CHEMICAL BONDING IMPORTANT DEFINITIONS: 1. A charged particle that form from an atom (or a group of atoms) by the loss or Gain of electrons is called an Ion. 2. A positively charged ion formed when an atom loses electron(s) is called Cation. 3. A negatively charged ion formed when an atom gains electron(s) is called Anion. 4. A simple ion is an ion formed from single atom, e.g., Na. A polyatomic ion is an ion Containing more than one kind of atoms, e.g. NH 4 or SO 4 2- ion. 5. The strong electrostatic force of attraction between positive and negative ions in an ionic compound is called Ionic Bond. e.g. NaCl,MgCl 2 etc. 6. The bond formed by the sharing of electrons between two non-metal atoms is called Covalent Bond. e.g.- CH 4, CO 2 etc. 7. A molecule is a small group of atoms held together by covalent bonds. 8. The force of attraction between positive metal ions and the sea of delocalized electrons is called Metallic Bond. Test Yourself: Describe the formation of ions by electron loss/gain in order to obtain the electronic structure of a noble gas 1.1 The Stable Noble Gas Structure: 1. The elements, helium (He), Neon (Ne), argon(ar), Krypton (Kr), xenon (e) and radon (Rn) are known as noble gases. The electronic structures (Electronic configuration) of helium, neon and argon are shown below. 2. Helium has two valence electrons (a duplet electronic structure). All other noble gases have eight valence electrons (an octet electronic structure). 3. An atom is stable if it has a full outer shell, that is, a duplet or octet electronic structure. Only the noble gas atoms have this stable electronic structure.
4. The atoms of all other elements have incomplete outer shells. Hence, the atoms of these elements will donate, accept or share electrons to achieve the stable noble gas structure (duplet or octet configuration). Atoms gain or lose electrons to obtain noble gases configuration. ------------------------------------------------------------------------------------------------------------------------------- 1.2 Formation of Positive Ions: -------------------------------------------- 1. All atoms are electrically neutral because they have equal numbers of protons and electrons. 2. An ion is a charged particle formed when an atom loses or gains electrons. 3. Atoms of metals tend to lose valence electrons to form positive ions or cations. 4. The Positive ion contains more protons than electrons. Example: Formation of Sodium Ion, Na (a) Only the valence electron is removed from the outermost shell. (b) The Na ion contains eight valence electrons (noble gas Structure). (c) The Na ion carries one positive charge because it has 11 protons and only 10 Electrons. 1.3 Formation of Negative Ions: 1. Atoms of non-metals tend to gain electrons to form negative ions or anions. 2. A negative ion contains more electrons than protons. Example: Formation of Chloride ion, Cl - (a) An electron is added to the outermost shell. (b) The Cl - ion has the same electronic structure as argon. (c) The Cl - ion carries one negative charge because it has 18 electrons and only 17 protons.
Tip for Students: When drawing the electronic structure of an ion,the charge must be put at the top right hand corner. You will not be given any marks if you put the charge in the nucleus. Example 1.1: Write down the formulae for the ions of the following elements. Element Li Mg O Be P 1 Number of 3 12 8 4 15 9 Protons Number of 4 12 8 4 16 9 Neutrons Number of Electrons 3 10 10 2 18 10 Solution: Ignore the number of neutrons. Calculate the total charge in the ion. Element Li Mg O Be P F Number of 3 12 8 4 15 9 Protons Number of 3 10 10 2 18 10 Electrons Total 0 2-2 2-3 -1 Charge Formulae Li Mg 2 O 2- Be 2 P 3- F- 1.4 Ionic Bond Electron Transfer: 1. An ionic bond is formed when a metal reacts with a non-metal to form a compound. The compound formed is called the ionic compound. 2. Ionic bonds are formed by the transfer of electrons from the metal atoms to the non-metal atoms.
