F3 Chemistry Notes - · PDF fileF3 Chemistry Notes 1. What is chemistry ? ... water can be...

F3 Chemistry Notes 1. What is chemistry ? What is chemistry ? Chemistry is the science of substances. What is chemistry concerned with ? • Chemistry is concerned with the composition and properties of substances, • how and when substances change, • and what new substances are made in the changes. What are physical properties ? Physical properties are properties that do not change after the observation. Examples of physical properties include colour, density, physical state, smell, hardness, etc., What are chemical properties ? Chemical properties are properties that can only be observed during a change. Examples of chemical properties include reactivity (for e.g. how fast a metal reacts with water), reaction towards air, acids, water and alkalis, etc., What is a physical change ? A physical change is one that does not cause any change to the nature of the substance. During a physical change, no new substance is (new substances are) formed. Most physical changes are reversible, and no new substances are formed during a physical change. For example, freezing water into ice does not change the nature of water (reversible), and tearing a piece of paper into two does not change the nature of the piece of paper, despite that in both cases the appearance of the object have changed (reversible if we recycle the paper). What is a chemical change ? A chemical change is one that changes the nature of the substance. During a chemical change, a new substance is (new substances are) formed. Most chemical changes are irreversible. For example, burning a piece of paper (carbon dioxide is formed) or dropping water onto a piece of sodium metal (hydrogen and sodium hydroxide are formed). Why do we have to study chemistry ? Chemistry can help us to • identify things already known to us (chemical tests, chemical analysis), • make new things that are useful to us (plastics, alloys, etc.), • make things change faster (increase the speed of making medicines, etc.) and • make things change slower (delay the turning bad of food in food preservation). • What is more, proper use of chemical techniques can help us to solve pollution problems. In conclusion, Chemistry is essential for the good living of people in the modern world.

F3 Chemistry Notes – Page 1

2. Mixtures and their separation One of the things that chemists must be able to do is to make pure things. For example, we need our food, water, medicine, etc. to be clean or pure. The fuel for cars and factories need to be free of impurities like sulphur for reducing air pollution (as a result of sulphur oxides). To make things pure we have to be able of separating things from one another. What is a pure substance ? A pure substance contains only one kind of matter. For example, iron (in the form of filings) is a pure substance, sand is another pure substance. What is a mixture ? A mixture contains more than one pure substance mixed together. For example, a mixture of solids is formed when iron filings and sand are mixed together. Air is a mixture of gases (nitrogen, oxygen, etc). Common Methods of Purification 1. Filtration can be used to separate an insoluble substance from a liquid or a solution. For example,

sand can be separated from water by filtration. After filtration, the clear liquid collected is called the filtrate, and the solid left on the filter paper is called the residue. However, filtration cannot separate a soluble substance from a solution. For example, when water is added to a mixture of sand and salt and then the mixture filtered, the residue is sand while the filtrate is salt solution (i.e., salt and water). Filtration does not remove the salt (which is soluble) from the mixture. To obtain dry, pure salt, either crystallisation or evaporation has to be employed.

2. Crystallisation can automatically exclude impurities when it crystallises, so it is a method of purifying things because the impurities are usually not soluble in a hot solvent. For example, pure sugar can be purified from a crude sugar solution (containing impurities and pure sugar) by crystallisation.

3. Evaporation can be used to remove water from a dissolved solid. For example, when sea water is evaporated, salt is left behind. The things to be separated must differ widely in their boiling points.

4. Distillation is evaporation followed by condensation. This can be used to purify a liquid that is made dirty with soluble impurities. For example, when sea water is distilled, pure water (distilled water) will be obtained as the distillate. Salts with high boiling point are not easy to evaporate and remain as residue. The things to be separated must have differ widely in their boiling points.

5. Fractional distillation can be used to separate two liquids with more or less the same boiling points. It is very similar to distillation, but is actually a series of repeated distillations. To simplify the


procedure, a fractionating column is usually used. For example, wine containing more water but less alcohol can be fractionally distilled to make it richer in alcohol (i.e., more concentrated in alcohol).

The method of purifying a mixture depends on the properties of the mixture to be separated. For example, iron filings can be separated from sand by using a magnet. However sawdust (wood powder) cannot be separated from sand by using a magnet. Sand can be separated from sawdust by first adding water and then filtering. (can you explain why?) 3. Elements and Compounds What is an element ? An element is a substance that cannot be further broken down to two or more simpler substances by chemical methods. • An element is one of the simplest substances on Earth, e.g., oxygen, nitrogen, carbon, iron, copper,

silver, gold, iodine, chlorine, neon, argon, etc., are elements.

• Scientists have discovered about 100 elements (some books list 103, some may list up to 108 elements) on Earth. Of these elements, about 90 are found in nature. The rest are made by scientists under high-energy conditions. It is possible that in the future scientists can discover more elements.

• Some elements are more abundant (more plentiful) on Earth. The most abundant elements in the Earth crust are oxygen (49.5%), silicon (25.7%), aluminium (7.5%) and iron (4.7%).

English Name Symbol Latin Name English Name Symbol Latin Name Carbon C Cobalt Co Calcium Ca Copper Cu Cuprum Sulphur S Iodine I Silicon Si Iridium Ir Silver Ag Argentum Iron Fe Ferrum

Sodium Na Natrium Phosphorus P Lead Pb Plumbum Potassium K Kalium



• Each element has a name and a symbol. The symbol is usually made up of a capital letter, or a capital letter followed by a small letter. The symbol of an element usually resembles its English name. In some cases it resembles its Latin name.

• Most elements can be classified as metals and non-metals. Many known elements are metals. The physical properties of metals and non-metals are summarised in the following table :

Metals Non-metals 1. Have bright shiny surfaces when freshly cut or

polished. 2. Are good conductors of heat and electricity. 3. Are malleable (can be hammered into sheets)

and ductile (can be drawn into wires). 4. Are usually solids (except mercury) with high

melting points (except sodium, potassium).

