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Ionic bonding
- 1. CHAPTER 6 Ionic Bonding © 2013 Marshall Cavendish International (Singapore) Private Limited 1
- 2. Chapter 6 Ionic Bonding 6.1 The Stable Electronic Configuration of a Noble Gas 6.2 Forming Ions 6.3 Ionic Bond: Transferring Electrons 6.4 Chemical Formulae of Ionic Compounds 6.5 Structure and Physical Properties of Ionic Compounds 2
- 3. 6.1 The Stable Electronic Configuration of a Noble Gas Learning Outcome At the end of this section, you should be able to: • describe the stable electronic configuration of a noble gas. 3
- 4. 6.1 The Stable Electronic Configuration of a Noble Gas What are Noble Gases? • Elements that belong to Group 0 of the Periodic Table • Examples: He, Ne, Ar, Kr and Rn • Atoms of noble gases are stable and unreactive. • They exist in nature as single atoms. 4
- 5. 6.1 The Stable Electronic Configuration of a Noble Gas What is the Noble Gas Structure? • Noble gases have full or complete outer shells. Helium has a duplet configuration (2 outer electrons). All other noble gases have an octet configuration (8 outer electrons). 5
- 6. 6.1 The Stable Electronic Configuration of a Noble Gas Why Do Atoms React? • Atoms of most other elements are reactive because they do not have the noble gas structure (i.e. their outer shells are not fully-filled). • Atoms of these elements lose, gain or share outer electrons to attain the noble gas configuration and form compounds. 6
- 7. 6.1 The Stable Electronic Configuration of a Noble Gas Chemical Bonding Atoms gain or lose electrons to attain noble gas configuration Ionic bonding Atoms share electrons to attain noble gas configuration Covalent bonding 7
- 8. Chapter 6 Ionic Bonding 6.1 The Stable Electronic Configuration of a Noble Gas 6.2 Forming Ions 6.3 Ionic Bond: Transferring Electrons 6.4 Chemical Formulae of Ionic Compounds 6.5 Structure and Physical Properties of Ionic Compounds 8
- 9. 6.2 Forming Ions Learning Outcome At the end of this section, you should be able to: • describe the formation of positive ions (cations) and negative ions (anions) to achieve the noble gas configuration. 9
- 10. 6.2 Forming Ions What is an Ion? Recall: Atoms have an equal number of protons and electrons. They are electrically neutral. • An atom loses or gains electrons to form ions. • Ions are charged particles. No. of electrons ≠ No. of protons 10
- 11. 6.2 Forming Ions What is an Ion? • • Positively-charged ions are called cations. • URL Ions can be positively- or negatively-charged. Negatively-charged ions are called anions. 11
- 12. 6.2 Forming Ions Formation of Cations • Atoms of metals lose electrons to form positively-charged ions called cations. •In this way, they achieve the noble gas configuration. 12
- 13. 6.2 Forming Ions Example 1: Formation of sodium (Na+) ion Na atom Electronic configuration: 2, 8, 1 Number of protons = 11 Number of electrons = 11 The Na atom loses one outer electron to form the Na+ ion. Why? To achieve stable octet (noble gas) configuration. Neon (2, 8) 13
- 14. 6.2 Forming Ions Example 1: Formation of sodium (Na+) ion 2, 8 2, 8, 1 sodium atom loses one outer electron Na atom: 11p, 12n, 11e Charge = 11p + 11e = (+11) + (–11) =0 Neutral Na atom + Na+ ion: 11p, 12n, 10e Charge = 11p + 10e = (+11) + (–10) = +1 Positively-charged Na+ ion 14
