STRUCTURE OF THE ATOM AND THE PERIODIC TABLE - Chemistry Notes Form 2

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The Atom

  • Refers to the smallest particle of an element hat can take part in a chemical reaction;
  • It has an average diameter of 10-8 cm with a nucleus of about 10-13 cm;

Parts of an Atom

  • The atom is made of two main parts:
    • The nucleus
    • The energy levels;

The nucleus

  • Is the positively charged part of an atom;
  • The nucleus contains two subatomic particles; neutrons and protons;
  • The positive charge is due to presence of protons;
  • The nuclei of all atoms contain neutrons except the hydrogen atom;
  • The protons and the neutrons are together referred to as the nucleons ;

The energy levels

  • They contain the electrons;
  • Electrons are so small and move so fast that their path cannot be traced directly;
  • Thus the energy level simple represents the region where the electrons are most likely to be found;

Structure of the Atom

structure of the atom
Note:

  • The atom can still however be split into smaller particles termed the sub-atomic particles;

The sub-atomic particles.

- Are generally three:

  • Protons;
  • Neutrons;
  • Protons;

Protons

  • Are the positively charged sub-atomic particles;
  • Are found in the nucleus and thus form part of the nucleons;
  • The number of protons in the nucleus is equal to the number of electrons in the energy levels;

Neutrons

  • Are neutrally charged sub-atomic particles found in the nucleus of the atom;
  • They are thought to probably prevent the positively charged protons from getting too close to each other;

Electrons

  • Are negatively charged sub-atomic particles found in the energy levels;
  • The number of electrons in the energy levels is equal to the number of protons in the nucleus;
  • This makes the atom to be electrically neutral;

 

 

Atomic Number and Mass Number

Atomic number

  • Refers to the number of protons in the nucleus of an atom;

Examples

  • Sodium has 11 protons in the nucleus and thus said to have atomic number 11;
  • Chlorine has 17 protons in the nucleus and thus said to have atomic number 17;

Mass number;

  • Refers to the sum of the number of protons and neutrons in an atom of an element;

Examples:

  • Sodium has 2 neutrons and 11 protons hence a mass number of 23;
  • Chlorine has 18 neutrons and 17 protons hence a mass number of 35.

Notation of Atomic Number and Mass Number

  • Both atomic number and mass number of an element can be written along with the symbol of an element;

Mass number;

  • Is conventionally represented as a superscript to the left of the symbol;

Examples:

  • Sodium; 23Na;
  • Magnesium 24Mg;

Atomic number;

  • Is conventionally represented as a superscript to the left of the symbol;

Examples:

  • Sodium; 11Na;
  • Magnesium 12Mg;

Thus the elements can be conventionally represented as:

  • Sodium 2311Na
  • Magnesium 2412Mg

Atomic Properties of the First 20 Elements.

Element Symbol Number of electrons Number of protons Number Of neutrons Atomic number Mass number
Hydrogen H 1 1 0 1 1
Helium He 2 2 2 2 4
Lithium Li 3 3 4 3 7
Beryllium Be 4 4 5 4 9
Boron B 5 5 6 5 11
Carbon C 6 6 6 6 12
Nitrogen N 7 7 7 7 14
Oxygen O 8 8 8 8 16
Fluorine F 9 9 10 9 19
Neon Ne 10 10 10 10 20
Sodium Na 11 11 12 11 23
Magnesium Mg 12 12 12 12 24
Aluminium Al 13 13 14 13 27
Silicon Si 14 14 14 14 28
Phosphorus P 15 15 16 15 31
Sulphur S 16 16 16 16 32
Chlorine Cl 17 17 18 17 35
Argon Ar 18 18 22 18 40
Potassium K 19 19 20 19 39
Calcium Ca 20 20 20 20 40

Isotopes

  • Are atoms of the same element with same atomic number but different mass number due to different number of neutrons.