1. Usually have dull surfaces. 2. Are poor conductors of heat and electricity

(except graphite) 3. Are brittle, i.e., break to pieces when

hammered or bent. 4. Usually have low melting and boiling points

(except diamond and graphite). Many are liquids (e.g., bromine) or gases (e.g., chlorine) at room temperatures.

• A few elements have properties that are in between those of metals and non-metals, and are called semi-metals or metalloids, e.g., silicon and germanium.

What is a compound ? A compound is a pure substance made up of two or more elements chemically combined together. • A pure compound has constant composition. This means that no matter where and how a compound

is obtained or made, it is always made up of the same elements combined in the same proportion by mass.

• A pure compound always has the same properties or characteristics (e.g., colour, smell, reactions, etc.). • The properties of a compound is different from those of the elements from which it is made. e.g.,

water can be used to put out a fire, but oxygen can be used to relight a glowing splint, while hydrogen can produce an explosion when it is lit.

• Each compound has a name and a formula. For example, the formula of water is H2O, which means that there are two hydrogen atoms and one oxygen atom in each molecule of water. The formula of carbon dioxide is CO2, and this means that there is one carbon atom and two oxygen atoms in each molecule of carbon dioxide. We can give a fixed formula to each compound because no matter where and how you obtain this compound it always has the same composition.

• Usually an energy change will accompany the formation of a compound. If heat energy is given out during the change, the change is described as exothermic (Greek : exo = outside, theroms = hot). If heat energy is absorbed during the change, the change is described as endothermic (Greek : endo = within).


4. Atomic Structure and the Periodic Table What is an atom ? An atom is the smallest particle of an element that can still have the properties of that element. If a piece of metal is cut into a smaller one, the smaller piece still has the same properties as the original piece of metal. Is an atom the smallest particle in the world ? No. Inside the atom, there are smaller particles called the sub-atomic particles (sub = below, under, or subordinate), namely the protons, the neutrons and the electrons. Properties of the sub-atomic particles are listed in the table below :

Sub-atomic particle Relative charge Relative mass proton +1 1 neutron 0 1

electron -1 11840 , (negligible)

• In an atom, the protons and the neutrons make up a small (diameter of about 10-15 m ) but heavy central particle for the atom, and this particle is called the nucleus (Latin : nucula = little nut). The nucleus carries positive charge.

• At a relatively long distance away, electrons move in orbits (Middle English : orbite = eye socket) around the nucleus. (The diameter of an atom is about 10-10m, 100,000 time the diameter of the nucleus. Imagine a football field : it is about 100m long. 100m × 10-5 = 1mm : if the atom is as big as a football field, the nucleus would be about the size of the head of a pin) The electrons carry negative charges. Between the nucleus and the electrons is an empty space. The electrons are extremely small in size.

Structure of a hydrogen atom :

What is the atomic number of an element ? The atomic number of an element is the number of protons in an atom of the element. • Charges on each proton and on each electron are equal in amount but opposite in sign. • An atom always has an equal number of protons and electrons, so an atom is always electrically neutral. • The atomic number of an element is also the number of electrons in a neutral atom of the element. • The atomic number is a characteristic of an element.


Element atomic number Element atomic number Element atomic number

hydrogen 1 oxygen 8 phosphorus 15 helium 2 fluorine 9 sulphur 16 lithium 3 neon 10 chlorine 17

beryllium 4 sodium 11 argon 18 boron 5 magnesium 12 potassium 19 carbon 6 aluminium 13 calcium 20

nitrogen 7 silicon 14 gold 79

• The atomic number of an element is given the symbol Z. Thus for hydrogen, Z = 1, and for helium, Z = 2.

What is the mass number of an atom ? The mass number of an atom is the sum of the number of protons and the number of neutrons in the atom. • Since the electrons have negligible mass, the mass of an atom is approximately the total mass of the

protons and neutrons. So the relative mass of an atom can be represented by the mass number of the atom.

• The mass number of an atom is given the symbol A. Thus for a helium atom with 2 protons and 2 neutrons, A = 4.

• Since the neutrons are electrically neutral, number of neutrons of an atom bears no direct relationship to the atomic number. For example, most hydrogen atoms have 1 proton but no neutron at all, and in most sodium atoms, there are 11 protons but 12 neutrons. However, many atoms have the same number of protons and neutrons in each atom, e.g., most carbon atoms have 6 protons and 6 neutrons in the nucleus, and most oxygen atoms have 8 protons and 8 neutrons in the nucleus.

• The number of neutrons in the nucleus of an atom can be calculated by using the formula : number of neutrons = A – Z

• The following notation (method of representation) is used to show the structure of the nucleus of atoms :

Mass number → A

X atomic number → Z e.g., the structure of the helium nucleus can be represented as : 4

He 2 What are isotopes ? Isotopes are atoms of the same element but with different number of neutrons in the nucleus of the atoms. In other words, isotopes have same atomic number, but different mass numbers. Isotopes are atoms.


Similarities Differences

of the same element with same atomic number with different mass number

with same number of protons with different number of neutrons

• Most elements are made up of a mixture of many isotopes. For each element with isotopes, the abundance of the isotopes are not equal. For example,

Element isotope % abundance Element isotope % abundancehydrogen 1H 99.985 nitrogen 14N 99.635

2H 0.015 15N 0.365 lithium 6Li 7.42 magnesium 24Mg 78.60

7Li 92.58 25Mg 10.11 carbon 12C 98.892 26Mg 11.29

13C 1.108 chlorine 35Cl 75.53

37Cl 24.47

What is the ‘carbon-12’ scale ? • Atoms are so small and light that it is not useful to show their masses in gram (g) or kilogram (kg).