- 15. 6.2 Forming Ions Example 2: Formation of calcium (Ca2+) ion 2, 8, 8 2, 8, 8, 2 calcium atom loses two outer electrons Ca atom: 20p, 20n, 20e Charge = 20p + 20e = 20(+1) + 20(–1) = (+20) + (–20) = 0 Neutral Ca atom 2+ Ca2+ ion: 20p, 20n, 18e Charge = 20p + 18e = 20(+1) + 18(–1) = (+20) + (–18) = +2 Positively-charged Ca2+ ion 15
- 16. 6.2 Forming Ions Common Cations and Their Charges Metal Ion Formula of ion sodium sodium ion Na+ potassium potassium ion K+ calcium calcium ion Ca2+ magnesium magnesium ion Mg2+ aluminium aluminium ion Al3+ 16
- 17. 6.2 Forming Ions Formation of Anions • Atoms of non-metals gain electrons to form negatively-charged ions called anions. •In this way, they achieve the noble gas configuration. 17
- 18. 6.2 Forming Ions Example 1: Formation of chloride (Cl–) ion Cl atom Electronic configuration: 2, 8, 7 Number of protons = 17 Number of electrons = 17 What happens in the formation of a chloride ion? The chlorine atom gains one electron in its outer shell to achieve a stable octet (noble gas) configuration. Argon (2, 8, 8) 18
- 19. 6.2 Forming Ions Example 1: Formation of chloride (Cl–) ion 2, 8, 7 2, 8, 8 chlorine atom gains one electron Cl atom: 17p, 18n, 17e Charge = 17p + 17e = (+17) + (–17) =0 Neutral Cl atom Cl– ion: 17p, 18n, 18e Charge = 17p + 18e = (+17) + (–18) = –1 Negatively charged Cl– ion 19
- 20. 6.2 Forming Ions Example 2: Formation of oxide (O2–) ion 2, 6 oxygen atom gains two electrons 2, 8 2– O atom: 8p, 8n, 8e O2– ion: 8p, 8n, 10e Charge = 8p + 8e = (+8) + (–8) =0 Charge = 8p + 10e = (+8) + (–10) = –2 Neutral O atom Negatively charged O2– ion 20
- 21. 6.2 Forming Ions Common Anions and Their Charges Non-metal Ion Formula of ion chlorine chloride ion Cl– bromine bromide ion Br– oxygen oxide ion O2– sulfur sulfide ion S2– 21
- 22. Why do metals lose electrons to form positive ions (cations) but non-metals gain electrons to form negative ions (anions)? 22
- 23. Chapter 6 Ionic Bonding 6.1 The Stable Electronic Configuration of a Noble Gas 6.2 Forming Ions 6.3 Ionic Bond: Transferring Electrons 6.4 Chemical Formulae of Ionic Compounds 6.5 Structure and Physical Properties of Ionic Compounds 23
- 24. 6.3 Ionic Bond: Transferring Electrons Learning Outcome At the end of this section, you should be able to: • describe how an ionic bonds are formed between metals and non-metals. 24
- 25. 6.3 Ionic Bond: Transferring Electrons Ionic Bonding • Ionic bonds are formed between metals and non-metals. Examples: Group VII: Fluorine, chlorine Group VI: Oxygen, sulfur Examples: Group I: Sodium, potassium Group II: Magnesium, calcium • This is done through the transfer of electron(s) from metals to non-metals. 25
- 26. 6.3 Ionic Bond: Transferring Electrons Ionic Bonding Metallic atom loses electron(s) Non-metallic atom gains electron(s) Positive ion Negative ion (cation) (anion) electrostatic forces of attraction (hold oppositely charged ions together) 26
- 27. 6.3 Ionic Bond: Transferring Electrons Formation of Ionic Compound Example 1: Sodium chloride Step 1: Formation of Positive Ions Each sodium atom (Na) loses its single outer electron to form a positively-charged sodium ion (Na+). Na 2, 8, 1 Na+ + e− 2, 8 27
- 28. 6.3 Ionic Bond: Transferring Electrons Step 2: Formation of Negative Ions Each chlorine atom gains an electron from a sodium atom to form a negatively-charged chloride ion (Cl−). Cl 2, 8, 7 + e− Cl – 2, 8, 8 28