Examples of isotopes

Element Isotope Atomic No. Number of protons Number of neutrons Mass number Isotopic
representation
Hydrogen  Protium 1 1 0 1 11H
Deuterium 1 1 1 2 21H
Tritium 1 1 2 3 31H
Carbon Carbon-12 6 6 6 12 126C
Carbon-13 6 6 7 13 136C
Oxygen Oxygen-16 8 8 8 16 168O
Oxygen-17 8 8 9 17 178O
Oxygen-18 8 8 10 18 188O
Chlorine Chlorine-35 17 17 18 35 3517Cl
Chlorine-37 17 17 20 37 3717Cl


Energy Levels and Electron Arrangements

Energy Levels

  • Are definite orbits in an atom that the electrons occupy.
  • The energy levels are numbered 1, 2, 3 starting with the one closest to the nucleus.
  • Electrons occupying the same energy level have approximately the same amount of energy.
  • Each energy level can only accommodate a given maximum number of electrons.

Maximum number of electrons per energy level

Energy level Maximum number of electrons
1st 2
2nd 8
3rd 8 (only for the first 20 elements)

Illustrations

Hydrogen

  • It has only one electron and thus this electron occupies the first energy level.
  • Since the first energy level is not yet full, hydrogen does not have the second energy level;
  • The electron arrangement of hydrogen is thus 1.

Helium

  • Helium is atomic number 2 and has only two electrons, which occupy the first energy level.
  • The first energy level is thus completely full, but since there are no other electrons lithium also has only one energy level;
  • The electron arrangement is thus 2.

Chlorine

  • Chlorine has atomic number 17 and thus has 17 electrons;
  • The first two electrons occupy the fist energy level which is thus completely filled up;
  • The remaining 15 electrons occupy the second energy level, which can however accommodate only 8 to be completely filled up;
  • Thus the remaining 7 electrons move to the third energy level; which needs 8 to be completely filled up;
  • Since the third energy level is not yet full chlorine does not have a fourth energy level;
  • The electron arrangement is thus 2.8.7.

Electron Arrangement.

  • Refers to the distribution of electrons in the energy levels of an atom.

Example: electron arrangement for the first 20 elements.

Element Symbol Atomic number No. of electrons Electron arrangement
Hydrogen H 1 1 1
Helium He 2 2 2
Lithium Li 3 3 2.1
Beryllium Be 4 4 2.2
Boron B 5 5 2.3
Carbon C 6 6 2.4
Nitrogen N 7 7 2.5
Oxygen O 8 8 2.6
Fluorine F 9 9 2.7
Neon Ne 10 10 2.8
Sodium Na 11 11 2.8.1
Magnesium Mg 12 12 2.8.2
Aluminium Al 13 13 2.8.3
Silicon Si 14 14 2.8.4
Phosphorus P 15 15 2.8.5
Sulphur S 16 16 2.8.6
Chlorine Cl 17 17 2.8.7
Argon Ar 18 18 2.8.8
Potassium K 19 19 2.8.8.1
Calcium Ca 20 20 2.8.8.2

Dot and Cross Diagrams

  • Is a diagrammatic representation of the electron arrangements in an atom in which the energy levels are represented by concentric lines while electrons are represented by dots or crosses.
  • However all electrons are the same regardless of whether they are represented as dots or crosses.

Examples: (Refer to the number of electrons, protons and neutrons in the tables above)

 

Lithium

lithium dot cross diagram

Magnesium

magnesium dot cross diagram

Aluminium

aluminum dot cross diagram

Carbon

carbon dot cross diagram



The Periodic Table

the periodic table

  • Is a table showing the arrangements of elements in order of their relative atomic masses.
  • Is based on the ideas of Dmitri Ivanovich Medeleev.
  • The modern periodic table is based on Meneleev's periodic law which states that:
    The properties of elements are a periodic functions of their relative atomic masses
  • The modern periodic law itself states that: The properties of elements are a periodic functions of their atomic numbers

Design of the Modern Periodic Table

It has vertical columns called groups and horizontal rows called periods.

