Chapter 2: Properties of Materials - Atomic Structure

Chapter 2 Properties and Materials  Learning Objectives 2.1 Understand that the structure of the Periodic Table is related to the atomic structure of the elements and the Periodic Table can be used to predict an element’s structure and properties. 2.2 Understand that th...

Chapter 2 Properties and Materials  Learning Objectives 2.1 Understand that the structure of the Periodic Table is related to the atomic structure of the elements and the Periodic Table can be used to predict an element’s structure and properties. 2.2 Understand that the groups within the Periodic Table have trends in physical and chemical properties, using group 1 as an example. 2.3 Understand that a molecule is formed when two or more atoms join together chemically, through a covalent bond. 2.4 Describe a covalent bond as a bond made when a pair of electrons is shared by two atoms (limited to single bonds). 2.5 Describe an ion as an atom which has gained at least one electron to be negatively charged or lost at least one electron to be positively charged. 2.6 Describe an ionic bond as an attraction between a positively charged ion and a negatively charged ion. 2.7 Know that elements and compounds exist in structures (simple or giant), and this influences their physical properties.  Learning Outcomes 2.1.1 Draw and describe a model of an atom. 2.1.2 Interpret information about the electronic structure of an atom. 2.1.3 Explain how the elements are arranged in the Periodic Table. 2.2.1 Explain how the structure of the Periodic Table is related to the atomic structure of the elements. 2.2.2 Identify the similarities and differences between the atomic structure of elements in the same group in the Periodic Table. 2.2.3 Predict the structure and properties of elements using the Periodic Table. 2.3.1 Explain what a molecules is. 2.4.1 Explain how covalent bond is formed. 2.4.2 Write the formulae of some covalent compounds. 2.4.3 Draw dot and cross diagram of covalent compound. 2.4.4 State that covalent bond is formed where electrons are shared.  Learning Outcomes 2.5.1 Explain the differences between the structure of an ion and an atom. 2.5.2 Draw the electronic structure of some ions (sodium ion, chloride ion). 2.5.3 Explain that positive ion is formed when an atom loses electron(s). 2.5.4 Explain that negative ion is formed when an atom gains electron(s). 2.6.1 Explain how ionic bond is formed. 2.6.2 Write the formulae of some ionic compounds. 2.6.3 Draw the electronic structure of some ionic compounds. 2.7.1 Explain how giant structures are formed. 2.7.2 State two differences between the properties of ionic and covalent compounds. 2.7.3 State the relationship between the structure of ionic and covalent compounds to their properties.  Content 01 02 Atomic Structure Trends in Periodic Table 03 04 05 Simple and Ionic Bond Covalent Giant Molecules Bond  01 Atomic Structure 2.1.1 Draw and describe a model of an atom. 2.1.2 Interpret information about the electronic structure of an atom. 2.1.3 Explain how the elements are arranged in the Periodic Table. STRUCTURE OF ATOM (Nucleus) Subatomic particles 6 protons 6 neutrons 6 electrons Electron shell CALCULATING THE SUBATOMIC PARTICLES: PROTONS, ELECTRONS AND NEUTRONS The masses of subatomic particles are very tiny. Instead of writing their actual masses in kilograms, we often use their relative masses. Protons and electrons have opposite electrical charges. Most of the mass of an atom is in the nucleus (protons and neutrons).  Atomic structure 6 Electrons 3 Electrons Nucleus (3 protons + 4 neutrons) Lithium atom Number of Protons = Number of Electrons  Atomic structure 6 Electrons Proton Electron (+ charge) (- charge) Electrostatic attraction between the opposite charges of protons and electrons hold them together in an atom.  