The mass of an atom is better compared with the mass of another atom. • The mass of an atom of the isotope of carbon, 12C is chosen as the standard for comparing the masses of

different atoms. The mass of a 12C atom is given a relative mass of 12.0000. This is called the carbon-12 scale.

• On the carbon-12 scale, the mass of the 12C atom is the same as its mass number.

• On the carbon-12 scale, the mass of 112 the mass of a 12C atom has a mass of 1.0000. This is the

same as the mass number of the lightest atom, 1H (the most abundant isotope of hydrogen). • Loosely speaking, the mass of an atom on the carbon-12 scale is the number of times that this atom is

as heavy as the hydrogen atom. This is a ratio, so it is a pure number. It has no unit. What is the relative atomic mass of an element ? The relative atomic mass of an element is the average mass of all the naturally occurring isotopes on the carbon-12 scale. • The relative atomic mass is given the symbol of Ar . • The ‘average’ taken is a ‘weighted average’. This is illustrated in the example below : Example : In chlorine, 75.53% of the atoms have the mass of 35 and 24.47% of the atoms have the

mass of 37 on the carbon-12 scale. The average mass of the chlorine atoms

= 35 × 75.53 + 37 × 24.47

100 = 35 × 75.53% + 37 × 24.47% = 35.49

• Notice that a relative atomic mass has no unit. • Notice that the average mass of the isotopes of chlorine, 35Cl and 37Cl, is not 36.


• The relative atomic masses of some elements are not whole numbers because of the existence of isotopes.

Can you calculate the relative atomic masses of (1) magnesium, and (2) hydrogen ? • The relative atomic masses of some elements (like hydrogen) are whole numbers not because there are

no isotopes, but because the abundance of the less common isotopes is negligible. In such cases, the relative atomic mass of the element is (approximately) the same as the mass number of the most abundant isotope.

Electronic Structure of Atoms • The atomic number of an atom is the number of protons in the nucleus of the atom. It is the same as

the number of electrons the atom (for the atom to maintain electrical neutrality). • For example, a hydrogen atom has only one electron. A carbon atom has 6 electrons. A potassium

atom has 19 electrons. • The electrons are placed inside spherical orbits, which we call shells (or electronic shells, energy levels). • The shell nearest to the nucleus has the lowest energy. Shells farther away have higher energy.

• The building-up principle states that electrons must first be filled into the shells of lowest energy

before the shells of higher energy are filled. Electrons are filled into another shell of higher energy when a shell of lower energy if full, until all the electrons are placed into the shells.

• Each shell can only hold a maximum of a certain number of electrons: 1. The 1st shell can hold a maximum of 2 electrons. 2. The 2nd shell can hold a maximum of 8 electrons.

• The next electrons will have to be filled to a higher energy shell when the following number of electrons are already present in the shells :

1. 2 electrons in the 1st shell. 2. 8 electrons in the 2nd shell. 3. 8 electrons in the 3rd shell. You may be interested to know : The 3rd shell can hold a maximum of 18 electrons. The 4th shell can hold a maximum of 32 electrons.

Representing the Electronic Structure of an Atom (1) By energy level diagrams (electronic diagrams) 1. The nucleus is represented by the atomic symbol. 2. The atomic number (= number of protons = number of electrons) of the element is first found.


3. Electrons are filled into the lowest energy shell first, and then to the higher energy shells one by one. 4. Starting from the 2nd shell, the electrons in a shell are shown singly if there are only 4 electrons or less

in the shell. From the fifth electron onwards, the newly added electrons have to be paired with another one of the first 4 electrons already present. Compare the electronic diagrams of carbon, fluorine and sodium below.

The electronic structures of the first 20 elements can be represented as below :

number of electrons in shell Element Z 1st 2nd 3rd 4th

H 1 1 He 2 2 Li 3 2 1 Be 4 2 2 B 5 2 3 C 6 2 4 N 7 2 5 O 8 2 6 F 9 2 7

Ne 10 2 8 Na 11 2 8 1 Mg 12 2 8 2 Al 13 2 8 3 Si 14 2 8 4 P 15 2 8 5 S 16 2 8 6 Cl 17 2 8 7 Ar 18 2 8 8 K 19 2 8 8 1 Ca 20 2 8 8 2

(2) By notation

1. The atomic number of the element is first found. 2. The symbol of the element is written on the left. 3. The number of electrons in each shell is listed, starting from the first shell, next to the symbol. Examples :

O 2,6 P 2,8,5 Ca 2,8,8,2


5. Fossil Fuels Coal, Petroleum and Natural gas as fossil fuels Fossil Fuels Energy is very important to all living things. An important way to obtain energy is by burning fuels. The commonest and most useful fuels are fossil fuels. We have three main fossil fuels : (1) Petroleum (Oil) (2) Coal (3) Natural Gas Coal, petroleum and natural gas are called fossil fuels (化石燃料) because they were formed from the remains of plants and animals that lived millions of years ago. Coal 1. Coal is formed from the remains of plants that lived hundreds of millions of years ago. 2. The plant remains were buried deep inside the earth crust. 3. Under high temperature and pressure the plant remains gradually turned to coal, which consists

mainly of carbon. Petroleum (Oil) and Natural Gas 1. Oil and natural gas come from the remains of plants and animals that lived in the sea hundreds of

millions ago. 2. These remains were buried deep beneath sand and mud. 3. Under high temperature and pressure, and assisted by bacterial action, the remains gradually turned

into oil and natural gas. How long will world fuel resources last Fossil fuels are non-renewable resources. Once burnt they are lost. Fossil fuels have taken millions of years to form and cannot be regenerated in a short period of time. There is a limit within which our fuel resources will be used up. Hence the concept of conservation of fuel is important. List some of the ways to conserve fuels Petroleum as a mixture of hydrocarbons and its separation into useful fractions by fractional distillation Crude oil is a mixture of hydrocarbons. Useful fractions of crude oil can be separated by fractional distillation. In the school laboratory, a simple distillation set-up could be used to separate crude oil. In the petroleum refinery industry, a fractionating tower is used.