Groups:

  • Are the vertical columns of a periodic table.
  • Are eight in number; and numbered in capital Roman numerals I all through to VIII.
  • Note: group VIII is also called group zero because the elements have little tendency to gain or lose electrons during chemical reactions
  • Between group 2 and group 3 is a group of elements called the transition metals;

Periods:

  • Are the horizontal rows in a periodic table.
  • They are 8 in number in a modern periodic table.

Transition metals.

  • Are elements that form a shallow rectangle between group II and group III.
  • These elements are generally metallic and hence the name transition metals
  • They are not fitted in any group because they have variable valencies.
  • Are sometimes called the d-block elements.
  • They are much less reactive than the elements in groups I and II
  • They have some unique characteristics that make them not fit in the 8 groups of the periodic table. These include:
    • Have variable valencies; hence show different oxidation states in their compounds;
    • They form coloured compounds as solids and in aqueous solutions;
    • Have very high melting and boiling points (than metals in groups I and II).
    • They do not react with water;
    • Have very high densities (compared to metals in groups I and II)

Lanthanides and Actinides.

  • They form a block of elements within the transition metals.
  • Are sometimes called the inner transition metals.
  • The lanthanides consist of 14 elements from Cerium (Ce) to Lutetium (Lu).
  • The Actinides are the 14 elements from Thorium (Tho) to Lawrencium (Lr).

Placing an Element in the Periodic Table

  • The position of an elements in the periodic table is governed by the atomic number and hence the electron arrangement.

The Period:

  • The period to which an element belongs is determined by the number of energy levels.
  • The number of energy levels is equal to the period to which an elements belongs.

Examples:

Elements Symbol Atomic number Electron arrangement Number of energy levels Period
Lithium Li 3 2.1 2 2
Sodium Na 11 2.8.1 3 3
Calcium Ca 20 2.8.8.2 4 4
Nitrogen N 7 2.5 2 2
Helium He 2 2 1 1

Group:

  • The group to which an element belongs to is governed by the number of electrons in the outermost energy level.
  • The number of electrons in the outermost energy level is equal to the group to which the element belongs.

Examples:

Elements Symbol Atomic number Electron arrangement Outermost electrons Group
Potassium K 19 2.8.8.1 1 1
Aluminium Al 13 2.8.3 3 3
Silicon Si 14 2.8.4 4 4
Oxygen O 8 2.6 6 6
Chlorine Cl 17 2.8.7 7 7

Note: In the modern periodic table atomic masses are used instead of mass numbers. The atomic masses are preferable because they take care of elements with isotopes unlike mass numbers.

Diagram: part of the periodic table showing the first 20 elements

part of the periodic table



Relative Atomic Mass and Isotopes.

Introduction:

  • The masses of individual atoms of elements are very negligible and thus quite difficult to weigh.
  • On average the mass of an atom is approximately 10 -22 g which cannot be determined by an ordinary laboratory balance.
  • For this reason the mass of atom has been expressed relative to that of a chosen standard element hence the term relative atomic mass.
  • The initial reference element was hydrogen which was later replaced with oxygen.
  • Later the oxygen scale was found unsuitable because oxygen exists in several isotopes and thus led to problems when deciding the mass of an oxygen atom.
  • For this reason oxygen was replaced with carbon as the reference atom and to date relative atomic masses of elements are based on an atom of carbon-12 (note that carbon is isotopic and exists as Carbon -12 or carbon-14).

Definition:

  • Relative atomic mass (R.A.M) of an element refers to the average mass of an atom of the element compared with a twelfth (1/12) of an atom of carbon-12.