Chemical symbol Atomic Chemical symbol number (number of protons) Name Atomic mass Elements are arranged in the Periodic Table according to the ascending order of atomic number (proton number). CALCULATING PROTONS, NEUTRONS AND ELECTRONS number of protons number of protons = atomic number = 4 = atomic number = 18 number of electrons number of electrons = atomic number = 4 = atomic number = 18 number of neutrons number of neutrons = mass number - atomic = mass number - atomic number number =9-4 =5 = 40 - 18 = 22  2.1.2 Interprete information about the electronic structure of an atom. How many protons, electrons and neutrons are there in the atoms below? Nitrogen Sodium Lithium Proton = 7 Proton = 11 Proton = 3 Neutron = 7 Neutron = 12 Neutron = 4 Electron = 7 Electron = 11 Electron = 3 In an atom, the number of protons = number of electrons.  Periodic Table Song There’s Hydrogen and Helium, Then Lithium, Beryllium, Boron, Carbon everywhere, Nitrogen all through the air, With Oxygen, so you can breath, And Fluorine for your pretty teeth, Neon to light up the signs, Sodium for salty times, Magnesium, Aluminium, Silicon, Phosphorus, then Sulfur, Chlorine and Argon, Potassium and Calcium, so you’ll grow strong. Atomic Element Chemical Atomic Element Chemical number symbol number symbol 1 Hydrogen 11 Sodium 2 Helium 12 Magnesium 3 Lithium 13 Aluminium 4 Beryllium 14 Silicon 5 Boron 15 Phosphorus 6 Carbon 16 Sulfur 7 Nitrogen 17 Chlorine 8 Oxygen 18 Argon 9 Fluorine 19 Potassium 10 Neon 20 Calcium 2.1.1 Draw and describe a model of an atom. The electronic structure of atoms The electrons occupy electron (max 2e) shells around the nucleus of an atom. ​ (max 8e) Shells are also called orbits/energy levels. (max 8e) The outermost electron shell has the highest energy level. (highest energy level) The electronic structure of Chlorine (Cl) atom Number of Electrons = 17 Electronic structure = 2,8,7 (1st shell = 2 e-, 2nd shell = 8 e-, 3rd shell = 7 e-) Dot and Cross diagram The electronic structure of Sodium (Na) atom Number of Electrons = 11 Electronic structure = 2,8,1 (1st shell = 2 e-, 2nd shell = 8 e-, 3rd shell = 1 e-) Dot and Cross diagram  Create a table with all the first 20 elements of the periodic table. Exercise The first 3 elements are done for you as an example Name of Atomic No. of proton No. of Electronic Electronic element Number or electron neutron structure arrangement Hydrogen 1 1 - 1 Helium 2 2 2 2 Lithium 3 3 4 2,1 2.1.3 Explain how the elements are arranged in the Periodic Table. Elements are arranged in ascending/increasing order of the atomic number (proton) in the Periodic Table. 2.1.3 Explain how the elements are arranged in the Periodic Table. Hydrogen, H Helium, He Atomic no = 1 Atomic no = 2 Elements are arranged in ascending/increasing order of the atomic number (proton) in the Periodic Table. 2.1.3 Explain how the elements are arranged in the Periodic Table. G1 G8 G2 G3 G4 G5 G6 G7 G R O U P Elements are arranged in Groups (vertical). Example: Group 1: Hydrogen (H), Lithium (Li), Sodium (Na) and Potassium (K). Group 8: Helium (He), Neon (Ne), and Argon (Ar). 2.1.3 Explain how the elements are arranged in the Periodic Table. G1 G8 G2 G3 G4 G5 G6 G7 P1 G R P2 O U P3 P P4 Period Elements are arranged in Period (horizontal). Example: Period 1: Hydrogen (H) and Helium (He). Period 2: Lithium (Li), Beryllium (Be), Boron (B), Carbon (C), Nitrogen (N), Oxygen (O), Fluorine (F) and Neon (Ne). 2.1.3 Explain how the elements are arranged in the Periodic Table. G1 G8 G2 G3 G4 G5 G6 G7 P1 G R P2 O U P3 P P4 Period Where is the position of the elements below in the Periodic Table? Lithium (Li) Phosphorus (P) Helium (He) Calcium (Ca) Group 1 Group 5 Group 8 Group 2 Period 2 Period 3 Period 1 Period 4 Learning Outcomes: 2.1.1 Draw and describe a model of an atom. Sodium atom (Na) Atomic number = 11 Electronic structure = 2.8.1 2.1.2 Interprete information about the electronic structure of an atom. Sodium atom = 2.8.1 There are 11 protons, 11 electrons, and 12 neutrons in an Sodium (Na) atom. The atomic number = 11 The atomic mass = 23  Learning Outcomes: 2.1.3 Explain how the elements are arranged in the Periodic Table. G1 G8 G2 G3 G4 G5 G6 G7 P1 G R O P2 U P P3 P4 Period Elements are arranged in: Groups (vertical) and Period (horizontal).  