1 Matter & the Atomic Structure Modul

Chapter 2 : The Structure of the Atom

A Matter

1. Matter is anything that occupies space and has mass. Matter exists in three states – solid, liquid and gas.

2. Matter is made up of tiny and discrete particles.

3. An atom is the smallest particle of an element that can participate in a chemical reaction.

4. An ion is a positively-charged or negatively-charged particle.

5. Particles in matter are in motion.

6. Diffusion occurs when particles of a substance move in between the particles of another subtance.

7. Diffusion of matter occurs most rapidly in gases, slower in liquids and slowest in solids, due to the different arrangement and movement of particles in the three states of matter.

The kinetic theory of matter.(diagram)

i) The change in heat changes the state of matter.ii) When a substance is heated, the particles gain kinetic energy and move faster.iii) When a substance is cooled, the particles loss their kinetic energy and move slower.

Do it yourself 2.1

State of MatterCharacteristic

Solid liquid gas

Arrangement of particles are packed closely together in an orderly manner

are packed closely together but not in orderly manner

Very far apart and in a random motion

Forces of attraction Strong forces Strong forces but weaker than the forces in a solid

Weak forces

Movement of particles Particles vibrate and rotate about their fixed positon

Particles vibrate, rotate and move throughout the liquid. They collide against each other

Particles vibrate, rotate and move freely. The rate of collision is greater than in liquid

Shape / volume - has a fixed volume and shape

- has a fixed volume and follows the shape of the container

Does not have a fixed shape of volume

compressibility Cannot be compressed Cannot be compressed easily

Can be compressed easily

1

1.Figure 1 below shows the heating curve of a pure substance at room temperature and pressure.

Temperature/oC U

S T

119 Q R

P

Time / minute

Figure 1

(a)Describe the movement of the particles of the pure substance at stage PQ of the curve ?

(b)Draw a diagram to show the arrangement of particles of the substance at stage QR in the box below.

(c)Explain why the pure substance is not water ?

(d)Samples of the pure substance at stage RS and TU are taken. Compare the movement of the particles of the substance at these two stages.

(e)After heating at 500oC, the substance is cooled. Draw and label the cooling curve.

2. P Q

2

U T S R

(a)Name the process in

P: Q :

R: S :

T: U :

(b)What will occur if matter undergoes a change of state ?.

(c)Compare the intermolecular distance and the packing of particles in the solid state and the liquid state.

B The Atomic Structure

3

Ice water Salt solution

water steamsalt water

Apply heat Apply

heat

Saturate it thencool the solution

Boil it, then cool the vapour

The historical development of atomic models.

Scientist Atomic Models

1. John Dalton

- imagined the atom as a small indivisible ball similar to a very tiny ball

2.J.J. Thomson

- described the atom as a sphere of positive charge which contains a few negatively-charged particles called electrons.

3. Ernest Rutherford

discovered proton the positive charge and most of the mass of the atom are concentrated in a small, central region called the nuclueselectrons move in a space that is larger than the space occupied by the nucleus

4.Neils Bohr

proposed that the electrons in an atom move in shells around the nucleus

5.James Chadwick

proved the existence of neutrons, the neutral particles in the nucleus. Neutrons contribute approximately to half the mass of an atom.

4

Protons, neutrons and electrons are subatomic particles of an atom.

i) Atoms are electrically neutral.ii) The number of protons is equal to the number of electrons.iii) The proton number of an element is the number of protons in its atom.iv) The nucleon number of an element is the total number of protons and neutrons in its atom.

Therefore,

v) Each element has its own proton number.vi) Each element is given a name and a symbol

Proton number

Element symbol Proton number

Element symbol

1 Hydrogen H 11 Sodium Na2 Helium He 12 Magnesium Mg3 Lithium Li 13 Aluminium Al4 Beryllium Be 14 Silicon Si5 Boron B 15 Phosphorus P6 Carbon C 16 Sulphur S7 Nitrogen N 17 Chlorine Cl8 Oxygen O 18 Argon Ar9 Flourine F 19 Potassium K10 Neon Ne 20 Calcium Ca

An atom of an element can be written as A X Z Where A is the nucleon number, X is the symbol of an element, Z is the proton number.

Subatomic particle

Symbol Relative mass

Charge Location

Proton p 1 +1 In the nucleus

Electron e 1/1840 -1In orbits around

the nucleus

Neutron n 1 0 In the nucleus

5

Nucleon number = proton number + number of neutrons

Do it yourself 2.2

Complete the table below.

Complete the table below

Element (symbol)

number of protons Number of neutrons Symbol of atoms

Lithium (Li)

Neon (Ne)Zinc (Zn)

Symbol of atom 27 Al 13

19 F 7

23 Na 11

Proton number

Nucleon number

Number of protons

Number of electrons

Number of neutrons

6

2.3 Isotopes and Their Importance

1. The isotopes of an element are the atoms of that element which contain a same

number of protons, but a different number of neutrons.

2. Isotopes of some element

Element IsotopesHydrogen

1 proton0 neutron

1 proton 1 neutron

1 proton 2 neutrons

Carbon

6 protons 6 neutrons

6 protons7 neutrons

6 protons8 neutrons

Oxygen

8 protons 8 neutrons

8 protons9 neutrons

8 protons10 neutrons

Sulphur



-

Bromine

35 protons44neutrons


-

3. The uses of isotopes in daily lifeField Isotopes applications

Medical Gamma rays from cobalt- 60 are used to kill cancer cell without surgery in patient. This treatment is known as radiotherapy.