RAM = Average mass of one atom of an element/1/12 of an atom of carbon-12

Measurement of Relative Atomic Mass

  • RAM of elements is determined by an instrument called Mass Spectrometer.
  • The instrument can also be used to determine the relative abundance of isotopes.
  • The use of a mass spectrometer in determining the RAM of elements is called mass spectrometry.

How mass spectrometry works

  • In the mass spectrometer atoms and molecules are converted into ions.
  • The ions are then separated as a result of the deflection which occurs in a magnetic field.
  • Each ion (from an atom, isotope or molecule) gives a deflection which is amplified into a trace.
  • The height of each peak measures the relative abundance of the ion which gives rise to that peak.
    Note: Generally the relative atomic mass of an element is closest in value to the mass of the most abundant isotope of the element.

Example: Diagram of a Spectrometer Trace for Lithium

spectrometer trace for lithium

Explanations:

  • The trace has two peaks indicating that there are two isotopes for lithium.
  • The fist peak occurs at a relative isotopic mass of 6 and the second at 7; these are the RAM of the two isotopes respectively.
  • The percentage abundance of the isotope with RAM of 6 (6Li) is 9 while the RAM of the isotope with RAM 7 ( 7Li) is 91.

Calculating relative atomic masses of isotopic elements.

  • Information form a spectrometer trace is usually extracted and used in calculation the relative atomic mass of elements.

Questions

  1. The mass spectrum below shows the isotopes present in a sample of lithium.

    lithim mass spectrum

    1. Use this mass spectrum to help you complete the table below for each lithium isotope in the sample. (3 marks)

      Isotope  Percentage composition  Number of 
      Protons Neutrons
      6Li      
      7Li      
    2. Calculate the relative atomic mass of this lithium sample. Your answer should be given to three significant figures. (3 marks)
  2. Element X with atomic number 16 has two isotopes. ⅔ of 33X and ⅓ of 30X. What is the relative atomic mass of element X? (2 marks)
  3. Calculate the relative atomic mass of an element whose isotopic masses and relative abundances are shown below. (2 marks)

    Relative abundance Isotopic mass
    69 63
    31 65
  4. A neutral atom of silicon contains 14 electrons, 92% of silicon - 28, 5% silicon 29 and 3% silicon 30
    1. What is the atomic number of silicon? (1 mark)
    2. Calculate the relative atomic mass of silicon. (1mark)
  5. Oxygen exists naturally as isotopes of mass numbers 16, 17 and 18 in the ratio 96:2:2 respectively. Calculate its R.A.M (2 marks)
  6. Calculate the relative atomic mass of potassium from the isotopic composition given below.

    Isotope Relative abundance
    39K 93.1
    40K 0.01
    41K 6.89

     
  7.  Sulphur and sulphur compounds are common in the environment.
    1. A sample of sulphur form a volcano contained 88% by mass of 32S and 12% by mass of 34S.
      1. Complete the table below to show the atomic structure of each isotope of sulphur.

        Isotope  Number of  
        Protons Neutrons Electrons
        32S      
        34S      
      2. Define relative atomic mass. (2 marks)
      3. Calculate the relative atomic mass of the volcanic sulphur. (2 marks)
  8. Iridium, atomic number 77, is a very dense metal. Scientists believe that meteorites have deposited virtually all the iridium present on earth. A fragment of a meteorite was analysed using a mass spectrometer and a section of the mass spectrum showing the isotopes present in iridium is shown below.

    iridium mass spectrum
    1. Explain the term isotopes. (1mark)
    2. Use the mass spectrum to help you complete the table below for each iridium isotope in the meteorite.