- 29. 6.3 Ionic Bond: Transferring Electrons Step 3: Formation of Ionic Bonds Loses one electron Sodium atom 2, 8, 1 Gains one electron Chlorine atom 2, 8, 7 Electrostatic forces of attraction Sodium ion Chloride ion 2, 8, 8 2, 8 Sodium and chlorine react in the ratio of 1 : 1 to form sodium chloride (NaCl). URL 29
- 30. 6.3 Ionic Bond: Transferring Electrons Example 2: Magnesium chloride Magnesium atom loses two electrons. Chlorine atoms gain one electron each. Chloride ion Magnesium Chloride ion 2, 8, 8 2, 8, 8 ion 2, 8 Magnesium reacts with chlorine in the ratio of 1 : 2 to form magnesium chloride (MgCl2). 30
- 31. Chapter 6 Ionic Bonding 6.1 The Stable Electronic Configuration of a Noble Gas 6.2 Forming Ions 6.3 Ionic Bond: Transferring Electrons 6.4 Chemical Formulae of Ionic Compounds 6.5 Structure and Physical Properties of Ionic Compounds 31
- 32. 6.4 Chemical Formulae of Ionic Compounds Learning Outcome At the end of this section, you should be able to: • deduce the chemical formula of an ionic compound from the charges on the ions and vice versa. 32
- 33. 6.4 Chemical Formulae of Ionic Compounds Chemical Formulae of Ionic Compounds • The formula of an ionic compound is constructed by balancing the charges on the positive and negative ions. • All the positive charges must equal all the negative charges in an ionic compound. 33
- 34. 6.4 Chemical Formulae of Ionic Compounds Example: Magnesium oxide Oxygen forms O2− ions. Magnesium forms Mg2+ ions. Mg2+ O2− Charge: +2 Charge: −2 Since 1 × (+2 charge) balances out 1 × (−2 charge), The formula is MgO. 34
- 35. 6.4 Chemical Formulae of Ionic Compounds Example: Copper(II) hydroxide Copper ion Hydroxide ion Cu2+ Charge: +2 OH− Charge: −1 To balance the charges, multiply the smaller charge (−1) by 2 to make it equal to +2. Since 1 × (+2 charge) balances out 2 × (−1 charge), The formula is Cu(OH)2. 35
- 36. 6.4 Chemical Formulae of Ionic Compounds Example 1 Write the chemical formula of aluminium oxide. oxide ion aluminium ion Al 3 + O 2− Charge: −2 Charge: +3 Al2O3 Therefore, the formula is Al2O3. 36
- 37. 6.4 Chemical Formulae of Ionic Compounds Example 2 Write the chemical formula of calcium carbonate. calcium ion Ca 2 + carbonate ion CO3 2 − Charge: +3 Ca2(CO3)2 Charge: −2 CaCO3 Since ‘2’ is a common factor, it can be removed. Therefore, the formula is CaCO3. URL 37
- 38. Chapter 6 Ionic Bonding 6.1 The Stable Electronic Configuration of a Noble Gas 6.2 Forming Ions 6.3 Ionic Bond: Transferring Electrons 6.4 Chemical Formulae of Ionic Compounds 6.5 Structure and Physical Properties of Ionic Compounds 38
- 39. 6.5 Structure and Physical Properties of Ionic Compounds Learning Outcomes At the end of this section, you should be able to: • state that ionic compounds form giant lattice structures; • deduce the formulae of ionic compounds from their lattice structures; • relate the physical properties of ionic compounds to their lattice structures. 39
- 40. 6.5 Structure and Physical Properties of Ionic Compounds Structure of Ionic Compounds Ionic compounds form giant ionic structures. Also known as giant lattice structures or crystal lattices Consist of an endlessly repeating three-dimensional lattice of positive and negative ions Ions are closely packed, arranged in an orderly manner and held in place by ionic bonds 40