Oil fractions and their uses

Name of fractions boiling point (°C) carbon number uses refinery gas below 40 C1-C4 LPG as fuel petrol motor car fuel naphtha

40-170

C5-C10 raw material to make town gas

kerosene 170-250 C10-C16 fuel for jet aircraft gas oil 250-350 C16-C25 raw material to make

more gasoline fuel oil fuel for buses, trucks

and ships lubricating oils and waxes

to make candles and polishes

bitumen

>350

> C25

road surfacing Fractions usually differ in terms of boiling points, colour, volatility, viscosity and burning characteristics. The following table compares the properties of the light and heavy petroleum fractions. Light Fraction Heavy Fraction Boiling points lower higher Colour colourless yellowish Volatility more volatile less volatile Viscosity watery more viscous Burning characteristics burns without residue burns with a smoky flame, leaving

a residue


6. The Mole Avogadro Number

H O

An oxygen atom is bigger in size, and is heavier, than a hydrogen atom. Actually an O atom is approx. 16 times more heavy than an H atom. It is difficult to weigh one atom because we could not have such an accurate balance. Chemists have introduced a quantity called mole so that if we weigh one mole of atoms, it would be 6.02 × 1023 atoms. Avogadro number, L is equal to 6.02 × 1023 mol-1. How heavy is one atom of hydrogen ? The approximate mass for 1 mole, i.e., 6.02 × 1023 atoms of hydrogen is 1 g. Calculate the mass of 1 atom of hydrogen.

Mass of 1 atom of hydrogen = 1

6.02 × 1023 = ________ g

Similarly calculate the mass for 1 atom of oxygen. The approximate mass for 6.02 × 1023 atoms of oxygen is 16 g. Mass of 1 atom of oxygen = __________ = ________ g Work out the ratio of mm

O

H =

where mO is the mass for one atom of oxygen and mH is mass for one atom of hydrogen Complete the following table

Isotope Relative mass of isotope Mass of 6.02 x 1023 particles12C 12 12 g 32S

35Cl 35 g 37Cl 37

207Pb2+ 37Cl-

Note : Pb2+ and Cl- are called ions. An ion is a charged particle made from an atom or a group of atoms. In these examples, Pb2+ is a lead atom which has lost 2 electrons, and Cl- is a chlorine atom that has gained 1 electron.


Relative Masses of Molecules A molecule is made up of two or more than two atom chemically combined together. We can work out the relative masses for molecules by simple addition. For HCl, the relative molecular mass = 1 + 35.5 = 36.5. For H2O, the relative molecular mass = 2(1) + 16 = 18. Complete the following table : (Given : H = 1, O = 16, Cl = 35.5, C = 12)

Molecule Relative Molecular Mass H2 (hydrogen molecule) 2(1) = 2 O2 (oxygen molecule)

Cl2 (chlorine molecule) O3 (ozone, found in upper layers of the atmosphere)

CO2 (carbon dioxide gas) CHCl3 (trichloromethane, or chloroform)

Molar Mass The relative molecular mass of H2 = 2. For one mole of H2, we simply add the unit grams to get the mass of H2 = 2 g Hence, the molar mass of H2 = 2 g (or : 2 g mol-1) What would be the molar mass for O2 ? ____ g ( ____ g mol-1) Relationship between Molar Mass and Mass The molar mass of H2 is 2 g. That means if we had 2 g of H2, we have one mole of H2. How many moles of H2 do we have if we have got only 1 g of H2 ? _____ A simple relationship between mass, molar mass and number of moles is as follows :

Amount in moles = mass

molar mass

Example (1) How many moles of molecules are there in 64 g of oxygen molecules, O2 ? Given : relative atomic mass of O = 16 Molar mass of O2 = 2 (16) = 32 g


molar mass = 6432 = 2

(2) Calculate the number of moles of H2O molecules in 72 g of water Given : relative atomic masses O = 16, H = 1 Molar mass of H2O = 2 (1) + 16 = 18 g


molar mass = 7218 = 4

(3) Calculate the mass of 1 atom of O2, given that the relative atomic mass of O = 16.


Chemical Formulae A chemical formula is an expression (or a symbol) to show the composition of a compound (or a substance). E.g.,

He H2 H2O NaCl an element, the formula shows that there is only

one helium atom per particle

an element, the formula shows that there are

two atoms of hydrogen per particle (molecule)

a compound, the formula shows that

there are two hydrogen atoms and one oxygen

atom per particle (molecule)

a compound, the formula shows that the ratio of sodium atom to

chlorine atom in the compound is

1 : 1 The chemical formula of a compound can be obtained by working out the ratio in moles for each element in the compound. The following diagram shows the apparatus used in an experiment to determine the formula of black copper oxide :

Town gas (containing mainly hydrogen by volume) was passed over the oxide before heating began. During heating, the copper oxide changed colour from black to light brown. After some time, heating was stopped, but town gas was allowed to pass over the copper until it was cold. The results of a typical run of the above experiment were as follows : Mass of tube = 20.10 g Mass of tube and copper oxide = 22.68 g Mass of tube and copper = 22.16 g (1) Town gas was burnt at the jet after it has passed over the copper oxide, otherwise the posionous

component (carbon monoxide) in town gas can kill, and the main component (hydrogen) may explode.