      Isotope  Percentage composition  Number of 
      Protons Neutrons
      6Ir      
      7Ir      
      1. Define the term relative atomic mass. (1mark)
      2. Calculate the relative atomic mass of the iridium in this meteorite. (3 marks)


Ion Formation

Introduction

  • Atoms whose outermost energy levels contain the maximum possible number of electrons are said to be stable.
  • Thus atoms with energy levels 2, 2.8 and 2.8.8 are said to be stable.
  • Electron configuration 2 is said to have a stable duplet state while electron configuration 2.8 and 2.8.8 is said to have a stable octet state.
  • These electron configurations resemble those of noble gases and as such they are stable and do not react.
  • Atoms without this stability acquire it by either electron gain or electron loss.
  • Whether an atom loses or gains electron(s) depend on the number of electrons in the outermost energy level.

Illustration

  • Take the case of sodium.
  • Atomic number is 11 with an electron configuration of 2.8.1.
  • Thus sodium has two options in to become stable:
    • to lose the single electron and acquire a stable electron configuration of 2.8.
    • to gain 7 electrons in its outermost energy level and acquire a stable electron configuration of 2.8.8
  • Gaining a single electrons and losing a single electrons requires equal amounts of energy.
  • Thus it is cheaper and faster in terms of energy for sodium to lose the single electron in the outermost energy level than to gain 7 electrons into its outermost energy level.
  • Thus sodium acquires a stable electron configuration 2.8 by losing the single electron in its outermost energy level.

Diagram

sodium atom sodium atom

Equation:

Na → Na+ + e-

Further examples:

Element Electron arrangement Options for stability Best (cheapest) option
Chlorine 2.8.7 2.8 or 2.8.8 2.8.8
Potassium 2.8.8.1 2.8.8 or 2.8.8.8 2.8.8
Aluminium 2.8.3 2.8 or 2.8.8 2.8
Magnesium 2.8.2 2.8 or 2.8.8 2.8
Carbon 2.4 2 or 2.8 2. or 2.8
Oxygen 2.6 2. or 2.8 2.8

Ion

  • Definition : an ion is a charged particle of an element.
  • Are formed when an atom of an element either loses or gains electrons.

Illustration:

  • For a neutral atom the number of electrons in the energy levels (negative charges) is equal and thus completely balances the number of protons in the nucleus (positive charges).
  • Thus the net charge in a neutral atom is zero (0).
  • When an atom gains electron(s), the number of electrons becomes higher than the number of protons resulting to a net negative charge hence an ion.
  • Oppositely when an atom loses electron(s) the number of protons becomes higher than the number of electrons resulting into a net positive charge hence an ion.
  • The charge on the ion is usually indicated as a superscript to the right of the chemical symbol.
  • Thus ions are of two types:
    • Cations
    • Anions

Cations

  • Are positively charged ions.
  • Are formed when atoms lose electrons resulting into the number of protons being higher than the number of electrons.
  • Are mostly ions of metallic elements since most metals react by electron loss.

Examples:

  1. Magnesium:
    • It has atomic number 12, with electron arrangement 2.8.2.
    • It has 12 protons and 12 electrons hence a net charge of 0 hence the atom is written simply as Mg.
    • It will form its ions by losing the two electrons from the outermost energy level.
    • Thus the number of electrons decreases to 10 while the number of protons remains 12.
    • This leads to a net charge of +2, giving the ion with the formula Mg 2+.
      Diagrammatic illustration:
      magnesium atom and ion dot cross diagram
  2. Aluminium
    • It has atomic number 13, with electron arrangement 2.8.3.
    • It has 13 protons and 13 electrons hence a net charge of 0 hence the atom is written simply as Al.
    • It will form its ions by losing three (3) electrons out of the outermost energy level.
    • Thus the number of electrons decreases by three to 10 while the number of protons remains 13.
    • This leads to a net charge of +3, giving the ion with the formula Al+3 .
      Diagrammatic illustration:
      aluminum atom vs ion dot cross diagram

Anions:

  • Are negatively charged ions.
  • Are formed when atoms gain electrons resulting into the number of electrons being higher than the number of protons.
  • Are mostly ions of non-metallic elements since most non-metals ionize (react) by electron gain.