02 Trends in Periodic Table 2.2.1 Explain how the structure of the Periodic Table is related to the atomic structure of the elements. 2.2.2 Identify the similarities and differences between the atomic structure of elements in the same group in the Periodic Table. 2.2.3 Predict the structure and properties of elements using the Periodic Table. 2.2.1 Explain how the structure of the Periodic Table is related to the atomic structure of the elements. Elements are arranged in ascending/increasing order of the atomic number (proton) in the Periodic Table. G P1 R O P2 U P P3 P4 Period Elements are arranged in Groups (vertical) and Period (horizontal).  Group 1 Elements → The alkali metals Lithium, Sodium, Potassium, Rubidium, Caesium, Francium  Group 1 Elements → The alkali metals Try to write and draw the electronic structures of Lithium (Li), Sodium (Na) and Potassium (K).  Group 1 Elements → The alkali metals Lithium (Li) Sodium (Na) Potassium (K) 2,1 2,8,1 2,8,8,1 They all have 1 electron at the outermost shell → 1 valence electron. Did you notice the similarities among these elements? 2.2.2 Identify the similarities and differences between the atomic structure of elements in the same group in the Periodic Table. *Take note that the number of electrons occupying the outermost shell is 1. Valence electron is 1. Number of shells filled with electrons increase down the group. Atoms get larger as you go down the group 2.2.1 Explain how the structure of the Periodic Table is related to the atomic structure of the elements. The Group Number → The number of valence electrons in the atom. Example: Group 1 = 1 valence electron Group 4 = 4 valence electrons. 2.2.1 Explain how the structure of the Periodic Table is related to the atomic structure of the elements. The Period Number → The number of shells filled with electrons. Example: Period 1 → only 1 shell filled with electrons. Period 2 → 2 shells filled with electrons. 2.2.1 Explain how the structure of the Periodic Table is related to the atomic structure of the elements. The Group Number → The number of valence electrons in the atom. The Period Number → The number of shells filled with electrons. Try this! Element X has an atomic number of 17. What is the position of the element in the Periodic Table? What is X? Electronic structure = 2.8.7 7 valence electrons → Group 7 Element X is Chlorine (Cl) 3 shells filled with electrons → Period 3 2.2.1 Explain how the structure of the Periodic Table is related to the atomic structure of the elements. Element Y has an atomic number of 20. What is the position of the element in the Periodic Table? What is Y? Electronic structure = 2.8.8.2 2 valence electrons → Group 2 Element Y is Calcium (Ca) 4 shells filled with electrons → Period 4 The Group Number → The number of valence electrons in the atom. The Period Number → The number of shells filled with electrons. The difference between Electronic Structure and Atomic Structure Atomic Structure Electronic Structure 11 protons 12 neutrons 6 electrons 11 electrons Sodium atom - Show electrons in - Show protons, neutrons and shells. electrons in shells. Electronic Structure/Configuration/Arrangement Writing Drawing 2,8,1 Group 1 Elements – The alkali metals What is the colour of each element? Are they shiny? Li - 3 The elements in the first group are called alkali metals. ○ They conduct electricity K - 19 ○ They are shiny when freshly cut Li, Na & K can be cut easily with a knife Na - 11 Group 1 Element – Physical Properties and pattern. Alkali metal Atomic Melting point (˚C) Boiling point number (˚C) Lithium, Li 3 180 1342 Sodium, Na 11 98 883 Potassium, K 19 63 759 What is the trend in the melting and boiling point? Going down the group, from Li to K, melting point and boiling point decrease.  