(2) The experimental results are interpreted as follows : (a) Mass of copper = 22.16 - 20.10 = 2.06 g (b) Mass of oxygen = 22.68 - 22.16 = 0.52 g (c) The moles of copper and oxygen can be worked out as follows :

Cu O

Mass (in g) 2.06 0.52

Number of moles 2.06/63.5 = 0.0324 0.52/16 = 0.0325 Mole ratio 1 1

(d) The mole ratio of Cu to O is 1 : 1, hence the formula of the black copper oxide is CuO


Empirical Formula The empirical formula of a compound is the formula that shows the simplest ratio of atoms of the elements present in the compound. Formulae that show the simplest ratio are called empirical formulae because they can be easily obtained from experiments. (empirical = relying on or derived from observation or experiment.) The chemical formula of a compound could be worked out if the mass of each element contained in the compound is worked out. The masses of each element are then converted into number of moles. The ratio in moles of each element gives the chemical formula of the compound. The chemical formula obtained from this method (from an experiment) is called the empirical formula. Example : Mass analysis of a compound containing C, H and O only gives the following information. Total mass of compound = 4.4 g Mass of C = 2.4 g Mass of H = 0.4 g Work out the empirical formula of the compound (Relative atomic masses : C = 12, H = 1, O = 16) Solution Mass of O = Total mass - Mass of C - Mass of H = 1.6 g

C H O

Mass (g) 2.4 0.4 1.6 Number of moles 2.4/12 = 0.2 0.4/1 = 0.4 1.6/16 = 0.1 Mole Ratio 2 4 1

The Empirical Formula of the compound is C2H4O Molecular Formula The empirical formula only shows the ratio of atoms in a compound. To find out the exact number of atoms in a compound, the molar mass of the compound has to be stated. Example A compound P has the empirical formula C2H4O. P has a molar mass of 88 g. Determine the molecular formula of P. Let the molecular formula of P be (C2H4O)n Molar mass of C2H4O = 2(12) + 4(1) + 16 = 44 44n = 88 ⇒ n = 2 Hence the molecular formula of P is (C2H4O)2. The formula simplifies to C4H8O2. Structural Formula Structural formulae give further information of compounds.

The arrangement of atoms in space is given by the structural formula. F3 Chemistry Notes – Page 16

7. The Periodic Table and Periodicity How are the elements arranged in the Periodic Table ? • In the Periodic Table, the elements are arranged in order of increasing atomic numbers. • The horizontal rows in the Periodic Table are periods. There are 7 periods. • The vertical column in the Periodic Table are groups. There are 8 main groups. • Starting from period 4, more elements appear between Group II and Group III. These elements are the

transition metals, e.g., iron (Fe), copper (Cu), mercury (Hg), silver (Ag) and gold (Au). What is the meaning of ‘periodicity’ ? • The Periodic Law states that when the elements are arranged in order of increasing atomic numbers,

the elements having similar properties reappear at regular intervals. Similar properties reappear at intervals of eight elements for the first twenty elements.

• The phenomenon of properties of elements recurring at regular intervals is called periodicity. • A characteristic or property that reappears at regular intervals is said to show periodic variation. E.g.

Li, Na and K, all react towards water similarly. What are the characteristics of a period ? • Atoms of elements in the same period have the same number of electronic shells. • The period number is the number of electronic shells. e.g., sodium (Na), which is in period 3, has 3

shells (2, 8, 1). • Metallic character decreases and non-metallic character increases across a period from left to right.

e.g., in period 3, sodium (Na, on the left) is a metal; chlorine (Cl, on the right) is a non-metal. • A diagonal line in the Periodic Table passing through boron and silicon roughly divides the metals

from the non-metals. Elements to the left of this diagonal line are metals and those to the right are non-metals. Elements lying just besides the diagonal line are semi-metals, e.g. silicon and germanium.

• Elements on the far left of the Periodic Table are very reactive metals while those on the far right are very reactive non-metals.

• The group 0 elements on the extreme right of the Periodic Table are very unreactive. They are called the noble gases. e.g., helium (He), neon (Ne), argon (Ar), krypton (Kr) and xenon (Xe).

What are the characteristics of a group ? • Elements in the same group have the same number of electrons in their outermost shells. • The group number of an atom is equal to the number of outermost shell electrons in the atom. • The chemical properties of an element depend on the number of electrons in its outermost shell. • As elements in the same group have the same number of outermost electrons, they have similar

properties.


Group similarities and trends in chemical properties What are noble gases and why are they unreactive ? • All noble gases (i.e. group 0 elements) have eight electrons in the outermost shell (an octet) except

for helium, which has 2 electrons in its only one shell (a duplet). • The special electronic structures of noble gases (duplet or octet) are very stable and that is why noble

gases are unreactive. i.e., noble gases do not want to have any change to their stable electronic structures.

What is metal character ? • Metals are elements on the left side of the Periodic Table. • The outermost electrons of the metals are not so firmly held. Metals tend to lose electrons from its

outermost shell to form positive ion (or cation, i.e., positively charged particle) which has the same electronic structure as the nearest noble gas. (Cation is the positive ion, it contains the letter “t” that looks like a “+” sign.)

• The atoms that tend to lose electrons from their outermost shells are said to show metallic character.

e.g., the Na atom tends to lose one electron to form a Na+ ion : Na → Na+ + e- 2,8,1 2,8 electron ion with stable noble gas structure

e.g., the Mg atom tends to lose two electrons to form a Mg2+ ion : Mg → Mg2+ + 2e- 2,8,2 2,8 electrons ion with stable noble gas structure What is the meaning of electropositive ? • Metals tend to form positively charged ions by losing electrons. They are said to be electropositive. (Question : Which is a more electropositive element, sodium or magnesium ?) .................................................................................................................................................................... • On going across a period, the electropositive (or metallic) character of the elements decreases. • On going down a group, the electropositive character of the elements increases. The outermost shell

electrons gets further away from the nucleus and thus the electrons are easier to be lost. (Question : Sodium and potassium are both group I metals, which is more electropositive ?) .................................................................................................................................................................... What is non-metal character ? • A non-metal atom tends to attract or gain electrons to its outermost electron shell to form a

negative ion (or an anion) which has the same electronic structure of the nearest noble gas.