Medical instrument are sterilized using gamma rays.

Radioactive materials such as iodide-131 are injected into patients to detect malfunction of thyroid glands.

Archeology Radioisotope carbon -14 is used to study the age of ancient artifacts.

Agricultural Carbon -14 is used to study the passage of

carbon in photosynthesis of green plants.

Industrial Isotope sodium-24 is used to detect leakage of underground pipes.

7

SPM

4. The electron arrangement of elements with proton number 1 to 20.( must know how to memorize)

2.4 Electron Arrangements / Electron Structures

Element Number of neutrons

Number of protons

Number of electrons

Number of nucleon

Electron arrangement

Number of valence

electronsHydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAluminumSiliconPhosphorusSulphurChlorineArgonPotassiumCalcium

First shell: 2 electrons

Second shell: 8 electrons

Third shell: 8 electrons

Nucleus ( contains protons and nucleus)

Last electron/s in the last outermost shell, we called as valence electron.

8

Chapter 3: Chemical Formulae and Equations

Subtitle 3.1: Relative Atomic Mass and Relative Molecular Mass

Concept:We can determine the mass of an atom relative to a standard atom

A Helium atom is 4 times heavier compare to a hydrogen atom.Helium is said to have relative atomic mass of 4

Important !!!Define:

Define

How to measure mass of an atom?

Standard atom1. hydrogen 2. oxygen

3. carbon-12✔ Solid & easy to

handle✔ Also used as

standard for mass spectrometer

Check to:page 176 of text book Look at Ar of all elements listed in

periodic table

from periodic table: Ar of Nitrogen atom is 14.The average mass of a nitrogen atom is 14 times larger than 1/12 of a carbon-12 atom.

Mr of Water Molecule is 18The average mass of one water molecule is 18 times larger than 1/12of a carbon-12 atom

9

Hydrogen as standard atom

✄ not use any more because gasseous form are difficult to handle

Relative atomic mass, A r

- of an element is the average mass of one atom of the element when compared with 1/12 of the mass of an atom of carbon-12

1/12 of one atom carbon-12

How many helium atoms are here?????

Relative atomic mass of an element= The average mass of one atom of an element 1/12 x the mass of an atom of carbon -12

Relative molecular mass of an element= The average mass of one molecule 1/12 x the mass of an atom of carbon -12

Relative molecular mass, Mr-of a molecule is the average mass of the molecule when compared with 1/12 of the mass of an atom of carbon-12

helium atom

Important!!!!Relative mass does not have any unit.

Numerical problemsA. About Relative Atomic Mass

1. How many times is copper atom heavier than two helium atom? Solution: Mass of a copper atom = Ar of copperMass of 2 helium atom 2 x Ar of helium = 64 2 x 4 = 8 times

2. How many magnesium atom have the same mass as two silver atoms ? Solution:

Lets the number of magnesium atoms = n Mass of n magnesium atoms = mass of 2 silver atoms So, n x Ar of magnesium = 2 x __________ n x 24 =

=

Do It Yourself

1. How many times is one atom of silicon heavier than one atom of lithium

2. Calculate the number of atoms of lithium that have the same mass as two atoms of nitrogen

3. The mass of one atom Y is A times larger than the mass of one nitrogen .Calculate the relative atomic mass of Y.

Get Ar value from periodic table

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Form 4 text bookQuick reviewpage 30

B. About Relative Molecular Mass

To determine Relative Molecular Mass, Mr

Molecular substance

Relative Molecular Mass

Carbon dioxide, CO2

Ar of C + ( 2 x Ar of O) = 12 + (2 x 16 ) = 44

Nitrogen gas, N2 2 x Ar of N = 2 x 14 = 28

Relative formula mass is used to replace Mr for ionic substances

Ionic substance Relative formula mass

Sodium Hydroxide, NaOH Ar of Na + Ar of O + Ar of H= 23 + 16 + 1 = 40

Aluminium sulphate, Al2 (SO4)3

2 x Ar of Al +3 ( Ar of S + 4 x Ar of O )=

Hydrated copper(II) sulphate, CuSO4. 5H2O

Ar of Cu + Ar of S + 4 x Ar of O + 5 ( Mr of H2O)=

Do it yourself

1. Calculate the relative molecular mass of

a) Bromine, Br2

c) Ammonia, NH3

b) Methane, CH4 d) Glucose, C6H12O6

Get Ar value from periodic table

11

2. Calculate relative formula mass of

a) Zinc oxide, ZnO c) Copper(II) hydroxide, Cu(OH)2

b) Magnesium nitrate, Mg(NO)3 d) Hydrated sodium carbonate, Na2CO3.10H2O

B. The Mole and the Number of Particles

Definition of mole

The word ‘pair’ and ‘dozen’ represent a fixed number of objects.

In chemistry, we use the unit ‘mole’ to measure the amount of substance. The symbol of mole is mol. 1 mol of substance = the number of particles in 12 g of carbon-12.

= 6.02 x 1023 particles.