Group 1 Element – Chemical Properties Reaction with water All the alkali metals are very reactive, that is they undergo chemical reactions readily. They fizz and zoom around on the surface of water. They get more vigorous going down the group. They react readily with water to form metal hydroxide (alkaline) and hydrogen gas. Universal indicator turns blue or purple. Sodium + Water Sodium hydroxide + Hydrogen  Potassium produces sparks/flames when reacting with water The reactivity of alkali metals with water increases down the group. Group 1 Element – Chemical properties Reaction with oxygen When the alkali metals are cut, they initially appear shiny grey but quickly become dull and white as they react with oxygen in the air. They react readily with oxygen in the air to form metal oxide. Universal indicator turns blue or purple. Sodium + oxygen sodium oxide  Group 1 Element – Chemical properties Reaction with oxygen Lithium burns with a red flame in excess oxygen. Sodium burns with a bright yellow flame in excess oxygen. Potassium burns with a very bright lilac flame in excess oxygen. All three reactions produce a white solid (metal oxide). They get more vigorous going down the group. 2.2.3 Predict the structure and properties of elements using the Periodic Table. Trends of elements in the same group (Group 1) Chemical Properties 1) Same number of valence electrons. (All Group 1 elements have 1 valence electron.) 1) Similar chemical properties (reactivity) a) Group 1 elements are reactive with water. Metal hydroxide and hydrogen gas are formed. The reactivity increases going down the group. b) Group 1 elements are reactive with oxygen. Metal oxide is formed. The reactivity increases going down the group. 2.2.3 Predict the structure and properties of elements using the Periodic Table. Trends of elements in the same group (Group 1) Physical Properties 1) The size of atoms increases going down the group. - The number of shells filled with electrons increases. 1) Low density -> lithium, sodium and potassium float on water. 2) Good conductors of electricity and heat. 3) Shiny when cut, but quickly tarnished upon exposure to air. 4) Soft and can be cut with a knife. 5) Melting point and boiling point decrease going down the group. 2.2.3 Predict the structure and properties of elements using the Periodic Table. Group 1 alkali metals Lithium, Sodium and Potassium react with water and oxygen. How is the reaction of Rubidium with water and oxygen? Rubidium reacts with water and oxygen. The reactivity of rubidium with water and oxygen is higher than potassium// Rubidium reacts more vigorously with water and oxygen than potassium. 2.2.3 Predict the structure and properties of elements using the Periodic Table. Group 1 alkali metals Lithium, Sodium and Potassium are good conductor of heat and electricity. Is Rubidium conductor of heat and electricity? Rubidium is a good conductor of heat and electricity.  Group Work: In your group, gather information about Group 7 and Group 8 of Periodic Table. You need to include: Examples of elements (Group 7 & 8) Reactivity (Group 7 & 8) Melting and boiling point (Group 7 & 8) Comparison table for Group 1, Group 7 and Group 8 Learning outcomes: 2.2.1 Explain how the structure of the Periodic Table is related to the atomic structure of the elements. The Group Number Valence electrons → The number of _____________________ in the atom. The Period Number → The number of ______________________________________. Shells filled with electrons Learning outcomes: 2.2.2 Identify the similarities and differences between the atomic structure of elements in the same group in the Periodic Table. Similarities: Elements of the same group will have same number of valence electrons. Differences: Elements of the same group different number of shells filled with electrons. The shells filled with electrons increases down the group. 2.2.3 Predict the structure and properties of elements using the Periodic Table.  03 Ionic Bond 2.5.1 Explain the differences between the structure of an ion and an atom. 2.5.2 Draw the electronic structure of some ions (sodium ion, chloride ion). 2.5.3 Explain that positive ion is formed when an atom loses electron(s). 2.5.4 Explain that negative ion is formed when an atom gains electron(s). Big numbers catch your audience’s attention 2.6.1 Explain how ionic bond is formed. 2.6.2 Write the formulae of some ionic compounds. 2.6.3 Draw the electronic structure of some ionic compounds.  