Examples

  1. Chlorine
    • It has atomic number 17, with electron arrangement 2.8.7.
    • It has 17 protons and 17 electrons hence a net charge of 0 hence the atom is written simply as Cl.
    • It will form its ions by gaining a single electron into the outermost energy level.
    • Thus the number of electrons increases to 18 while the number of protons remains 17.
    • This leads to a net charge of -1, giving the ion with the formula Cl- .
      Diagrammatic illustration:
      chlorine atom vs ion dot cross diagram
    • Note: the electros gained must be represented by a different notation from the initial electrons in the atom. E.g. if the initial electrons are represented with crosses(x) then the gained electrons should be represented by dots (.) and vise versa.
  2. Phosphorus
    • It has atomic number 15, with electron arrangement 2.8.5.
    • It has 15 protons and 15 electrons hence a net charge of 0 hence the atom is written simply as P.
    • It will form its ions by gaining three (3) electrons into the outermost energy level.
    • Thus the number of electrons increases by three to 18 while the number of protons remains 15.
    • This leads to a net charge of -3, giving the ion with the formula P-3 .
      Diagrammatic illustration:
      phosphorus atom vs ion dot cross diagram

 

Electron Transfer During Chemical Reactions

  • Atoms react either by electron gain or electron loss.
  • Generally metals react by electron gain while non-metals react by electron loss.

Illustration

  • Consider the reaction between sodium and chlorine.
  • Sodium attains stability // reacts by losing the single electron form its outermost energy level.
  • Chlorine attains stability // reacts by gaining a single electron into its outermost energy level.
  • Thus during the reaction between the two elements the single electron lost by the sodium atom to form the sodium ion is the same one gained by the chlorine atom to form the chloride ion.

Valence electrons

  • Refers to the number electrons in the outermost energy level.

Examples:

  • Calcium, with electron arrangement 2.8.8.2 has 2 valence electrons.
  • Oxygen with electron arrangement 2.6 has 6 valence electrons.
  • Phosphorus with electron arrangement 2.8.5 has 5 valence electrons.

Valency

  • Refers to the number of electrons an atom loses or gains during a chemical reaction.
  • Valency is also known as the combining power of an element.

Examples:

  • Calcium, with electron arrangement 2.8.8.2 loses 2 electrons during chemical reactions and hence has a valency of 2.
  • Oxygen with electron arrangement 2.6 gains 2 electrons during chemical reactions and thus has a valency of 2.
  • Phosphorus with electron arrangement 2.8.5 has gains 3 electrons during chemical reactions and hence has a valency of 3.
  • Aluminium, with electron arrangement 2.8.3 loses 3 electrons during chemical reactions and hence has a valency of 2.

Note: Some elements have variable valencies and are usually termed the transitional elements (metals)

Examples:

  • Iron can have valency 2 or 3;
  • Copper can have valency 1 or 2
  • Lead can have valency 2 or 4.

Summary on valencies of common elements

  Valency 1 Valency 2 Valency 3
Metals Sodium
Potassium
Calcium
Barium
Magnesium
Zinc
Iron
Lead
Copper
Aluminium
Iron
Non-metals Nitrogen
Chlorine
Fluorine
Hydrogen
Nitrogen
Oxygen
Sulphur
Nitrogen
Phosphorus


Radicals

  • Are groups of atoms with a net charge that exist and react as a unit during chemical reactions.
  • Radicals also have a valency, which is equivalent to the value of its charge.

Summary on valencies of some common radicals

  Valency 1 Valency 2 Valency 3
Radicals Ammonium (NH4+)
Hydroxide (OH-)
Nitrate (NO3- )
Hydrogen carbonate (HCO3-)
Hydrogen sulphate (HSO4-)
Carbonate (CO32-)
Sulphate (SO42-)
Sulphite (SO32-)
Phosphate (PO43-)

Oxidation Number

  • Refers to the number of electrons an atom loses or gains during a chemical reaction.
  • In writing the oxidation number the sign (+ or -) to show gain or loss is written followed by the number of electrons lost or gained respectively.