The value of 6.02 x 1023 is called as the Avogadro constant (NA). To determine the number of moles or the number of particles:

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Form 4 practical bookTry this 3.1page 17

Number of particles = Number of moles 6.02 x 1023 Practical Book Activity 3.2, page 17

Example 1:i. 1 mol of iron atom = 6.02 x 1023 iron atomsii. 1 mol of hydrogen molecule = 6.02 x 1023 hydrogen moleculesiii. 1 mol of sodium chloride = 6.02 x 1023 formula units of sodium chloride

Example 2:A closed glass bottle contains 0.5 mol of oxygen gas, O2.i. How many oxygen molecules, O2 are there in the bottle?ii. How many oxygen atoms are there in the bottle?

Solution:i. Number of oxygen molecules = 0.5 x 6.02 x 1023

= 3.01 x 1023

ii. 1 oxygen molecule, O2 has 2 oxygen atoms.Therefore, number of oxygen atoms = number of oxygen molecules 2= 3.01 1023 2= 6.02 1023

Example 3:Find the number of moles of molecules in a sample containing 9.03 × 1023 molecules of carbon dioxide, CO2.

Solution:

Number of moles =

= 1.5 mol.

Do it yourself

[Avogadro constant = mol-1]

1 Define a mole?

A mole is the amount of substance which has the same number of particles as there in 12 g carbon -12.

2 Calculate the number of atoms in 2 mol carbon.

Number of atoms = 2 × 6.02 x 1023 = 1.2 ×1024 atoms.

3 How many ions are there in 1.5 mol sodium chloride, NaCl?

1 formula unit sodium chloride, NaCl has 2 ions which are 1 sodium ion and 1 chloride ion.Thus, number of ions = number of formula units x 2

= 1.5 × 6.02 x 1023 × 2

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= 1.806 × 1024 ions.4 Calculate the number of moles of bromine molecules which consists 1.5 × 1022 of bromine molecules.

Number of moles = 1.5 ×10 22 6.02 × 1023

= 0.025 mol.

5 How many atoms are there in 2 mol of ammonia, NH3?

1 ammonia molecule, NH3 has 4 atoms which are 1 nitrogen atom and 3 hydrogen atoms.Thus, number of atoms = number of molecules x 4

= 2 × 6.02 x 1023 × 4 = 4.8 × 1024 atoms.

C. The mole and the mass of substances

Molar mass is

Unit of molar mass is g mol-1 or grams per mole.

The molar mass of a substance = the mass of 1 mol of the substance= the mass of NA number of particles= the mass of 6.02 x 1023 particles

Example:

Element/ Compound Relative mass Mass of 1 mol

Molar mass

Lithium, Li 7 7g 7g mol-1

Iron, Fe 56 56gMagnesium oxide, MgO 24+16=40 40g mol-1

Carbon dioxide, CO2 12+16x2=44

*1 : The value of molar mass of an element is equal to its relative atomic mass*2 : The value of molar mass of a compound is equal to its relative molecular or formula mass

Formula: Number of moles = mass

Relative atomic mass

(or relative molecular mass or relative formula mass)

14

*1

*2

Example:

1. Calculate the number of moles found in 20g of magnesium oxide, MgO. (Relative atomic mass: Mg, 24; O, 16) Solution:

Number of moles = mass Relative formula mass = 20 24 + 16 = 0.5 mol

2. Calculate the mass in gram found in 0.2 mol of magnesium oxide, MgO. (Relative atomic mass: Mg, 24; O, 16) Solution:

Number of moles = mass

Relative formula mass

Mass = number of moles x relative formula mass = 0.2 x (24 + 16)g

= 8g

3. How many magnesium ions are there in 30g of magnesium oxide, MgO. (Relative atomic mass: Mg, 24; O, 16. Avogradro constant: 6.02 x 1023) Solution:

The relative formula mass of magnesium oxide, MgO = 24 + 16 = 40

Therefore, the molar mass of magnesium oxide, MgO = 40g mol-1

Number of moles of 30g magnesium oxide, MgO = mass of MgO

Relative formula mass of MgO = 30g

40 g mol-1 = 0.75 mol

The number of formula units of MgO = 0.75x 6.02 x 1023

= 4.515 x 1023

Each formula units of MgO has 1 magnesium ions.

Therefore, the number of magnesium ions = the number of formula units of MgO x 1 = 4.515 x 1023 x 1 = 4.515 x 1023

4. Calculate the mass in gram of 3 x 1022 units of magnesium oxide, MgO. (Relative atomic mass: Mg, 24; O, 16. Avogradro constant: 6.02 x 1023)

15

Solution:

Number of moles = number of particles NA

Mass = number of particlesRelative formula mass NA

Mass = number of particles x relative formula mass NA

Mass of 3x1022 units of magnesium oxide, MgO= 3 x 10 22 x (24+16) 6 X 1023

= 0.05 X 40 = 2 g

Do It Yourself

1. Calculate the number of moles found in 9.5g of magnesium chloride, MgCl2. (Relative atomic mass: Mg, 24; Cl, 35.5)

2. Calculate the mass in gram found in 0.3 mol of magnesium chloride, MgCl2.

(Relative atomic mass: Mg, 24; Cl, 35.5)

3. How many chloride ions are there in 19g of magnesium chloride, MgCl2. (Relative atomic mass: Mg, 24; Cl, 35.5. Avogradro constant: 6.02 x 1023)

4. Calculate the mass in gram of 3 x 1022 units of magnesium chloride, MgCl2 . (Relative atomic mass: Mg, 24; Cl, 35.5. Avogradro constant: 6.02 x 1023)

E. Chemical Formulae

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Form 4 TextBookWork This Out 3.2Page 35

Quick Review CPage 35

A chemical formula is a representation of a chemical substance using letters for atoms and subscript numbers to show the numbers of each type of atoms that are present in the substance. Examples : (a) Glucose

C6 H12O6

(b) Sodium hydroxide

Mg (OH)2

(1) Empirical Formulae

(i) The empirical formula of a compound gives the simplest whole number ratio of atoms of each element present in the compound.