Stable Electronic Structure Helium (2) Neon (2,8) Argon (2,8,8) Why are these 3 atoms said to be stable? Their outermost electron shell is completely full of electrons. Group 8 elements are stable -> inert (not reactive) -> also known as Noble gas  Unstable Electronic Structure Sodium (2,8,1) Carbon (2,4) Chlorine (2,8,7) Why are these 3 atoms said to be unstable (reactive)? Their outermost electron shell is NOT full of electrons. They can react with other atom(s) to form molecule or compound.  Unstable Atoms + Sodium (2,8,1) Chlorine (2,8,7) Sodium chloride compound How do the unstable atoms become stable? They can react with other atom(s) to form molecule or compound. → Form Chemical Bond  Formation of Chemical Bonds Two ways to form chemical bonds: Atoms can lose or gain electrons. Atoms can share electrons. → Ionic Bond → Covalent Bond Sodium chloride compound Water molecule 2.5.1 Explain the differences between the structure of an ion and an atom. 2.5.3 Explain that positive ion is formed when an atom loses electron(s). Positive Ion The outermost + electron Loses 1 electron Atomic number = 11 [2,8]+ [2,8,1] The outermost electron shell is The outermost electron shell is full NOT full → Stable (not reactive) → Unstable (reactive) → becomes an ion (sodium ion, Na+), → Lose 1 electron (outermost) with positive charge. 2.5.1 Explain the differences between the structure of an ion and an atom. 2.5.3 Explain that positive ion is formed when an atom loses electron(s). Positive Ion The outermost The outermost electron electron Loses 2 electrons Atomic number = 12 Magnesium ion, Mg2+ Magnesium atom, Mg [2,8]2+ [2,8,2] The outermost electron shell is The outermost electron shell is full NOT full → Stable (not reactive) → Unstable (reactive) → becomes an ion (magnesium ion, Mg2+), → Lose 2 electrons (outermost) with positive charge. 2.5.3 Explain that positive ion is formed when an atom loses electron(s). 2.5.2 Draw the electronic structure of some ions (sodium ion, chloride ion). Positive Ion Practise: Atomic Atomic number Question 1 number Question 2 = 19 Aluminium atom, Al = 13 Potassium atom, K + 3+ -1e- -3e- Aluminium Aluminium Potassium Potassium atom, Al ion, Al3+ atom, K ion, K+ [2,8,3] [2,8]3+ [2,8,8,1] [2,8,8]+ 2.5.1 Explain the differences between the structure of an ion and an atom. 2.5.4 Explain that negative ion is formed when an atom gains electron(s). Negative Ion Gain 1 e- Atomic number = 17 Chlorine atom, Cl Chloride ion, Cl- [2,8,7] [2,8,8]- The outermost electron shell is The outermost electron shell is full NOT full → Stable (not reactive) → Unstable (reactive) → becomes an ion (chloride ion, Cl-), → Gain 1 electron (outermost) with negative charge. 2.5.1 Explain the differences between the structure of an ion and an atom. 2.5.4 Explain that negative ion is formed when an atom gains electron(s). Negative Ion Gain 2 e- Atomic number =8 Oxygen atom, O Oxide ion, O2- [2,6] [2,8]2- The outermost electron shell is The outermost electron shell is full NOT full → Stable (not reactive) → Unstable (reactive) → becomes an ion (oxide ion, O2-), → Gain 2 electrons (outermost) with negative charge. 2.5.4 Explain that negative ion is formed when an atom gains electron(s). 2.5.2 Draw the electronic structure of some ions (sodium ion, chloride ion). Negative Ion Practise: Atomic Atomic number Question 1 number Question 2 =7 Sulfur atom, S = 16 Nitrogen atom, N 3- +3e - +2e - Sulfur Sulfide Nitrogen Nitride atom, S ion, S2- atom, N ion, N3- [2,8,6] [2,8,8]2- [2,5] [2,8]3- 2.5.3 Explain that positive ion is formed when an atom loses electron(s). 2.5.4 Explain that negative ion is formed when an atom gains electron(s). What kind of elements/atoms tend to become positive or negative ion? Metals lose electrons to Non-metals gain electrons become positive ions. to become negative ions. Group 1 → 1 valence e- Group 5 → 5 valence e- → lose 1 e- → X+ → gain 3 e- → Y3- Group 2 → 2 valence e- Group 6 → 6 valence e- → lose 2 e- → X2+ → gain 2 e- → Y2- Group 3 → 3 valence e- Group 7 → 7 valence e- → lose 3 e- → X3+ → gain 1 e- → Y- 2.5.3 Explain that positive ion is formed when an atom loses electron(s). 