- 41. 6.5 Structure and Physical Properties of Ionic Compounds Structure of NaCl Sodium chloride crystal Three-dimensional arrangement of sodium ions and chloride ions Sodium ions and chloride ions alternate with each other. 41
- 42. 6.5 Structure and Physical Properties of Ionic Compounds Structure of NaCl Strong forces of attraction between ions in crystal lattice Na+ Cl – Na+ Cl– A large amount of energy is required to overcome these forces of attraction between ions. Cl– Na+ Na+ Cl– Na + Na+ Cl– Cl– Na+ Na+ Cl– Cl– Na+ Cl – Cl– Na+ Na+ Cl– Na+ Cl – Na + Na+ Cl– Cl– Cl– Na+ Na Na+ Cl– + Na+ Cl– Cl– Na+ Na+ Cl– Na+ Cl– Na+ Cl– Cl– Na+ Cl– Na+ Cl– Na+ 42
- 43. 6.5 Structure and Physical Properties of Ionic Compounds Structure of NaCl Cl− ion Each chloride ion is surrounded by six sodium ions. Na+ ion Each sodium ion is surrounded by six chloride ions. The ratio of sodium ions to chloride ions is 1 : 1. Hence, the formula unit of sodium chloride is NaCl. 43
- 44. 6.5 Structure and Physical Properties of Ionic Compounds Melting and Boiling Points of Ionic Compounds • High melting and boiling points Na+ Cl – Na+ • Non-volatile Cl– Cl– Na+ Na+ Cl– Na + Na+ Cl– • Exist as solids at room temperature Cl– Na+ Na+ Cl– Cl– Na+ Cl – Cl– Na+ Na+ Cl– Na+ Cl– Na + Na+ Cl– Cl– Cl– Na+ Na Na+ Cl– + Na+ Cl– Cl– Na+ Na+ Cl– Na+ Cl– Na+ Cl– Cl– Na+ Cl– Na+ Cl– Na+ 44
- 45. 6.5 Structure and Physical Properties of Ionic Compounds Solubility of Ionic Compounds • Usually soluble in water Cl– Na+ Cl– Na+ Cl – Na + Cl dissolve in water Na+ – Na+ Cl– Na+ Na+ Water molecules • URL Cl– Usually insoluble in organic solvents E.g. ethanol, turpentine, petrol 45
- 46. 6.5 Structure and Physical Properties of Ionic Compounds Electrical Conductivity of Ionic Compounds solid NaCl aqueous NaCl molten NaCl 46
- 47. 6.5 Structure and Physical Properties of Ionic Compounds Electrical Conductivity of Ionic Compounds • Ionic compounds conduct electricity in the molten and aqueous states. • They do not conduct electricity in the solid state. • In the molten and aqueous states, mobile ions are present. • Mobile ions conduct electricity. 47
- 48. Chapter 6 Ionic Bonding Concept Map 48
Hinweis der Redaktion
- Provide a brief introduction to the idea of noble gases. Have students recall from Chapter 4 that neon, which is a noble gas, is monoatomic. Have students recognise that noble gases exist in the monoatomic state.
- State that noble gases are inert, i.e. their atoms do not combine with other atoms. Have students understand that the 8 electrons in the outer shell (2 in the case of helium) makes the atoms stable and, therefore, inert. A duplet or octet configuration is also known as a noble gas configuration. Ask students why atoms of other elements react with one another. Lead them to understand that atoms of other elements react in order to achieve the stable noble gas structure.
- Briefly introduce the two types of bonding.
- Define Ions as charged particles formed when atoms lose or gain electrons.
- Some students may have the misconception that when an ion gains an electron, it becomes positive and vice versa. Highlight to them that electrons are negatively charged, so when an atom gains an electron, a negative ion is formed. Clicking on the URL button will link you to <http://www.bbc.co.uk/schools/gcsebitesize/science/add_aqa/atomic/ionicact.shtml>, a website with an interactive animation introducing the formation of ions and ionic bonds. (The video is approximately 3.5 minutes long.)