(ii) Steps in determining the empirical formula of a compound.i. find the mass of each element in the compoundii. convert the masses to the numbers of moles of atomsiii. find the simplest ratio of moles of the elements

Example : 2.24 g of iron combines chemically with 0.96g of oxygen to form an oxide. What is the empirical formula of the oxide ?

[ Relative atomic mass: O, 16; Fe, 56 ]

Element Iron, Fe Oxygen, O

Mass (g) 2.24 0.96

Number of moles of atoms 2.24 = 0.0456

0.96 = 0.0616

Ratio of moles 0.04 =10.04

0.06 =1.50.04

Simplest ratio of moles 1 × 2 = 2 1.5 × 2 = 3

The empirical formula of the oxide is Fe2O3.

Do it Yourself

1. The table below shows the relative atomic mass and the mass of elements V and O in an oxide.Element V O

17

Show the symbols for carbon, hydrogen and oxygen

Show the numbers of carbon, hydrogen and oxygen

Show the symbols for magnesium, oxygen and hydrogen.

Show the numbers of magnesium, oxygen

and hydrogen.

Relative Atomic Mass 56 16Mass(g) 5.6 2.4

What is the empirical formula of this compound ?

element V Oxygen, OMass (g) 5.6 2.4Number of moles of atomsRatio of moles

Simplest ratio of moles

The empirical formula of the oxide is ……………………

2. Copper (II) iodide constains 20.13% of copper by mass. Find its empirical formula. [ Relative atomic mass : Cu,64 ; I, 127 ]

Based on its percentage composition, 100g of copper(II) iodine contains 20.13g of copper. So, taking 100g of the compound.

element K ClMass (g)Number of moles of atomsRatio of moles


The empirical formula of the oxide is ………………….

3. A potassium compound has a percentage composition as the following K, 31.84% ; Cl, 28.8% ; O, 39.18% What is the empirical formula of the potassium compound ? [ Relative atomic mass : O, 16; Cl,35.5; K, 39 ]

Based on its percentage composition, 100g of compound contains 31.84g of potassium, 28.98g of chlorine and 39.18g of oxygen. So, by taking 100g of the compound:

element K Cl OMass (g)Number of moles of atomsRatio of moles


1 mole of potassium atoms combines with 1 mole of chlorine atoms and 3 moles of oxygen atoms.Therefore, the empirical formula of the potassium compound is KClO3.

(2) Molecular Formulae

(i) The molecular formula of a compound gives the actual number of atoms of each element present in a molecule of the compound.

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Form 4 TextBook Work this out 3.7Page 42

(ii) The molecular formula of a compound is a multiple of its empirical formula.

(iii) Relating empirical formula to molecular formula Compound Empirical formula Molecular formula n

Water H2O H2O = (H2O)1 1Ethene CH2 C2H4 = (CH2)2 2Ethane CH3 C2H6 = (CH3)2 2propane CH2 C3H9 = (CH3)3 3glucose CH2O C6H12O6 = (CH2O)2 6

(iv) Calculation involving molecular formulae

Example :The empirical formula of a compound is CH. Its relative molecular mass is 78. Find its molecular formula. [ Relative atomic mass : H, 1; C, 12 ]

Let the molecular formula be (CH)n.

The relative molecular mass = n[ 12 + 1 ] = 13n

However, its molar mass is 78. Therefore, 13n = 78 n = 78/13 = 6

Hence, the molecular formula of the compound is (CH)6 or C6H6.

Do it yourself

1. A carbon compound has an empirical formula of CH2 and a relative molecular mass of 70. Find the molecular formula of the compound. [ Relative atomic mass : H, 1; C, 12 ]

Hence, the molecular formula of the compound is (CH2)5 or C5H10.2. 2.07 g of element Z reacts with bromine to form 3.67g of a compound with an

empirical formula of ZBr2. Find the relative atomic mass of element Z. [ Relative atomic mass: Br, 80 ]

19

Molecular formula = ( Empirical formula )n

element Z BrMass (g)

Number of moles of atoms

Simplest ratio of moles (from the emp for given)

Based on the empirical formula ZBr2 , the ratio of atoms of Z : Br is 1 : 2 herefore,

2.07 : 0.02 = 1 : 2 z 2.07/0.02z = ½ z = 207

The atomic mass of the element Z is 207.

(3) Ionic Formulae

(i) Ionic compounds are compounds consisting of anions and cations.