2.5.4 Explain that negative ion is formed when an atom gains electron(s). Element Positive ion Name of ion Element Negative ion Name of ion H H+ hydrogen F F- fluoride Li Li+ lithium Cl Cl- chloride Na Na+ sodium O O2- oxide K K+ potassium S S2- sulfide Be Be2+ beryllium N N3- nitride Mg Mg2+ magnesium P P3- phosphide Ca Ca2+ calcium Al Al3+ aluminium 2.6.2 Write the formulae of some ionic compounds. Writing chemical formula of ionic compounds: Positive ion Negative ion Example: Px+ Qy- Na+ S2- Charge x y 1 2 Simplest x y 1 2 ratio Chemical formula PyQx Na2S 2.6.2 Write the formulae of some ionic compounds. Writing chemical formula of ionic compounds: Positive ion Negative ion Example: Px+ Qy- Ca2+ S2- Charge x y 2 2 Simplest x y 1 1 ratio Chemical formula PyQx CaS 2.6.2 Write the formulae of some ionic compounds. Practice: 1) Sodium chloride (Na+, Cl-) 3) Aluminium sulfide (Al3+, S2-) NaCl Al2S3 (1 sodium ion, Na+ and 1 chloride ion, Cl-) (2 Aluminium ions, Al3+ and 3 sulfide ions, S2-) 2) Magnesium chloride (Mg2+, Cl-) 4) Calcium oxide (Ca2+, O2-) MgCl2 CaO (1 Magnesium ion, Mg2+ and 2 chloride ions, Cl-) (1 calcium ion, Ca2+ and 1 Oxide ion, O2-) 2.6.1 Explain how ionic bond is formed. 2.6.3 Draw the electronic structure of some ionic compounds. Ionic Bond Chemical bond formed by the attraction between the positively charged ion and negatively charged ion. [Electrostatic force] Positive ion (metal) Negative ion (non-metal) Na+ Cl- Sodium chloride, NaCl NaCl is an ionic compound. 2.6.1 Explain how ionic bond is formed. 2.6.3 Draw the electronic structure of some ionic compounds. Ionic Bond Formation of Sodium chloride compound: + Chlorine atom [2,8,1] [2,8,7] Outermost electron shell is full (stable). → lose 1 e- → gain 1 e- Ionic compound of sodium chloride, NaCl is formed. 2.6.1 Explain how ionic bond is formed. 2.6.3 Draw the electronic structure of some ionic compounds. Ionic Bond Formation of Magnesium chloride compound: + Magnesium ion, Mg2+ Magnesium atom Chlorine atom [2,8]2+ [2,8,2] [2,8,7] Outermost electron shell is full (stable). → lose 2 e- → gain 1e- Ionic compound of magnesium chloride, → 2 chlorine MgCl2 is formed. atoms are needed 2.6.1 Explain how ionic bond is formed. 2.6.3 Draw the electronic structure of some ionic compounds. Group work: Show the formation of ionic compounds below by showing their electronic structures: 1) Potassium chloride 2) Magnesium fluoride 3) Calcium oxide 4) Lithium sulfide 2) Magnesium fluoride 4) Lithium sulfide  Extension They are from Group 1. Which element is the most reactive out of all? Explain. Potassium is the most reactive one. → The 1 valence electron is the furthest from the nucleus. → the attraction between the valence electron (- charge) and the nucleus (+positive charge) is the weakest. → Easiest to lose the 1 valence electron. Conclusion: The reactivity of Group 1 elements increases going down the group.  Extension They are from Group 7. Which element is more reactive? Explain. Fluorine is the more reactive. → The 7 valence electrons is the closer from the nucleus. → the attraction between the valence electron (- charge) and the nucleus (+positive charge) is the stronger. → Easiest to attract (gain) the 1 more electron. Conclusion: The reactivity of Group 7 elements decreases going down the group.  04 Covalent Bond 2.3.1 Explain what a molecule is. 2.4.1 Explain how covalent bond is formed. 2.4.2 Write the formulae of some covalent compounds. Big numbers catch your audience’s attention 2.4.3 Draw dot and cross diagram of covalent compound. 2.4.4 State that covalent bond is formed where electrons are shared. 2.3.1 Explain what a molecule is. All substances are made of atoms All substances are made of very tiny particles called atoms. Many substances are made up of different types of atoms. Carbon dioxide Oxygen Copper atoms 2.3.1 Explain what a molecule is. Atom Molecule The smallest unit of a matter 2 or more atoms bonded together Oxygen gas molecule (oxygen atoms) Oxygen Carbon Hydrogen atom atom atom Water molecule (hydrogen and oxygen atoms) Carbon dioxide gas molecule (carbon and oxygen atoms) 2.4.1 Explain how covalent bond is formed. 