• Non- metals are elements on the right side of the Periodic Table.

e.g., the Cl atom tends to gain one electron to form a Cl- ion : Cl + e- → Cl- 2,8,7 electron 2,8,8 ion with stable noble gas structure

e.g., the S atom tends to gain two electrons to form a S2- ion. S + 2e- → S2- 2,8,6 electrons 2,8,8 ion with stable noble gas structure • The atoms that tend to gain electrons to form negative ions are said to show non-metallic character.

(Question : Which is a more electronegative element, chlorine or sulphur ?) .................................................................................................................................................................... • On going across a period, the non-metallic character (or electronegativity) increases. • On going down a group, the non-metallic character (or electronegativity) decreases. The attraction

between the nucleus and the electrons in the outermost electron shell is weaker because the number of electronic shells increases. Thus the atoms have weaker tendencies to gain electrons into their outermost shell.

(Question : Fluorine and chlorine both are group VII elements, which is more electronegative ?) .................................................................................................................................................................... Electropositivity and Electronegativity In general, metals are electropositive and non-metals are electronegative. (Why ?)


8. Conductors, Non-conductors and Electrolytes In Topic 3, we have classified elements into conductors and non-conductors (or insulators). This classification cannot be applied to compounds. None of the compounds are conductors of electricity, whereas some compounds can conduct electricity under certain specific conditions, which we call them electrolytes. Compounds that cannot conduct electricity under all conditions are called non-electrolytes. What is an electrolyte ? A compound that can conduct electricity when molten or in aqueous solution is called an electrolyte. What are conductors ? • They are elements. • They conduct electricity but are not chemically changed after the conduction of electricity. • Metals are examples of conductors. They conduct electricity in both solid and liquid states because

they have free electrons in both physical states. What are insulators ? • They include both elements and compounds. • They do not conduct electricity in any state. • They include all non-metal elements (except graphite) and compounds that have no free electrons. Characteristics of electrolytes : • Electrolytes include ionic compounds (compounds made up of both positive and negative ions). • They do not conduct electricity in the solid state but conduct in the molten or aqueous (dissolved in

water. Latin : aqua = water) state. • During conduction, they decompose (break down; de- = opposite, Latin : com- = together, p¢nere =

put.). What is electrolysis ? • Electrolysis is the decomposition of an electrolyte by the passage of electric current through the

electrolyte in the molten or aqueous state (lysis :Greek , lusis = to loosen) Examples of conductors, insulators and electrolytes :

Conductors (elements) Insulators (elements or compounds)

Electrolytes (compounds)

mercury iron

copper lead

graphite (non-metal)

sugar distilled water

plastics wax

sulphur (an element)

sodium chloride lead(II) bromide

copper(II) sulphate copper(II) chloride potassium iodide


How do electrolytes conduct electricity ? • Electrolytes are made up of charged particles called ions. Thus they are ionic compounds. • Ions are electrically charged particles made from atoms or groups of atoms. • Positively charged ions are also called cations. • Negatively charged ions are also called anions. • On passing electricity through molten or aqueous electrolytes, the (+ve) cations move to the negative

electrode (cathode) and gain electrons from the electrode. The (-ve) anions move to the positive electrode (anode) and lose electrons to the electrode.

• Electrons do not flow directly through the electrolyte, but electric current is carried through the electrolyte by the movement of ions.

• Movement of ions in the liquid state or the molten state is also known as the migration of ions. Why electrolytes cannot conduct electricity in the solid state ? • In the solid state, the cations (+ve) and anions (−ve) attract each other strongly. The ions are held

together firmly in fixed positions. The ions are therefore not free to move and the electrolyte cannot conduct electricity.

Why can electrolytes conduct electricity in the molten or aqueous state ? • In molten or in aqueous state, the strong attractive forces between ions are weakened. Thus, the ions

become free to move (i.e., become mobile) and conduct electricity. An example of the conduction of electricity through an electrolyte molten lead(II) bromide

• The positive lead ions gain electrons from the cathode and form neutral lead atoms. Metallic lead

is seen around the cathode.

positive lead ions + electrons → neutral lead atoms (electrons from cathode)

• The negative bromide ions lose electrons to the anode and form neutral bromine atoms, which then further reacted to become bromine molecules. Red bromine vapour is seen around the anode.

negative bromide ions → neutral bromine atoms + electrons (electrons carried away by the anode)


9. Chemical Bonding (Ionic Compounds and Covalent Substances)

Why do elements combine to form compound ? Example : Two elements, sodium and chlorine, combine to form sodium chloride.

• When sodium is ignited and lowered into a jar of chlorine, it continues to burn and heat is given out. A white solid is left behind which is sodium chloride.

• Sodium and chlorine have higher potential energy compared with sodium chloride. When ignited, the elements combine and release energy. A chemical bond which links sodium and chlorine atom together is formed. A compound, sodium chloride (i.e. the white salt), is also formed which has lower potential energy and is more stable than the two elements.

sodium + chlorine → sodium chloride (elements : higher potential energy) (compound : lower potential energy)

IONIC BONDING What are ionic compounds ? Compounds made up of both positive and negative ions are called ionic compounds. Are there any evidence for the existence of ions ? Evidence of the existence of ions migration of ions


• In the above experiment, the U-tube is filled with a solution of copper(II) sulphate and potassium dichromate in gel. Water containing some dilute sulphuric acid for better conduction of electricity is added carefully to each limb of the U-tube, so as to avoid mixing of the layers.

• The circuit is closed. After some time, a blue colour appears around the cathode region and an orange colour appears around the anode region.