(ii) The formulae of some common cations

Cation ( positive ion ) Formula of cation Charge of cationSodium ion Na+ +1Potassium ion K+ +1Silver ion Ag+ +1Hydrogen ion H+ +1Ammonium ion NH4

+ +1Copper (II) ion Cu2+ +2Calcium ion Ca2+ +2Magnesium ion Mg2+ +2Zinc ion Zn2+ +2Barium ion Ba2+ +2Iron (II) ion Fe2+ +2Copper (I) ion Cu+ +1Tin (II) ion Sn2+ +2Lead (II) ion Pb2+ +2Aluminium ion Al3+ +3Iron (III) ion Fe3+ +3Chromium (III) ion Cr3+ +3

(iii) The formulae of some common anions

Anion ( negative ion ) Formula of anion Charge of anion

20

FORM 4 Textbook Work this out 3.8Page 44

Fluoride ion F- -1Chloride ion Cl- -1Bromide ion Br- -1Iodide ion I- -1Hydroxide ion OH- -1Nitrate ion NO3

- -1Nitrite ion NO2

- -1Hydride ion H- -1Oxide ion O2- -2Phosphate ion PO4

3- -3Carbonate ion CO3

2- -2Sulphate ion SO4

2- -2Chromate (VI) ion Cr2O7

2- -2

(iv) The chemical formulae of ionic compounds are electrically neutral because the total of positive charges are equal to the total of negative charges

(v) The chemical formula of an ionic compound can be constructed as the following :i. identify and write down the formula of its cation and anionii. determine the number of cations and anions by balancing the positive and

negative charges.iii. Write the formula of the compoundiv. The number of cations and anions are written as subscript numbers.

MgCl2

Do it yourself

1. magnesium chloride

21

Magnesium chloride

Magnesium ion, Mg2+ Chloride ion, Cl-

1 magnesium ion, Mg2+

Total of positive charges=1 (+2)=+2

2 chloride ions, Cl-

Total of negative charges= 2 (-1)= -2

2. aluminium oxide

3. aluminiuim hydroxide

4. sodium sulphate

(4) Naming of chemical compounds

1. Chemical compounds are named systematically according to the guidelines given by the International Union of Pure and Applied Chemistry (IUPAC).

2. For ionic compounds, the name of the cation comes first, followed by the name of anion.

cation anion Name of ionic compoundSodium ion Chloride ion Sodium chloride

Magnesium ion Oxide ion Megnesium oxideAluminium ion Oxide ion Aluminium oxide

Zinc ion Sulphate ion Zinc sulphate

3. Transition metals can form more than one ions, Roman numerals ( such as I, II, III ) are used to differentiate the ions.

Fe2+ - iron (II) ion Fe3+ - iron (III) ion

4. For simple molecular compounds, the name of the first element is maintained. However, the name of the second element is added with an “ –ide “.

Examples : HCl – hydrogen chloride HF - hydrogen flouride 5. Greek prefixes are used to show the number of atoms of each element in a compound. Examples : CO – carbon monoxide CO2 – carbon dioxide CCl4 – carbon tetrachloride / tetrachloromethane SO3 – sulphur trioxide

6. Table below shows the meaning of the prefixes.

prefix meaning prefix meaningMono- 1 Hexa- 6

di- 2 Hepta- 7Tri- 3 Octa- 8

Tetra- 4 Nona- 9Penta- 5 Deca- 10

F. CHEMICAL EQUATION

A) Qualitative aspect of chemical equation

♣ A chemical equation is a shorthand description of a chemical reaction.

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Form 4 TextbookWork This Out 3.9Page 46


♣ The starting substances are called reactants.♣ The new substances formed are called products.

♣ The reactants are written at the left-hand side of the equation.♣ The products are written at the right-hand side of the equation.

♣ A chemical equation also shows the states of each substance.

Symbol Physical states of substancess Solidℓ Liquidg Gas

aq Aqueous solution

Example :

Do It Yourself 3f

Identify the reactants, products and the state of each substance. Present your answer in the form of a table.

Solution :

Reactants Products123

B) Writing chemical equation

A chemical equation must be balanced. There must always be the same number of atom of each element on each side of the equation.

23

Reactants Products

C (s) + O2 (g) CO2 (g)

Zn (s) + Cl2 (g) ZnCl2 (s)

1. HCl (aq) + NaOH (aq) NaCl (aq) + H2O (i)

2. CuCO3 (s) CuO (s) + CO2 (g)

3. HCl (g) + NH3 (g) NH4Cl (s)


Example :

Magnesium reacts with dilute hydrochloric acid, HCl to produce magnesium chloride, MgCl2 and hydrogen gas, H2. Write an equation to represent the reaction.

Solution :

Do It Yourself 3.f B

Write a chemical equation for each of the following reactions.

1. A solution of silver nitrate is added to a solution of sodium chloride. A precipitate of silver chloride and a solution of sodium nitrate are produced.

2. Nitrogen gas reacts with hydrogen gas to produce ammonia gas.

3. When solid lead (II) carbonate is heated strongly, it decomposes into solid lead (II) oxide and carbon dioxide gas is released.

C) Quantitative aspect of chemical equation

The coefficients in a balanced equation tell us the exact proportions of reactants and products in a chemical reaction.

24

STEP 1 Write the equation in words. The reactants are written on the left whereas the products are written on the right.

STEP 2 Write the correct chemical formula for each reactants and products.

STEP 3 Balance the equation. You just need to adjust the coefficients in front of the chemical formulae and not the subscripts in the formulae.

STEP 4 Put the state symbols in the equation.