Example: (a) Sodium (Metal) reacts with chlorine (non-metal) by the transfer of electrons to form sodium chloride (ionic compound). (b) There are three parts involved in the formation of sodium chloride. 1: Each sodium atom loses its single valence electron to form a positively charged Na Ion. 2.Each chlorine atom gains an electron to form a negatively charged Cl-ion. Na Na e - (2,8,1) (2,8) Cl e- Cl - (2,8,7) 3. The Na ions and Cl- ions are attracted to each other by strong electrostatic attraction to form sodium chloride. This electrostatic attraction between positive and negative ions is called the ionic bond. An ionic bond may also be known as an electrovalent bond. Na Cl - NaCl The formation of an ionic bond can be represented by a dot and cross diagram. 3. An ionic bond can also be formed between a simple ion and a polyatomic ion, or between polyatomic ions. For example, ionic bonds can be formed in magnesium sulfate (Mg 2 SO 4 2 -) and ammonium nitrate (NH 4 NO 3 - ). 4. (a) The chemical formula of an ionic compound is constructed by balancing the charges on the Positive ions with those on the negative ions. (b)all the positive charges must equal all the negative charges in an ionic compound. (c) The Chemical formula of an ionic compound formed between the cation, M a and the anion, b - is:
M a b - M b a Example: Fe 3 O 2- Fe 2 O 3 5. The formulae of some ionic compounds are given in the table below. Cation Anion Formula of ionic Compound Na Cl - NaCl Na OH - NaHO K O 2 K 2O Na CO 2 Na 2 CO 3 Cation Anion Formula of ionic Compound Mg 2 NO - 3 Mg(NO 3 ) 2 Ca 2 OH - Ca(OH) 2 Mg 2 2- SO 4 MgSO 4 Al 3 - No 3 Al(NO 3 ) 3 Example 1.2: The formula of bismuth oxide is Bi 2 O 3 and calcium phosphate is Ca 3 (PO 4 ) 2. What is the charge on (a) bismuth ion? (b) phosphate ion? Analysis: An ionic compound is a neutral substance. Total positive charge = total negative charge Thus, for an ionic compound with the formula, (A n ) x (B m- ) y : x(n) y(-m) = 0 Solution: (a) Let the charge on bismuth ion be. (b) Let the charge on phosphate ion be. The charge on oxygen ion,o 2- = -2 The charge on Ca 2 = 2 2 ( ) = 3(2) 2 = 0 = =
Formula of bismuth ion : Bi 3 Formula of phosphate ion: PO 4 3- Analysis: Recognise that ionic compound contain a giant lattice in which the ions are held by electrostatic attraction. Deduce the formula and names of ionic compounds from their lattice structures. 1.5 Structure of ionic Compounds: Ionic compounds form giant ionic structures because the positive ions and negative ions are very strongly attracted to one another, the ions are arranged to form a giant ionic lattice or structure. Example: Sodium chloride is made up of Na and Cl - ions. In the solid state, these ions are very strongly attracted to one another. Analysis: Relate the physical properties of ionic compounds to their lattice structure 1.6 Properties of ionic compounds: Some of the physical properties of simple ionic substances are shown in the table below. Property Solubility: Ionic compounds are usually sol;uble in water, but insoluble in organic solvents like ethanol, petrol and turpentine. Explanation Water molecules can separate the positive ions From the negative ions, Causing them to dissolve. Ionic compounds such as silver chloride and calcium carbonate are
insoluble in water. Volatility: Ionic compounds have high melting and Boiling points. Electrical conductivity: Ionic compounds are non-conductors of electricity when molten or in aqueous solution. The forces of attraction between positive and negative ions are very strong. A large amount of heart energy is needed to overcome these strong bonds during melting or boiling. In the solid state, the ions are not free to move about. In the molten state or in aqueous solution, the ions can move freely to conduct electricity. Tip for Students: it is the ions (and not electrons) that conduct electricity in molten or aqueous ionic compounds. You will get no mark if you say sodium chloride conducts electricity when molten because the ions and the electrons can move : The incorrect use of the word electrons negates the mark for using the word ions. Test yourself: Describe the formulation of a covalent bond by the sharing of a pair of electrons Deduce the arrangement of electrons in covalent molecules 1.7 Covalent Bond Sharing of Electrons: 1. A covalent bond is the chemical bond formed by the sharing of electrons between two non-metal atoms. After bonding, each atom attains the electrons the electronic structure of a noble gas. 2. When atoms combine by sharing electrons, molecules are formed. A molecule is a group of two or more Atoms held together by covalent bonds.