• The reason for the above observation: cations (+ve), Cu2+ ions, move towards the cathode (-ve) and the anions (-ve), Cr2O7

2- ions, move towards the anode (+ve). How and why ions are formed ? • Except helium, all atoms of noble gases have 8 electron in the outermost shell, i.e. an octet structure

of electrons in the outermost shell. It is believed that noble gases are unreactive because their electronic structures are very stable.

• Atoms of elements other than noble gases do not have such stable electronic arrangements. They can become stable by gaining or losing electron(s) to attain the nearest noble gases structure.

• A simple ion is formed when an atom loses or gains electron(s). • Ions are electrically charged particles. Positive ions (Cations) • Metal atoms lose their outermost shell electron(s) to form positive ion. • The positive metal ion has the stable electronic structure of the nearest noble gas. • e.g., each sodium atom loses one electron and becomes a sodium ion.

sodium atom, Na(2, 8, 1)

sodium ion, Na+(2, 8)

11 protons = 11+10 electrons = 10-

__________________overall charge = +1

loses1 electron

+

+ 1e-. .. .

.... ..11p

12n. .

..

.... ...11p

12n

Negative ions (Anions) • Non-metal atoms gain electron(s) until there are eight electrons in the outermost shell to form negative

ions (except hydrogen). The negative non-metal ion has the stable electronic structure of the nearest noble gas.


• e.g., each chlorine atom gains one electron to become a chloride ion.

chlorine atom, Cl(2, 8, 7)

chloride ion, Cl-(2, 8, 8)

17 protons = 17+18 electrons = 18-

__________________overall charge = -1

gains1 electron

-

+ 1e- .17p18n

x x

x x

x x

xxx

x x

x x

xx

x x

17p18n

x x

x x

x x

xxx

x x

x x

xx

x x

What is an ionic bond ? • An ionic bond is the strong electrostatic attractive force between cations (+ve) and anions (-ve). • A complete transfer of electrons from a metal atom to a non-metal atom can form an ionic bond. • Example: sodium chloride 1. The sodium ions Na+ and chloride ions Cl- in sodium chloride are formed by the transfer of electrons

from sodium atoms to chlorine atoms. Both the ions formed have stable noble gas electronic structures.

2. The positive Na+ ions and the negative Cl- ions are then attracted and held together by strong electrostatic forces.

3. Electronic diagram for sodium chloride :

sodium atom

Na

(2, 8, 1)

. .. .

.... ...Na

chlorine atom

Cl

(2, 8, 7)

sodium ion

[Na]+

(2, 8)

chloride ion

[ Cl ]-

(2, 8, 8)

. xx x

xx

x xx

xx

x

xx x

.

-+

. .. .

.... ..Na .Clx x xx

xx

x

x x

x xxx

x

x

xx

Clx x xxx

x

x

x x

x xxx

x

x

xx

Notice that in the second representation above, only the outermost shells are shown.


•Example: Magnesium fluoride

. .. .

.... ...Mg

2+

. .. .

.... ..Mg.

Magnesium atom

Mg

(2, 8, 2)

fluorine atoms

2 F

(2, 7)

magnesium ion

[Mg]2+

(2, 8)

fluoride ions

2 [ F ]-

(2, 8)

.. xx x

xx

x x

-

-

.x

xx

x

xx x

Fx x x

x

x

x

xx

Fx x x

x

x

x

xx

. Fx x x

x

x

x

xx

. Fx x x

x

x

x

xx

In the space below, sketch electronic diagram to show how the compound Al2O3 could be formed from its elements Al and O. Electronic diagram of Al2O3 What is a simple ion ? • A simple ion is a monatomic ion formed from a single atom, e.g., Cl-, Na+. What is a polyatomic ion ?

• A polyatomic ions is an ion made from two or more atoms, e.g., OH-, SO42-.


NAMES AND FORMULAE OF COMMON IONS Cations Anions

Formula Name Colour Formula Name Colour

Na+ sodium ion colourless Cl- chloride ion colourless K+ potassium ion colourless Br- bromide ion colourless Ag+ silver ion colourless I- iodide ion colourless H+ hydrogen ion colourless OH- hydroxide ion colourless NH4

+ ammonium ion colourless NO3- nitrate ion colourless

Cu+ copper(I) ion red MnO4- permanganate ion purple

Hg+ mercury(I) ion colourless NO2- nitrite ion colourless

HCO3- hydrogencarbonate ion colourless

HSO4- hydrogensulphate ion colourless

OCl- hypochlorite ion colourless H- hydride ion colourless

Mg2+ magnesium ion colourless O2- oxide ion colourless Ca2+ calcium ion colourless CO3

2- carbonate ion colourless Ba2+ barium ion colourless SO4

2- sulphate ion colourless Pb2+ lead(II) ion colourless Cr2O7

2- dichromate ion orange Fe2+ iron(II) ion pale green CrO4

2- chromate ion yellow Co2+ cobalt(II) ion pink S2- sulphide ion colourless Ni2+ nickel(II) ion blue / green SO3

2- sulphite ion colourless Mn2+ manganese(II) ion colourless (Mn2+ ions are actually very pale pink in colour) Cu2+ copper(II) ion blue Zn2+ zinc ion colourless Hg2+ mercury(II) ion colourless

Al3+ aluminium ion colourless PO43- phosphate ion colourless

Fe3+ iron(III) ion brown (Dilute Fe3+ ion solutions are yellow in colour) Cr3+ chromium(III) ion green Formulae and naming of ionic compounds How can the formulae of an ionic compound be accurately predicted ? • The formula of an ionic compounds shows the ratio of different ions in the compound. • The symbol of the cation (+ve) is always written first, followed by the symbol of the anion (-ve). • All ionic compounds are overall electrically neutral, thus in an ionic compound the total positive

charge must be equal to the negative charge.