STEP 1 Magnesium + hydrochloric acid Magnesium chloride + hydrogen gas

Reactants Products

STEP 2 Mg + HCl MgCl2 + H2

STEP 3 Mg + 2HCl MgCl2 + H2

STEP 4 Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g)

Example :

The equation tell us that 2 moles of hydrogen reacts with 1 mole of oxygen to produce 2 moles of water.

Or

The equation tell us that 2 molecules of hydrogen reacts with 1 molecule of oxygen to produce 2 molecules of water.

D) Numerical problems involving chemical equation

Stoichiometry is a study of quantitative composition of substances involved in chemical reactions.We can always make use of the stoichiometric coefficients in a chemical equation to solve various numerical problems.

Generally the steps involved in stoichiometric calculations are as follows.

Example :

Copper (II) oxide, CuO reacts with aluminium according to the following equation.

Calculate the mass of aluminium required to react completely with 12 g of copper (II) oxide, CuO.[Relative atomic mass : O, 16 ; Al, 27 ; Cu, 64]

Solution :

The number of moles of 12g of Copper (II) oxide, CuO = 12 g (64 + 16) g mol-1

= 12 g = 0.15 mol 80 g mol-1

Based on the chemical equation, 3 mole of Copper (II) oxide, CuO requires 2 mole of aluminium. Therefore, the number of aluminium required by 0.15 mole of Copper (II) oxide, CuO

= 0.15 mole x 2 mole = 2 mole

25

2H2 (g) + O2 (g) 2H2O (l)

STEP 1 Write the balanced equation of the reaction.

STEP 2 Compare the mole ratio.

STEP 3 Identify the information given and you want to find.

STEP 4 Calculate the number of moles.

3CuO (s) + 2Al (s) Al2O3 (s) + 3Cu (s)

3CuO (s) + 2Al (s) Al2O3 (s) + 3Cu (s)

3 mole 2 mole

3 mole

Thus the mass of aluminium required

= 0.1 mol x 29 g mol-1 = 2.7 g

Do It Yourself 3f (D)

How many moles of potassium are needed to reacts with 0.5 mole of bromine gas ?

Solution :

Information : ? mole 0.5 mole

Based on the equation, 1 mole of bromine gas reacts with 2 moles of potassium.

Therefore, 0.5 mole of bromine gas will react with

2 x 0.5 = 1 mole of potassium.

2. 1.35 g of aluminium reacts with excessive copper (II) oxide powder to produce aluminium oxide powder and copper. Find the number of copper atoms produced.[Relative atomic mass : Al, 27 ; Avogadro constant : 6.02 x 1023 mol-1]

3.

What is the mass of zinc needed to produce 2.4 dm3 of hydrogen gas at room conditions ?[Relative atomic mass : Zn, 65 ; Molar volume 24 dm3 mol-1 at room conditions]

26

1. 2K (s) + Br2 (g) 2KBr (s)

Zn (s) + 2HNO3 (aq) Zn(NO3)2 (aq) + H2 (g)

2K (s) + Br2 (g) 2KBr (s)

2 mole 1 mole

More Exercises:

1. CuCO3 CuO + CO2

In this reaction, 3.1 g of copper(II) carbonate are heated in a laboratory. Find :

(a) the mass of copper (II) oxide that being produced.(b) the volume of carbon dioxide gas produced at s.t.p

2. CaCO3 CaO + CO2

In this reaction, 300 cm3 gas carbon dioxide are produced at room temperature, when calcium carbonate are heated. Find:

(a) the mass of calcium carbonate used.(b) mass of calcium oxide produced.

3. 2Na + 2H2O 2NaOH + H2

When 0.23 g of sodium is added to water, the metal will react vigorously at the surface of the water, find(a) the mass sodium hydroxide produced.(b) volume of hydrogen gasses being produced at temperature room.

4. 2Mg + O2 2MgO

A strip of magnesium has a weight of 1.2 g are being burn with sufficient oxygen to produced magnesium oxide. Find:

(a) the mass magnesium oxide being produced.(b) the mass of oxygen that needed for this reaction.

5. C3H8 + 5O2 3CO2 + 4H2O

Propane gas was burned in oxygen follow as equation above. If 3.36 dm3 of carbon dioxide gas are produced in this reaction at s.t.p, find

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(a) the mass of propane burned(b) volume of oxygen gas that reacted

6. 2Al + 3CuO Al2O3 + 3Cu

1.35g of aluminium powder and copper (II) oxide was heated strongly in laboratory to produced aluminium oxide and copper. Find

(a) the mass of copper (II) oxide reacted(b) the mass of aluminium oxide produced.(c) the mass of copper produced.

Chapter 4.

PERIODIC TABLE OF ELEMENTSA. Historical Development of the Periodic Table

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Group I Group II Group III Group IV

Oxygen NitrogenHidrogenLightHeat

SulphurPhosporusCarbon ChlorinFluorin

ArsenicBismutCobaltLeadZincNikelStanumArgentum

Potassium oxideBarium oxideSilicon(IV) oxsideMagnesium oxideAluminium oxide

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Scientist like to find patterns. In the 18th and 19th centuries, scientist discovered many elements. The elements found were classified through many stages of hard work by scientist. This led to the development of the Periodic Table of Elements that we use today.

Here is the history ofPeriodic Table of Elements.