Compounds such as hydrogen chloride (HCl) and carbon dioxide (CO 2 ), which contain covalent bonds, are called Covalent compounds. 3. Covalent bonds can be formed between (a) atoms of the same element ( e. g. hydrogen, H 2, and chlorine, Cl 2 ), (b) atoms of different elements (e. g. ammonia, NH 3, and water,l H 2 O). 4. Covalent bonds in molecules of elements (a) A single covalent bond or a single bond is formed when a pair of electrons are shared between two atoms. It is represented by a single line - in the structural formula. Example: Hydrogen molecule, H 2 H xh H x H (or H H) Two hydrogen atoms One hydrogen molecule By sharing their valence electrons each hydrogen atom achieves the stable electronic structure by sharing their valence electrons, each hydrogen atom achieves the stable electronic structure. (b) A double colavent bond or a double bond is formed when two pairs of electrons are between two atoms. It is represented by = in the structure formula. Example: Oxygen molecule, O 2 O O O O An oxygen atoms has six valence electrons. It needs two more electrons to have the stable octet electronic structure. Each oxygen molecule contains a double bond. O O O 2 (or O =O)
5. Covalent bonds in molecules of compounds (a) Molecules made from two or more different types of atoms linked together by covalent bonding are called molecular compounds or covalent compounds. (b) Water, methane and carbon dioxide are example of covalent compounds. Example: Water molecules, H 2 O O H H H H O O 2H H 2 O ( H O H ) Water is formed by the reaction of hydrogen with oxygen such that all there atoms attain noble gas structures. Each water molecule contains two single bonds. Methane contains two elements carbon and hydrogen. In a methane molecule, the carbon atoms has an octet structure while each hydrogen atoms has a duplet structure. Each methane molecule contains four single bonds. Methane molecule, CH 4 Carbon dioxide is formed when carbon reacts with oxygen. The carbon atom must share two electrons each with two oxygen atoms in order to achieve a stable octet of a electrons. Each oxygen atoms shares two electrons and attains an octet structure.
H H C 4H H C H ( or H C H ) H H Each carbon dioxide molecules contains two double bonds. Carbon dioxide molecules,co 2 7. An ionic compounds can contain covalent bonds,but a covalent compound cannot contain ionic bonds. Example: ammonium chloride is an ionic compound that contains covalent bonds. The bonding between NH 4 and Cl - ions is an ionic bond. The bonding between N and H atoms in NH 4 is a covalent bond. Tip for Students: Compounds of metals with non-metals are usually ionic but compounds of non-metals with non-metals (same element or different elements) are always covalent.
Relate the physical properties of covalent substances to their structure and bonding 1.8 Properties of Covalent Substances: 1. In general,covalent substances are insoluble in water but soluble in organic solvents. 2. There are exception to ths rule. For example, hydrogen chloride (HCl), ethanol (C 2 H 5 OH ) and ethanoic acid (CH 3 COOH ) are covalent substances that are soluble in water. 3. Other properties of simple covalent substances are shown in the table below. Test Yourself: Property Low melting and boiling points Do not conduct electricity During the melting of simple covalent molecules, heat energy is absorbed to break the covalent bonds between the molecules. Simpe covalent molecules have low melting and boiling points because covalent bonds are weak. Explanation The covalent molecules are held together by weak intermolecular forces of attraction. Only a small amount of heat energy is needed to overcome these weak forces of attraction. Covalent substances do not contain positive and negative ions that can move about freely. Hence, covalent substances are non-conductors of electricity (except graphite) During the melting of simple covalent molecules, heat energy is absorbed to overcomes the attractive forces between molecules. The covalent bonds within the molecules are strong. Simple covalent molecules have low melting and boiling points because the attractive forces between the covalent molecules are weak. Compare the structure of simple molecular substances with those of giant molecular substances in order to deduce their properties. Deduce the properties of Diamond, Graphite and Silicon(IV) Oxide from their bonding and structures.
Compare the physical properties of ionic compounds with of covalent molecules 1.9 Types of Covalent Substances: 1. Covalent substances can be divided into two categories as shown in the table below. Simple covalent molecules Examples: hydrogen, chlorine, oxygen, water, methane and carbon dioxide Macromolecules Example: diamond, graphite and sand (silicon dioxide) 2. Simple molecular substances: (a) These are made up of independent molecular units. As there are on ions formed, the forces of attraction between molecular in solid, covalent compounds like iodine and sulfur are much weaker. They are called vander Waals forces and produce a weak molecular lattice with low melting points. (b) In covalent liquids like water, the molecules are even further apart, so the vander Waals forces are weaker still. In covalent gases like methane and ammonia,these forces are almost non-existent. 3. Giant molecular substances: (a) Molecules with giant structures are called macromolecules. (b) The structure, properties and uses of diamond, graphite and silicon(iv) oxide are shown below.