2.4.4 State that covalent bond is formed where electrons are shared. Covalent Bond Chemical bond formed by the sharing of electrons between a non-metal and another non-metal. Non-metal A Non-metal B H Cl Hydrogen chloride, HCl HCl is a covalent compound (molecule). 2.4.1 Explain how covalent bond is formed. 2.4.2 Write the formulae of some covalent compounds. 2.4.3 Draw dot and cross diagram of covalent compound. Covalent Bond + Hydrogen atom, H Chlorine atom, Cl Hydrogen atom, H Chlorine atom, Cl [2,8,7] [2,8,8] Hydrogen atom and Chlorine atom share 1 pair of → needs 1 more e- → needs 1 more e- e-. Outermost electron shell is full (stable). Covalent compound of hydrogen chloride, HCl is formed. 2.4.1 Explain how covalent bond is formed. 2.4.2 Write the formulae of some covalent compounds. 2.4.3 Draw dot and cross diagram of covalent compound. Covalent Bond Formation of hydrogen molecule, H2 + Hydrogen atom, H Hydrogen atom, H Hydrogen atom, H Hydrogen atom, H → needs 1 more e- → needs 1 more e- Two Hydrogen atoms share 1 pair of e-. Outermost electron shell is full (stable). Covalent molecule (hydrogen molecule, H2) is formed.  Covalent Bond Formation of ammonia molecule, NH3 + Hydrogen atom, H Nitrogen atom, N Hydrogen atom, H Nitrogen atom, N [2,8] [2,5] 1 Nitrogen atoms share 3 pairs of e- with 3 Hydrogen atoms. Outermost electron shell is → each H atom → needs 3 more e- full (stable). Covalent compound (ammonia needs 1 more e- molecule, NH3) is formed. More examples of covalent molecules: 2.4.3 Draw dot and cross diagram of covalent compound. Individual work: Show the formation of covalent bonds below by showing their electronic structures: 1) Chlorine molecule, Cl2 2) Water molecule, H2O 3) Methane molecule, CH4  05 Simple and Giant Molecules 2.7.1 Explain how giant structures are formed. 2.7.2Big State two differences numbers between catch your the attention audience’s properties of ionic and covalent compounds. 2.7.3 State the relationship between the structure of ionic and covalent compounds to their properties. 2.7.1 Explain how giant structures are formed. Giant structures in Ionic compounds Sodium chloride, NaCl is an ionic compound. The oppositely charged ions (Na+ and Cl-) are strongly attracted to each other Lattice structure of NaCl by electrostatic force. Forms giant structure known as lattice. Ions are arranged in a regular pattern. → Sodium chloride forms crystals with regular shape. Sodium chloride crystal Simple covalent molecules Many covalent molecules are simple molecules. Example: oxygen, carbon dioxide, methane. The force holding the molecules together are very strong. But the force between the molecules are very weak → intermolecular force. Weak intermolecular force → easy to break → low melting/boiling point → appear as liquid or gas 2.7.1 Explain how giant structures are formed. Giant covalent molecules The carbon atoms in diamond form a giant structure. Each carbon atom forms four strong covalent bonds. Diamond - hardest material on Earth. Strong, rigid, 3-dimensional structure of lattice. Used for jewellery, cutting and drilling tools 2.7.1 Explain how giant structures are formed. Giant covalent molecules Large structure → Macromolecules A diamond lattice structure; each carbon atom forms four strong covalent bonds. 2.7.1 Explain how giant structures are formed. Giant covalent molecules The carbon atoms in graphite form a giant structure. Graphite is used for the “lead” in pencils, and for lubricating moving parts in machines. Carbon atoms each make bonds with three other atoms. They form layers, which can easily slide over one another → it makes the surface very soft and easily comes away. 2.7.1 Explain how giant structures are formed. Giant covalent molecules Weak bond between layers Strong covalent bond → easily slide over one another within layer 2.7.2 State two differences between the properties of ionic and covalent compounds. 