Antoine Lavoisier First chemist who classify the element into 4 group. The 4 group consisted of gases, metal, non-metal and metal

oxide. Element in the group is classify into metal and non-metal.

Element in triadLithium

LiSodium

NaPotassium

K

Average atomic mass

Li and K


7 23 39

7 + 39 = 23 2

Element in triadChlorine

ClBromine

BrIodine I

Average atomic mass Li and K


35 80 12735 + 127 = 2 81

Ca Sr Ba (40 + 137) ÷ 2 = 88 Li Na K Cl Br I 7 23 39 35 80 127

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Johann W. Dobereiner

Classify the element with same chemical properties into a few group Each of group consisted from 3 element called triad. He found that the relative atomic mass of the element in the middle of

each triad is approximately equal to the average atomic mass of other two elements.

Triad law show the relationship between the relative atomic mass of elements with it chemicals properties.

This law cannot be use to most of the other element.

John Newlands Arranged 62 known elements in order of increasing nucleon

number (atomic weights ) in horizontal rows. He noted that after interval of eight elements similar

physical/chemical properties reappeared. He was the first to formulate the concept of periodicity in the

properties of the chemical elements. He proposed the Law of Octaves:

Elements exhibit similar behavior to the eighth element following it in the table.

He was not successful because;i. Was only accurate for the first 16 elements (from

hydrogen to potassium)ii. There were no gaps allocated from the elements yet

to be discovered.

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Lothar Meyer

Determine the volume of an atom of an element. Formula;

Volume of an atom = mass of one mole-atom of the element Density of the element

He plotted a graph of volume of atoms of elements against their relative atomic masses to produce meyer’s atomic volume curve.

From the graph he found elements occupying the corresponding positions of the curve exhibit similar chemical properties. example

(a) Li, Na, K, Rb : located at the peak of the curve

(b) Be, Mg, Ca, Sr : located after the maximum point

Like Newlands, Meyer showed the properties of the elements recured periodically.

Dimitri Mendeleev

Arranged the elements in order of increasing atomic weights and properties.

He left gaps for elements yet to be discovered.He arranged the element that have the same properties in group.

Henry Mosely

He was able to derive the relationship between x-ray frequency and number of protons. and obtained a straight line graph.

When Moseley arranged the elements according to increasing atomic numbers and not atomic masses, some of the inconsistencies associated with Mendeleev's table were eliminated.

The modern periodic table is based on Moseley's Periodic Law (atomic numbers/proton number). He suggest proton number determine the position of elements in periodic table. He arranged elements in periodic table in order of increasing proton number. He also left gaps for the elements yet to be discovered.

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Henry Mosely Periodic Table

Arrangements of elements in the Periodic Table

Elements in the Periodic Table are arranged in an increasing order of proton number.

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What is the basic principle applied in arranging the elements in the Periodic Table today?

Elements with similar chemical properties are placed in the same vertical column. There are 18 vertical column of elements in the Periodic Table. Each column is called group. The

vertical columns are known as Group 1 to Group 18. There are 7 horizontal rows of elements in the Periodic Table. Each of these horizontal rows of

elements is called a period. The horizontal rows are known as Period 1 to Period 7.

1. The number of valence electrons in an atom decides the position of the group of an element.

The number of valence electron

1 2 3+10 4+10 5+10 6+10 7+10 8+10

Group in The Periodic Table

1 2 13 14 15 16 17 18

2. The number of shells occupied with electrons in its atom decides the period number of an element.

Example 1;

Example 2;

X40

20

Number of proton = 20Number of electron = 20Number neutron = 20

Electron arrangement = 2.8.8.2The number of valence electrons = 2The number of shells = 4

Group in the Periodic Table = 2Period in the Periodic Table = 4

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Do you know how the electron arrangement of the atom of an element related to its group and period?????

Example 3;

Hw: WTO 4.3 pg. 62 no. 1,2,3

PERIODIC TABLE OF ELEMENTS

B. GROUP 18 ELEMENTS

http://periodictable.com/

Y16

8


Electron arrangement = 2.6The number of valence electrons = 6The number of shells = 2

Group in the Periodic Table = 16Period in the Periodic Table = 2

Z40

18


Electron arrangement = The number of valence electrons = The number of shells =

Group in the Periodic Table = Period in the Periodic Table =

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GROUP 18

1.The elements in Group 18 are Helium 2Neon 2.8Argon 2.8.8Krypton 2.8.18.8Xenon2.8.18.18.8Radon2.8.18.32.18.8

1. They are also known as noble gases, which are chemically unreactive. Noble gases are monoatomic.

2. Helium has two valence electrons. This is called duplet electron arrangement.

3. Other noble gases have eight valence electrons. This is called octet electron arrangement.

4. Duplet and octet electron arrangements are very stable because the outermost occupied shells are full.

5. All nobles gases are inert which means chemically unreactive.

BECAUSE THE OUTERMOST OCCUPIED SHELLS ARE FULL

Physical Properties of Group 18 Elements1. Group 18 elements have very small atom.

2. They are colourless gases a room temperature and pressure.

3. They have low melting and boiling point.

4. They have low densities.Elements/ symbol


Atomic radius (nm)

Melting points (°C)

Boiling points (°C)

Density(g cm-3)

Helium 2 0.050 -270 -269 0.17Neon 2.8 0.070 -248 -246 0.84Radon 2.8.8 0.094 -189 -186 1.66Krypton 2.8.18.8 0.109 -156 -152 3.45Xenon 2.8.18.18.8 0.130 -112 -107 5.45Radon 2.8.18.32.18.8 - -71 -62 -

Table 1: Physical Properties of Group 1 Elements

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Why noble gases exist as monoatomic gases and are chemically unreactive?