Examples: (a) sodium sulphate is made up of Na+ and SO4

2- ions. One negative SO42- ion needs two positive

Na+ ions to make the compound overall electrically neutral. Ions Na+ SO4

2-

Ratio 2 : 1 Formula Na2SO4

Note : the number “2” in the above formula is called the subscript. The subscript tells how many times the ion occurs in each formula.

(b) magnesium hydroxide contains Mg2+ and OH- ions. One positive Mg2+ ion needs two negative

hydroxide ions to make the overall charges become zero. As the OH- ion is a polyatomic ion, the symbol of the OH- ion is placed in brackets and the subscript is written outside the bracket.

Ions Mg2+ OH- Ratio 1 : 2 Formula Mg(OH)2

Note : the bracket is needed when both of the following conditions are satisfied --

1. the ion in the bracket is a polyatomic ion, and 2. the ion occurs more than once in each formula of the coumpound. Do not use a bracket if any one of the above conditions are not satisfied.

(c) Some more examples (Can you explain why these compounds should have such formulae ?) :

K2Cr2O7 (NH4)2SO4 Al2O3 Pb(NO3)2 How do we give names to ionic compounds ? • The cation (+ve) is named first, followed by the name of the anion (-ve). • The name of a simple non-metal ion always ends in the letters -ide. For example, ZnO is called zinc

oxide, FeS is called iron sulphide, NaCl is called sodium chloride. • The name of polyatomic ions containing oxygen always end in -ate or ite. For example, MgSO4 is

called magnesium sulphate, KMnO4 is called potassium permanganate, (NH4)2Cr2O7 is called ammonium dichromate.

• Many metals, especially transition metals, can form ions with different ionic charges. The charge is then indicated by Roman numerals written in brackets after the name of the metal. For example, Fe(OH)2 is iron(II) hydroxide and Fe(OH)3 is iron(III) hydroxide.


COVALENT BONDING How and why are covalent bonds (covalent molecules) formed ? • Instead of transferring electrons to become ions, non-metal elements can also get the stable electronic

structure of noble gas by sharing electrons. • Bonds formed by sharing electrons are called covalent bonds. • Covalent bonds are found in non-metallic elements and compounds formed between non-metallic

elements. • In the example below, each atomic nucleus (positively charged) has an attraction for the shared electron

pair (negatively charged). Therefore, the two atoms are held firmly together.

or : H + H H H or H--H

.H H H.H

. xx

shared electron pair: a covalent bond

+

hydrogen atoms hydrogen molecule

.

xx

What is a molecule ? A molecule is formed when two or more atoms combine together by covalent bonds only. What is the atomicity of a molecule ? The number of atoms present in each molecule is called its atomicity. • Molecules formed by combining two atoms only are called diatomic molecules. e.g., O2, H2. • Molecules formed by combining three atoms are called triatomic molecules. • Molecules formed by combining many atoms together are called polyatomic molecules. Examples for formation of covalent bond : The chlorine molecule • The electronic structure in notation of chlorine is 2, 8, 7. The outermost electronic shell of chlorine

needs one more electron in order to have a stable noble gas structure. • Two chlorine atoms can each uses one electron for sharing. Each chlorine atom now has 8 electrons in

its outermost shell and it can then be considered to have a stable noble gas structure. • The bond in the chlorine molecule is a single covalent bond (or single bond) since it is formed by one

shared electron pair (or two shared electrons).

. .

.... .Cl

chlorine atoms chlorine molecule

or : Cl + Cl Cl Cl or Cl--Cl

. .

.... .Cl

.............. x

x x

xxx

x

xx x

xxx

x

...... x x

xxx

x

Clx

x x

xx

x x

Clx

x x

xx

x x


The oxygen molecule • The electronic structure in notation of oxygen is 2, 6. The outermost shell of oxygen needs two more

electrons in order to have a stable noble gas structure. • Two oxygen atoms can each use two electrons for sharing. Each oxygen atom now has 8 electrons in

its outermost shell and it can then attain the stable noble gas electronic structure. • The bond in the oxygen molecule is a double bond since it is formed by two shared electron pairs (or

four shared electrons).

. .

....O O

oxygen atoms oxygen molecule

or : O + O O O or O = O

. .

....O

....... xx

xx x

xx

x

..

..x. x x

x

x x

x x....

Ox

x x

x

x x

x x

xx

x x

The nitrogen molecule • The electronic structure in notation of nitrogen is 2,5. The outermost shell of nitrogen needs three

more electrons in order to have stable noble gas structure. • Two nitrogen atoms can each use three electrons for sharing. Each nitrogen atom now has 8 electrons

in its outermost shell and it can attain the stable noble gas electronic structure. • The bond in nitrogen molecule is a triple bond since it is formed by three shared electron pairs (or

six shared electrons).

. .. ..N

nitrogen atoms nitrogen molecule

or : N + N N N or N N

. .. ..

N

.. .. .. x x

x xxx x

..x. xx xx.. .

Nx

x

xxx

Nx

x

xxx

More examples : The water molecules, H2O • Hydrogen atom needs one more electron to become stable and the oxygen atom needs two more

electrons to become stable. • A hydrogen atom and an oxygen atom each uses one electron for sharing. The hydrogen atom then

can be considered to have a stable structure as helium. The oxygen atom, however, needs one more hydrogen atom for sharing electrons in order to have a stable noble gas structure. Thus, two single bonds are formed .

• The electron pairs in the outermost shell not involved in the bonding are called lone pairs or non-bonding electron pairs.


bondpair

H

The water moleculehas 2 bond pairs

and 2 lone pairs of

lone pair .H

..H . H

x

O

x

x x

xx

x

O

x

x x

xx

formation of water (H2O)

The CH4 and NH3 molecules

C.H

H

.. H

.H

N.H . H

H

.

x

x

x

x

xx

x x

x

methane (CH4) ammonia (NH3)

End of F3 Chemistry Notes


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