Giant molecular structure Diagrammatical representation Properties Diamond Graphite Silicon(IV) oxide Three valence electrons in each carbon atom are used for covalent bonding and the fourth valence electron is not involved in chemical bonding. This gives hexagonal rings of six atoms that join together to form flat layers. These layers of carbon atoms, lie on top of each and are held together by vander Waals forces. Each carbon atom has four valence electrons. Each carbon atoms is joined to four other carbon atoms by strong covalent bonds in a tetrahedral arrangement. Hard Very high melting and boiling points Non-conductor of electricity Soft and slippery high melting and boiling points Good conductor of electricity Each silicon atom has four valence electrons and forms four covalent bonds with four oxygen atoms in a tetrahedral structure. each oxygen atom is bonded to two silicon atoms. High melting and boiling points Non-conductor of electricity Uses As gemstones As tips of grinding, cutting and polishing tools In pencils As a dry lubricant Brushes for electric motor As inert electrodes Main component of concrete and glass production As an abrasive in sandblasting and in media filters for filtering water 4. The table below compares the physical properties of ionic compounds and covalent substances:
Property Test yourself: Ionic compound Example: Sodium chloride Covalent substances Simple molecules Giant molecule Example :Iodine and Example: Diamond and carbon dioxide silicon dioxide Volatility Non-volatile Very volatile Non-volatile Melting and boiling High Low High points Solubility in water Soluble (there are some exception) Insoluble (there are some exception) Insoluble solubility in organic solvents Electrical conductivity Insoluble Soluble Insoluble Non-conductor in solid state Conductor electricity in molten state or aqueous solution Non-conductor in solid and liquid states (except graphite) Describe metals as a lattice of positive ions in a sea of electrons Define metallic bond and relate the electrical conductivity of metals to the mobility of the electrons in the structure 1.10 Metallic Bond: 1. In a metal, the atoms are packed tightly together in a regular pattern. 2. Metals consist of a lattice of positive ions surrounded by a sea of electrons. 3. This sea of electrons comes from the valence electrons of each atom. They are said to be delocalised. They move freely among the metal ions like a cloud of negative charge. 4. Therefore, metallic bond is the force of attraction between positive metal ions and the sea of delocalised electrons. 5. The lattice structure of the metallic bond is shown in the diagram below. 6. Properties of metals: (a) Electrical conductivity: (i) Metals are good conductors of electricity in the solid and molten states. This is due to the sea of electrons that can move through the crystal lattice of metals.
(ii) When a metal is used in an electrical circuit, electrons entering one end of the metal cause a similar number of electrons to be displaced from the other end. (iii) The valence electrons move from the negative terminal of the electrical circuit and hence the metal is able to conduct an electric current. Tips for Students: The electrical conductivity in metals is due to the moving valence electrons. You should not say that it is due to the movement of electrons as it could Include other electrons also in the metal atoms and not just the valence electrons. TERMINAL FLOW OF ELECTRONS TERMINAL ( - ) ( ) FREE ELECTRONS METAL ION ELECTRICAL CONDUCTIVITY OF METALS (b) Malleability and Ductility: Metals are malleable (can be bent or flattened) and ductile (can be drawn into wires) Because the layers of metal ions can slide over each other when a force is applied.
The effect of a force being applied on a metal lattice can be seen in the diagram below. METAL ION VALENCE ELECTRON APPLIED FORCE SLIP PLANE (c) Melting and boiling points: (i) Metallic bonds are strong bonds. Hence, metals have high melting and boiling points. (ii) There are few exceptions. For example, mercury has a low melting point, and is a liquid at room temperature. Similarly, sodium and potassium have low melting and and boiling points. The melting points of sodium and potassium are 98 C and 64 C respectively. Test yourself: Deduce the physical and chemical properties of substances from their structures and bonding and vice versa. 1.11 Deducing Bonding and Structure of a Substance: We can deduce the bonding and structure of a substance from its physical properties as shown in the in the flow chart below.