Illustration

  • Atoms are electrically neutral and are thus assigned an oxidation state of 0 since the number of protons in the nucleus is equal to the number of electrons in the energy levels.
  • However when atoms react they either lose or gain electrons and thus acquire a new state.
  • This new state is a new oxidation state and the atom thus acquires a new oxidation number.

Examples:

Atom E. arrangement Ion formula Valency Oxidation number
Sodium 2.8.1 Na+ 1 +1
Magnesium 2.8.2 Mg2+ 2 +2
Aluminium 2.8.3 Al3+ 3 +3
Nitrogen 2.5 N3- 3 -3
Sulphur  2.8.6 S2- 2 -2
Chlorine 2.8.7 Cl-1 1 -1

Further examples:

Particle Oxidation number
Copper metal, Cu 0
Lead (II) ion, Pb2+ +2
Bromide ion, Br- -1
Aluminium ion, Al2+ +2
Sulphide ion, S2- -2
Magnesium metal, Mg 0

Note: oxidation number (state) and charge of an element.

  • Oxidation state is written with the positive or the negative sign coming before the element. Examples: -2, 3, +1, -1 etc.
  • Charge on an element is write as a superscript of the element with the number coming before the positive r the negative sign
    Examples: Mg2+ , Al3+ , Na+ , Cl- etc.

Chemical Formulae

  • Refers to a representation of a chemical substance using chemical symbols.
  • In a single atom it is equivalent to the chemical symbol of the element.
  • In a compound it shows the constituent elements and the proportions in which they are combined.

Deriving the Chemical Formula of Compounds.

  • In order to write the correct formula of a compound the following must be known:
    • The symbols of the constituent elements or radicals.
    • The valencies of the elements or radicals
  • The chemical formula should start with the element which is more likely to lose electron (s) followed by the element that is more likely to gain.

Worked Examples

  1. Deriving the formula of sodium chloride.
    Elements Sodium Chlorine
    Formula Na Cl
    valencies 1 1
    balancing x1 x1
    Balancing ratios as subscripts : Na1Cl1
    Formula: NaCl
    Explanation:
    • For sodium to combine with chlorine to form sodium chloride, sodium loses an electron while chlorine gains an electron.
    • Thus every sodium atom needs only a single chlorine atom for both to be fully stable
      Note: When the balancing ratio // subscript is 1, it is usually not written since the symbol of the element itself represents a single atom.

  2. Deriving the formula of magnesium chloride.

    Elements Magnesium Chlorine
    Formula Mg Cl
    valencies 2 1
    balancing x1 x2

    Balancing ratios as subscripts : Mg1Cl2
    Formula: MgCl2
  3.  Deriving the formula of magnesium oxide.

    Elements Magnesium Oxygen
    Formula Mg O
    valencies 2 2
    balancing x1 x1
    Balancing ratios as subscripts : Mg1O1
    Formula: MgO

Questions.

  1. Derive the chemical formula of each of the following compounds.
    • Calcium fluoride
    • Carbon (II) oxide
    • Carbon (IV) oxide
    • Aluminium nitrate
    • Calcium hydrogen carbonate
  2. Complete the table below for elements A, B and C
    Element Valency Chemical formula of various compounds     
    hydroxides Sulphates carbonates nitrates phosphates Hydrogen carbonates
    A 1            
    B 2            
    C 3            

Chemical Equations

  • Refers to representations of a chemical reaction by means of chemical symbols and formula.

Key Features of a Chemical Equation

  • The correct formulae of the reactants are on the left of the equation.
  • The correct formulae of the products are on the right of the equation.
  • The reactants and products are separated by an arrow pointing to the right.
  • The state symbols of the reactants and products must be stated as subscripts to the right of the symbols
  • The number of each atom on the reactants side must be equal to the number of the same atom on the products side.