2.7.3 State the relationship between the structure of ionic and covalent compounds to their properties. Melting point and Boiling point Ionic compound - High melting point and boiling point. - Strong electrostatic force holding the ions together. - High amount energy is needed to overcome these forces in order to melt or boil them. 2.7.2 State two differences between the properties of ionic and covalent compounds. 2.7.3 State the relationship between the structure of ionic and covalent compounds to their properties. Melting point and Boiling point Covalent molecule/ compound - Simple covalent molecules have low melting point and boiling point. - The forces between the molecules (intermolecular force) are weak. - Small amount energy is needed to Giant covalent molecules have overcome these forces in order to melt high melting point and boiling or boil them. point.  1) Is magnesium chloride an ionic compound or a simple covalent molecule? Explain. Ionic compound. It has high melting point and boiling point. It is also formed by a metal and a non-metal. 2) Is ammonia an ionic compound or a simple covalent molecule? Explain. Covalent compound. It has low melting point and boiling point. It is formed by two non-metals. 3) Is ammonia a solid, liquid or gas at room temperature? Gas. 4) Why do magnesium chloride and calcium oxide have high melting points? They are ionic compounds. The electrostatic forces between the ions in these compounds are very strong so, high amount of energy is need to break these forces to melt them. 5) Why do methane and chlorine have low melting points? They are simple covalent molecules. The forces between the molecules (intermolecular force) are weak so, less energy is need to break these forces to melt them. Conducting electricity Ionic compound - Can conduct electricity because ions have electrical charges. - But the ions must be free to move to carry the charges. - In solid form, ionic compounds cannot conduct electricity (ions do not move). - When dissolved in water or melt to form liquid, ionic compounds can conduct electricity.  Conducting electricity Covalent compound - Do not conduct electricity. - Do not have free moving ions. - Some covalent compounds are weak conductors of electricity. Example: Water, acid.  Comparison of Ionic substances and Covalent substances Ionic substances Covalent substances Giant lattices made from ions. Simple molecules. Some are giant molecules. Properties High melting point and boiling Simple molecules have low point melting point and boiling point. Can conduct electricity when Do not conduct electricity. dissolved in water or melted.  Learning Outcomes 2.1.1 Draw and describe a model of an atom. 2.1.2 Interpret information about the electronic structure of an atom. 2.1.3 Explain how the elements are arranged in the Periodic Table. 2.2.1 Explain how the structure of the Periodic Table is related to the atomic structure of the elements. 2.2.2 Identify the similarities and differences between the atomic structure of elements in the same group in the Periodic Table. 2.2.3 Predict the structure and properties of elements using the Periodic Table. 2.3.1 Explain what a molecules is. 2.4.1 Explain how covalent bond is formed. 2.4.2 Write the formulae of some covalent compounds. 2.4.3 Draw dot and cross diagram of covalent compound. 2.4.4 State that covalent bond is formed where electrons are shared.  Learning Outcomes 2.5.1 Explain the differences between the structure of an ion and an atom. 2.5.2 Draw the electronic structure of some ions (sodium ion, chloride ion). 2.5.3 Explain that positive ion is formed when an atom loses electron(s). 2.5.4 Explain that negative ion is formed when an atom gains electron(s). 2.6.1 Explain how ionic bond is formed. 2.6.2 Write the formulae of some ionic compounds. 2.6.3 Draw the electronic structure of some ionic compounds. 2.7.1 Explain how giant structures are formed. 2.7.2 State two differences between the properties of ionic and covalent compounds. 2.7.3 State the relationship between the structure of ionic and covalent compounds to their properties. End of Chapter 2