- Mention that cations have a stable octet configuration in their outer shells.
- Ask students to identify the noble gas configuration achieved by the sodium ion in its outer shell.
- Ask students if they can identify the noble gas configuration of each ion, i.e. sodium ion has neon’s electronic configuration, calcium ion has argon’s electronic configuration. Highlight to students that although the ions have the same electronic configuration as noble gases, they are not the same. For example, although sodium ion has the same configuration (2, 8) as neon, they are not identical particles because sodium ion has 11 protons, whereas neon has 10 protons.
- Mention that anions have a stable octet configuration in their outer shells.
- Ask students to identify the noble gas electronic configuration achieved by the chloride ion in its outer shell. Emphasise that chlorine forms a chloride ion, not a chlorine ion.
- Ask students if they can identify the noble gas configuration of each ion, e.g. chloride ion has argon’s electronic configuration.
- Ask students: What is another name for positive ions? What is another name for negative ions? Get students to recall that a metallic atom loses electron(s) to form a positive ion while a non-metallic atom gains electron(s) to form a negative ion. The oppositely charged ions are then held together by strong electrostatic forces of attraction. Recap: Why do atoms form ions? Answer: To achieve stable noble gas electronic configuration
- Get students to identify the metal and non-metal elements in sodium chloride. Sodium is the metal, chlorine is the non-metal.
- Clicking on the URL button will link you to <http://www.youtube.com/watch?v=QqjcCvzWwww&NR=1>, a website with a video on chemical bonding. The first part of the video (up till 0.51 seconds) can be used to demonstrate ionic bonding. The second part of the video (0.52 seconds onwards) may be used to demonstrate covalent bonding. Alternative resource: Go to <http://www.bbc.co.uk/schools/gcsebitesize/science/add_aqa/atomic/ionicrev4.shtml>, a website with an animation on ionic bonding of sodium chloride, magnesium oxide and calcium chloride.
- Explain that although ions are charged particles, the compounds they form are neutral. Thus, we can infer that the positive and negative charges must be balanced.
- When writing chemical formulae of ionic compounds, the metal is usually written first. When there is only 1 of each ion, the subscript ‘1’ is not written. E.g. MgO and not Mg1O1
- Highlight to students that polyatomic ions need to be enclosed within brackets when there is a subscript (in the case where there is more than one polyatomic ion).
- Usually, the simplest set of whole numbers is written, e.g. CaCO3 and not Ca2(CO3)2. However, there are exceptions, e.g. H2O2 and N2O4. Clicking on the URL button will link you to <http://www.learner.org/interactives/periodic/groups_interactive.html>, a website with an interactive game on ionic bonding. Students will get to match different ions to form compounds and see how well they score. (Note: There is an error in the description of aluminium phosphate – its chemical formula should be AlPO4, not AlPO2.)
- The giant ionic structure is held together very tightly due to the strong electrostatic forces between the oppositely charged ions.
- Another representation of the 3-D arrangement of Na+ and Cl− ions
- Ionic compounds have a giant lattice structure, which is held very tightly by strong attractive forces. A large amount of energy is required to overcome these forces to change an ionic compound from the solid state to the liquid state. Therefore, ionic compounds have high melting and boiling points and are solids at r.t.p.
- Clicking on the URL button will link you to <http://www.youtube.com/watch?v=EBfGcTAJF4o&feature=related>, a website with a video simulation on how water molecules “pull” the positive and negative ions in an ionic compound away from each other.
- Sodium chloride does not conduct electricity in the solid state, but it can conduct in the molten and aqueous states.
- A salt in the solid state, as shown earlier, has its ions held rigidly in fixed positions. It does not conduct electricity simply because the ions cannot move around. However, if we take the salt and dissolve it in water, the water molecules will pull the positive and negative ions away from each other. As a result, the ions will be free to move around to conduct electricity. Similarly, in the molten state, the ions will overcome some of the attractive forces between them and be free to move around. The presence of mobile ions enables electricity to be conducted.