Example

Reaction between hot copper metals and oxygen gas.
✓ Word equation : Copper + oxygen → Copper (II) oxide.
✓ Chemical equation : 2Cu(s) + O2(g) → 2CuO(s)

Balancing Chemical Equations

  • A chemical equation is only valid if it is balanced.
  • A chemical equation is said to be balanced if the number of each atom on the reactants side is equal to that on the products side.
  • This is because atoms are neither created nor destroyed during a chemical reaction.

Rules and Guidelines in Balancing Chemical Equations

Step 1: Write the chemical equation in words.

  • Example: Copper metal + oxygen gas

Step II: Write the correct formulae of both reactants and products

  • Example: Cu + O2 → CuO

Step III: Check whether the number of atoms of each element on the reactants side is equal to that on the products side.

  • If equal proceed to step (V);
  • If not equal proceed to step (IV).
    Example: Cu + O2 → CuO
  • In this case there are two oxygen atoms on the reactants side yet there is only one oxygen atom on the products side. Thus we proceed to step IV

Step IV: Multiply the chemical formula containing the unbalanced atoms with the lowest common multiple.

  • Example: Cu + O2 → CuO
  • In this case the chemical formula with the unbalanced atom is CuO on the products side. We thus multiply it by 2.
  • The new equation now reads: Cu + O2 → 2CuO

Step V: check again to ensure that all atoms are balanced.

  • If all atoms are balanced proceed to step VI.
  • If not then repeat step IV until all atoms are balanced.
    Example: Cu + O2 → 2CuO
  • In this case multiplying CuO by 2 offsets the balancing of Cu; which is now unbalanced!
  • We therefore repeat step IV in order to balance Cu.
  • There is only 1 Cu atom on the reactants side yet there are 2 Cu atoms on the products side.
  • We thus multiply the formula with the unbalanced atom (s) by the lowest common multiple, in this case 2.
  • The new equation at this step thus becomes: 2Cu + O2 → 2CuO
  • We then repeat step V; in this case all atoms are now balanced.

Step VI: The physical states of the reactants and the products are then indicated.

  • If this is not done the chemical equation is considered incorrect.

Types of State Symbols.

  • There four main state symbols.
    1. Solid; denoted as (s)
    2. Liquid; denoted as (l)
    3. Aqueous (in solution in water); denoted as (aq)
    4. Gaseous; denoted as (g)
  • In a chemical equation the state symbols are written with their denotations as subscripts to the right of the chemical formulae.
    Example: 2Cu(s) + O2(g) → 2CuO(s)

Thus the balanced chemical equation for the reaction between copper metal and oxygen is:
✓ 2Cu(s) + O2(g) → 2CuO(s)

Exercise

  1. Balance equations for each of the following reactions.
    -Sodium hydroxide and dilute hydrochloric acid     
    -Zinc oxide and dilute sulphuric (VI) acid
    -Zinc metal and dilute nitric (V) acid                     
    -Calcium carbonate and dilute sulphuric (VI) acid
    -Sodium and water
  2. Balance each of the following equations.
    • Mg(s) + HCl (aq) →MgCl2(aq) + H2(g)                                             
    • Na(s) + H2O(l) → NaOH(aq) + H2(g)
    • NaOH(aq) + H2SO4(aq) →Na2SO4(aq) + H2O(l)
    • CuCO3(s) + HNO3(aq) → Cu(NO3)2(aq) + CO2(g) + H2O(l)
    • H2S(g) + O2(g) → O2(g) + H2O(l)                                                   
    • C2H6(g) + O2(g) → CO2(g) + H2O(l)
    • Pb(NO3)2(s) → PbO(s) + NO2(g) + O2(g)
    • Fe(s) + Cl2(g) → FeCl3(s)
    • Al(s) + H2SO4(aq) → Al2(SO4)3(aq) + H2(g) .

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