L oxygen at STP? [0.29 mol]

7. How much space is needed to fill 3.2 moles of methane gas at room temperature? [76.8]

8. An antacid tablet contains 450 mg Na2CO3. When swallowed, the Na2CO3 reacts with gastric secretion which contains hydrochloric acid (HCl), according to the reaction equation, Na2CO3 + 2 HCl  2 NaCl + CO2 + H2O How many grams of HCl were neutralized by the tablet? [0.031g]

9. Calculate the mass of (NH4)2CO3 that contains a) 0.300 mol NH4+ b) 6.02 x 1023 H atoms 10. How many moles of water can fill a half litre bottle? (the density of water is 1.00 g/mL) 11. Determine the density of oxygen and helium gas at STP. (nHe = 4.003 g/mol; nO = 16.00 g/mol) 12. Seawater is typically 3.5% sea salt and has a density of about 1.03 g/mL. How many grams of sea salt would be needed to prepare enough seawater solution to completely fill a 62.5 L aquarium?

[2.25 X 103 g]

13. An experiment calls for a 0.150 m solution of sodium chloride in water. How many grams of NaCl would have to be dissolved in 500.0 g of water to prepare a solution of this molality? [4.38 g]

14. What is the molality of 10.0% (w/w) aqueous NaCl? [1.90 m]

15. Ascorbic acid (vitamin C) cures scurvy and may help prevent the common cold. It is composed of 40.92% carbon, 4.58% hydrogen and 54.50% oxygen by mass. The molar mass of ascorbic acid is 176 g mol1. Determine its empirical formula and molecular formula. [C6H8O6]

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© Lau Kah Pew STPM 2006 16. A compound Y with chemical formula as shown below: CH2=CHCOOCH3 Write the empirical formula and molecular formula of the compound. Calculate the percentage composition of carbon in the compound Y. 17. Urea, (NH2)2CO, is prepared by reacting ammonia with carbon dioxide. 2 NH3(g) + CO2(g)  (NH2)2CO(aq) + H2O(l) In one process, 637.2 g of ammonia are allowed to react with 1142 g of CO2. a) Which of the two reactants is the limiting reagent? b) Calculate the mass of (NH2)2CO formed and, c) Determine the amount of excess reagent (in grams) that is left at the end of the reaction. [NH3 = limiting reagent, 1124 g (NH2)2CO, 319 g CO2]

18. Titanium is a strong, lightweight, corrosion-resistant metal that is used in rockets, aircraft, bicycle frames, and even sports cars. It is prepared by the reaction of titanium(IV) chloride with molten magnesium between 950o to 1150oC. TiCl4 (g) + 2 Mg (l)  Ti (s) + 2 MgCl2(l) In a certain industrial operation, 3.54 X 107 g of TiCl4 are reacted with 1.13 X 107 g of Mg. Calculate the theoretical yield of Ti in grams. If 7.91 X 106 g of Ti is actually obtained, determine the percent yield. [8.93 X 106 g Ti, 88.6 %]

19. What mass of KI is required to make 500. mL of a 2.80 M KI solution? 20. A 16.42 mL volume of 0.1327 M KMnO4 solution is needed to oxidise 20.00 mL of a FeSO4 solution in an acidic medium. What is the concentration of the FeSO4 ? A 16.42 mL volume of 0.1327 M KMnO4 solution is needed to oxidise 20.00 mL of a FeSO4 solution in an acidic medium. What is the concentration of the FeSO4? [0.5450 M]

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STPM Past Year Questions Q3-P2-2003 b) Kaolinite, Al2Si2O5(OH)4, is a hydrated aluminosilicate mineral clay which is able to absorb cations in aqueous solutions on its surface according to the following equilibrium equation: M+ + Al2Si2O5(OH)3–OH Al2Si2O5(OH)3–O–M + H+ (i) state the oxidation number of silicon in kaolinite.

Q9-P2-2000 a) (i) Define relative atomic mass. (ii) The relative atomic mass of X is 30.97. How many times is one atom of X heavier than one atom of carbon-12? (iii) Naturally occurring sulphur consists of four isotopes with relative abundance shown in the table below. Relative Isotopic Mass

Relative Abundance (%)

31.97

95.02

32.97

0.75

33.96

4.21

35.96

0.02

Calculate the relative atomic mass of sulphur b) The mass spectrum of 1,2-dichloropropane shows peaks at mass/charge ratio of 112, 114, and 116. (i) If 1,2-dichloropropane consists of isotopes: hydrogen-1, carbon-12, chlorine-35 and chlorine-37, give the formulae of the ions responsible for the peaks. (ii) If the ratio of chlorine-35 to chlorine-37 is 3:1, determine the relative abundance of the three peaks. (iii) Sketch and label the mass spectrum of 1,2-dichloropropane showing the peaks concerned.

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Q10-P2-2002 a) An organic compound, P, with a relative molecular mass of 126.5, has the following composition by mass: carbon, 66.4%; chlorine, 28.1%; and hydrogen, 5.5%. Determine the molecular formula of P.

Q7-P2-1999 a) A hydrocarbon with an empirical formula C7H6, and a relative molecular mass of 180 contains 93.33% carbon, and 6.67% hydrogen by mass. Determine the molecular formula of this hydrocarbon.

Q5-P1-Nov 1973 Figure 2 shows the distances of ions in the mass spectrometer of bromine gas. The bromine used consists of its isotopes with a nucleon number of 79 and 81. The atomic mass of bromine is 79.9. The three groups of lines; A, B and C, is produced by ions Br+(g), Br2+(g) and Br2+(g). State which of the ions give the following lines. (a) Group A (b) Group B (c) Group C

1 2 A

1 2 B

1 2 C

Identify all the lines in groups B and C. Underline the spectrum which has the highest abundance.

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Matriculation Past Year Questions Jan 1999 Two common isotopes of chlorine are 35Cl and 37Cl. The relative abundance of natural occurring isotopes is as follows: 35 Cl = 3.127 37 Cl a) Using chlorine as an example, explain the meaning of isotope. b) With reference to 12C = 12.00 scale, the relative isotopic mass of 35Cl and 37Cl are 34.9689 and 36.9659 respectively. Calculate the relative atomic mass of chlorine.

June 1999 Air contains 21% oxygen gas. Natural oxygen consists of three isotopes; 16O, 17O, and 18O. a) Write all possible molecular formulae of oxygen that might exist in our air. b) A mixture of these isotopes was analysed in a mass spectrometer. A stream of univalent positive ions produced is deflected by the magnetic field and is detected by the ion detector. Which ion would deflect least? Explain. c) Calculate the molar mass of the heaviest oxygen gas.

Mar 2002 Analysis of mass spectrometer shows that copper consists of two naturally occurring isotopes: 63Cu and 65Cu. If the ratio of relative abundance of these two isotopes is 63 Cu = 2.235 65 Cu Calculate the percentage of relative abundance of each copper isotope.

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Introduction To Matter: 1.1 Atoms & Molecules 1.1.1 Classification of Matter

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