4. From Table 1, when going down the group, atomic size and density increase.5. When going down the group, melting points and boiling points decrease

Uses of Group 18 elements

Helium Used to fill airships and weather balloons, because the gas is very light. The diver’s oxygen tank contains a mixture of helium (80%) and oxygen (20%).

Neon Advertising lights. Television tubes. Airport landing bulb to help aero plane landing safely.

Argon To fill light bulbs, it can last longer To provide inert atmosphere for welding at high temperature.

Krypton Used in lasers to repair the retina of the eye. To fill photographic flash lamps.

Radon Used in treatment of cancer.

Xenon Used in bubble chambers in atomic energy reactors.

Hw: QR B pg. 65 no. 1,2

C. GROUP 1 ELEMENTS

http://periodictable.com/

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GROUP 1

6. The elements in Group 1 are

Lithium 2.1 Sodium 2.8.1Potassium 2.8.8.1Rubidium 2.8.18.8.1Caesium 2.8.18.18.8.1Francium 2.8.18.32.18.8.1

7. They are also known as alkali metals which react with water to form alkaline solutions.

8. All Group 1 elements have one valence electron in their outermost occupied shells.

Physical Properties of Group 1 Elements1. Group 1 elements are soft metals with low densities and low melting points as compared to other metals such as iron and copper.

2. They have silvery and shiny surfaces .

3. They are good conductor of heat and electricity.

Elements/ symbol


Atomic radius (nm)

Melting points (°C)

Boiling points (°C)

Density(g cm-3)

Lithium, Li 2.1 0.15 180 1336 0.57Sodium, Na 2.8.1 0.19 98 883 0.97Potassium, K 2.8.8.1 0.23 64 756 0.86Rubidium, Rb 2.8.18.8.1 0.25 39 701 1.53

Table 1: Physical Properties of Group 1 Elements

6. From Table 1, when going down the group, atomic size and density increase.

7. When going down the group, melting points and boiling points decrease

Chemical Properties of Group 1 Elements

Lithium, sodium and potassium have similar chemical properties but differ in reactivity.

Let us carry out this Experiment!

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Practical Book Experiment 4.1, page 35 Activity 4.3, page 38

1. Alkali metals react vigorously with water to produce alkaline metal hydroxide solutions and hydrogen gas.

[Video]

Chemical equation;

2Li + 2H2O → 2LiOH + H2 Lithium Water Lithium Hydrogen

hydroxide gas

2Na + 2H2O → 2NaOH + H2 Sodium Water Sodium Hydrogen

hydroxide gas

2K + 2H2O → 2KOH + H2 Potassium Water Potassium Hydrogen

hydroxide gas

2. Alkali metals react rapidly with oxygen gas, to produce white solid metal oxides. Chemical equations;

4Li + O2 → 2Li2O Lithium Oxygen Lithium

gas oxide

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4Na + O2 → 2Na2O Sodium Oxygen Sodium

gas oxide

4K + O2 → 2K2O Potassium Oxygen Potassium

gas oxide

3. Alkali metals burn in chlorine gas to form white solid metal chlorides.

Chemical reaction;

2Li + Cl2 → 2LiCl Lithium Chlorine Lithium

gas chloride

2Na + Cl2 → 2NaCl Sodium Chlorine Sodium

gas chloride

2K + Cl2 → 2KCl Potassium Chlorine Lithium

gas chloride

4. Alkali metals burn in bromine gas to form metal bromides.

For example,

2Li + Br2 → 2LiBrLithium Bromine Lithium

gas bromide

2Na + Br2 → 2NaBrSodium Bromine Sodium

gas bromide

2K + Br2 → 2LiBrPotassium Bromine Potassium

gas bromide

5. Therefore, alkali metals have similar chemical properties.

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Alkali metals have one valence electron in their outermost occupied shells.

Each of them reacts by donating one electron from its outermost occupied shell to form an ion with a charge of +1, thus achieving the stable electron arrangement of the atom of noble gas.

Li Li+ + 1e-

2.1 2

Na Na+ + 1e- 2.8.1 2.8

K K+ + 1e- 2.8.8.1 2.8.8

6. The reactivity of Group 1 elements increases down the group.

Going down Group 1, the atomic size (atomic radius) increases.

The single valence electron in the outermost occupied shell becomes further away from the nucleus

Hence, the attraction between the nucleus and the valence electron becomes weaker

Therefore, it is easier for the atom to donate the single valence electron to achieve the stable electron arrangement.

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Why alkali metals have similar chemical properties?

Why The reactivity of Group 1 elements increases down the group?

Safety precautions in handling Group 1 elementsAlkali metals are very reactive. Safety precautions must be taken when handling alkali metals. The elements must be stored in paraffin oil in bottles Do not hold alkali metals with your bare hands Use forceps to handle them Wear safety goggles Wear safety gloves Use a small piece of alkali metal when conducting experiments

Hw: QR C pg. 69 no. 1,2,3

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Documents

1 Matter & the Atomic Structure Modul