Yes Metal (or graphite) Substances Does it conduct electricity in in the solid state? No Ionic compound or covalent molecule Yes Ionic compound Does it conduct electricity in the Liquid state? No Covalent molecule Is it melting point high or low Simple covalent Giant covalent molecule STRUCTURAL QUESTIONS AND ANSWERS 1. (a) Draw a dot and cross diagram of the electronic structure of carbon dioxide. Include all its electrons. (b)using carbon dioxide as an example, explain what is meant by covalent bonds. Answer:
1. a) : Electrons of Carbon : Electrons of Oxygen C b) For carbon dioxide, electrons are shared between carbon and oxygen. The pairs of shared electrons from covalent bonds. 2.) Element Q is found in Group VI of the Periodic Table while element R is found in Group VII of the Periodic Table. Element Q reacts with element R to give a covalent compound. (a) Draw a dot and cross diagram to show the arrangement of electrons in one molecule of the compound formed. (b) Write the molecular formula of the compound formed. Answer: 2. a) : Electrons Of Q : Electrons Of R
Q R R b) QR2 3.)Explain each of the following as fully as you can. Answer: (a)hydrogen sulfide has a low boiling point of 60. (b)iodine does not conduct electricity. 3.)a) Hydrogen sulfide has a simple molecular structure in which the hydrogen sulfide molecules are held together by weak intermolecular forces. A small amount of energy is needed to overcome the weak intermolecular forces (holding discrete hydrogen sulfide molecules together ). Hence, hydrogen sulfide has a low boiling point. b) Iodine has a simple molecular structure in which the iodine molecules are heldtogether by weak intermolecular forces. Iodine does not conduct electricity because of the absence of mobile charged particles
(such as ions or electrons ). Answer: 4.) Hydrogen exists as three isotopes: hydrogen, deuterium and tritium. (a) Using D to represent deuterium, construct a dot and cross diagram to show the bonds in the oxide of deuterium. (b) Using T to represent tritium, construct a dot and cross diagram to show the arrangement of electrons in a tritium molecule. Answer: 4.) a) : Electrons of Oxygen : Electrons of Deuterium D D b) T T
5.) (a) Write the formula of a chloride ion. 5.) a) Cl - (b) Draw a diagram to illustrate the formation of a chloride ion from a chlorine atom. Show the electronic structures of both particles. Answer: b) Gains An 17p 18n Electron 17p 18n Cl Cl - 6.) (a) Draw a dot and cross diagram of the electronic structure of magnesium oxide.include all its electrons. (b) Using magnesium oxide as an example, explain what is meant by ionic bonds.
Answer: 6.) a) 2 2- Mg O Mg 2 O 2- b) In magnesium oxide, two is the electrons from a magnesium atom are transferred to an oxygen atom. The ionic bond is the electrostatic force of attraction between the positively charged magnesium ion and the negatively charged oxide ion. 7.) Explain each of the following as fully as you can. (a) Magnesium oxide, which has a melting point of 2800, is used as a refractory material. (b) Sodium chloride conducts electricity in the molten state but does not conduct electricity in the solid state. (c) Calcium fluoride is a brittle material. Answer: 7.) a) Magnesium oxide has a giant ionic structure ( lattice ) in which the positively charged magnesium ions ( Mg 2 ) and the negatively charged oxide ions ( O 2- ) are held together by strong electrostatic attraction
( ionic bonds ). A lage amount of energy is needed to overcome the strong ionic bonds holding the Mg 2 and O 2- ions together. Hence, magnesium oxide has a high melting point and can be used as a refractory material. b) In the solid state, sodium chloride has a giant ionic structure ( lattice ) in which the positively charged sodium ions and the negatively charged chloride ions are held together by strong electrostatic attraction. These ions are held in fixed positions and are not fro to move about. Thus, sodium chloride cannot conduct electricity in the solid state. In the molten state, electrostatic attraction is slightly weaker and the ions are free to move about (slide around one another ). Thus, sodium chloride in the molten state conducts electricity because of the presence of mobile ions. c) Calcium fluoride has a giant ionic structure ( lattice ) in which the positively charged calcium ions and the negatively charged fluoride ions are held together by strong electrostatic attraction ( ionic bonds ). When a force is applied, a slight displacement in the lattice structure occurs and brings similarly charged ions together. Repulsion between the like charges fractures the lattice of calcium fluoride easily. Hence, calcium fluoride is a material. 8.)Explain each of the following as fully as you can. (a) Iron has a high melting point of 1538. (b) Aluminium is malleable and ductile. (c) Copper is a good conductor of electricity. Answer: 8.) a) Iron has a giant metallic structure in which a lattice of positive ions and delocalized electrons are bound together by strong metallic bonds. A large amount of energy is needed to overcome the strong metallic bonds in iron. Hence, iron has a high melting point. b) Aluminium has a giant metallic structure in which a lattice of positive
ions and delocalized electrons are bonded by strong melting bonds. The layers of positive ions can slide over each other without breaking the metallic bonds. After being beaten into different shapes or drawn into wires, the neighbouring positive ions are still held together by the Sea of electrons. c)copper has a giant metallic structure in which a lattice of positive ions and delocalised electrons are bound together by strong metallic bonds. Copper is able to conduct electricity because of the presence of the sea of mobile delocalised electrons.