Cambridge Lower Secondary Year 9 Science Chapter 2 Properties of Materials PDF

Cambridge Lower Secondary Year 9 Science (0893) Chapter 2 Properties of Materials MR. VENTUS TAN WAI SHAN MHSc Biomedical Science (UKM), BSc Microbiology (USM) Email: [email protected] Name the Scientists Who Discovered 1. Atom: John Dalton 2. Proton: 3. Neutron: 4. Electron: Atomic Structure All substances are made of matter called atoms Each atom is made of subatomic particles called protons, neutrons, and electrons The protons and neutrons are located at the Nucleon centre of the atom, which is called the nucleus The electrons move very fast around the nucleus in orbital paths called shells The mass of the electron is negligible, hence the mass of an atom is contained within the nucleus where the protons and neutrons are located Types of Subatomic Particles Subatomic particles Charges Mass Location Proton +1 1 In nucleus NUCLEON Neutron 0 1 In nucleus Electron -1 1/1840 Shell Number of nucleon = Number of proton + Number of neutron Atomic Number The atomic number (proton number) is the number of protons in the nucleus of an atom The symbol for atomic number is Z. It is also the number of electrons present in a neutral atom and determines the position of the element on the Periodic Table Charges 2+ Atomic Mass The nucleon number (or mass number) is the total number of protons and neutrons in the nucleus of an atom The symbol for nucleon number is A The nucleon number minus the proton number gives us the number of neutrons of an atom Atomic Notation The atomic number and mass number of an element can be shown using atomic notation The Periodic Table shows the elements together with their atomic (proton) number at the top and relative atomic mass at the bottom. NOTATION FORM PERIODIC TABLE FORM Finding Protons, Neutrons and Electrons Finding the protons The atomic number of an atom and ion determines which element it is The number of protons equals the atomic (proton) number Number of protons = mass number – number of neutrons Finding the electrons An atom is neutral and therefore has the same number of protons and electrons Finding the neutrons The mass and atomic numbers can be used to ﬁnd the number of neutrons in ions and atoms: Number of neutrons = mass number – number of protons Atomic Number and Mass An atom has a nucleus. Nucleus are made up of two subatomic atoms: - The positively charged protons - The neutrons which do not have any charge Each element has a diﬀerent number of protons in its atoms. Helium atoms have two protons. Lithium atoms have three protons and gold have 79 protons. The number of proton in an element represents its proton number. Poroton number can also be called as atomic number. Periodic Table Notation Proton number: 6 Nucleon number: 12 6 12 Neutron number: 6 C C Number of electron: 6 Electronic conﬁguration: 2, 4 12 6 Group: 4 Period: 2 The atomic number is the number of protons in the nucleus of a atom Shell The number of proton plus neutron 0 1 +1 1 29 34 26 82 30 36 20 16 20 2 5 18 Try! Periodic Notation Proton Nucleon Neutron Electron Electronic Table [proton + neutron) Configuration 7 14 N 14 7 N 7 14 7 7 14 Si 28 si 14 28 14 14 40 Ca 20 20 40 20 20 13 14 13 HI 39 27 20 19 19 The amount positive charge in an atom Ba.ro 8 17 47 vanadium Argonyar Iron Fe 4 Structure of Boron 5 X B 11 X Proton= 5 Neutron = 6 X Nucleon = 11 X X 6 4 6 Aluminium Electronic Conﬁguration Electrons Shell and The Periodic Table The electronic conﬁguration will show the number of occupied shells of electrons the atom has, showing the period in which that element is in The last notation shows the number of outer electrons the atom has, showing the group that element is in (for elements in Groups I to VII) Elements in the same group have the same number of outer shell electrons which makes the same group elements to share similar chemical properties. Exercise Particles Proton Electronic conﬁguration Na 11 2.8.1 F 9 2.7 Cl 17 2.8.7 Al 13 2.8.3 O 8 2.6 N 7 2.5 Ca 20 2 8 8 2 I 1 I Practice Element Proton Electron Electron arrangement Hydrogen 1 1 I 2 2 2 Helium Lithium 3 3 2 1 6 6 214 Carbon Nitrogen 7 2 2 5 9 9 Fluorine 2 7 10 10 2 Neon 8 11 11 2 81 Sodium Aluminium 13 13 2 8 3 Silicon 14 14 2 8 4 Phosphorus 15 15 2 8 5 16 16 8 6 Sulfur Argon 2 8 8 Outer Shell Electrons Atoms of element in Group 1 which is in the left column have only one electron in their outer shell. Atoms of all element in the same group in the periodic table have the same number of outer shell electrons which is also called valence electron. Electron conﬁguration gives an element its structure and as a result its chemical properties. 0 1 2 at 4 3 2 1 the arrangement of electron show 7 8 9 2.1 2.8.1 2.8.8.2 Forming Ions Ion is a particle with a positive or negative charge. Ions are formed when the atom gains or loses electrons. - If an atom gains one or more electrons, it becomes a negatively charged ion (anion). - If an atom loses one or more electrons, it becomes a positively charged ion (cation). A sodium ion forms when a sodium atom gives one electron to a non-metal atom. - The sodium atom has 11 positive protons and 11 negative electrons. It has no net charge. - The chlorine atom has 17 positive protons and 17 negative electrons. It has no net charge. In the reaction, one electron moves from sodium atom to the chlorine atom to achieve stable electron arrangement. Sodium has the electron arrangement of 2,8,1, by which it needs to donate 1 electron to achieve stability of 2,8. - The sodium ion has 11 positive protons and 10 negative electrons. Its charge +1. Chlorine has the electron arrangement of 2,8,7 by which it needs to gain 1 electron to achieve stability of 2,8,8. - The chloride ion has 17 positive protons and 18 negative electrons. Its charge -1. Proton Element Symbol Electronic Ions Period Group Number structure formed 1 Hydrogen H 2 3 Helium Lithium He Li 7 2 I 1 2 1 4 Beryllium Be 2 2 2 2 2 5 Boron B 2 2 3 3 3 6 Carbon C 2 4 41 4 2 4 7 Nitrogen N 2 5 3 2 5 8 Oxygen O 2 6 2 2 6 9 Fluorine F 2 7 I 2 7 10 Neon Ne 2 8 0 2 8 Proton Element Symbol Electronic Ions Period Group Number structure formed 11 Sodium Na 2 8 1 1 3 1 12 Magnesium Mg 2 8.2 2 3 2 13 Aluminium Al 2 8.3 3 3 14 Silicon Si 2 84 141 4 3 4 15 Phosphorous P 2 8.5 3 3 5 16 Sulphur S 2 8.6 2 3 6 17 Chlorine Cl 3 18 Argon Ar 2187 1 7 2 8 8 3 8 19 Potassium K 28.8.1 11 4 1 2 8 8.2 20 Calcium Ca 2 4 2 Electronic Conﬁguration of Ions The electronic conﬁgurations of ions are usually complete and stable. Negative ion will gain electrons while positive ion will donate electrons to achieve complete electron arrangement as in: 2 2, 8 2, 8, 8 Exercise Particles Electronic conﬁguration Na 2 8.1 Na+ 2 8 F 2.7 Cl- 2.8.8 Al3+ 2 8 2 8 O2- N3- 2 8 Ca2+ 2 8.8 XX xx the atom losses 3 electrons to become stable 11 12 sodium losses one electron the valence shell two atom combine together A pure substance that cannot be separated either by chemically or physica 10 22 Aand B A B 11 11 14 16 17 17 Determine the Group and Period for the Following Elements. 1. Silicon: 2, 8, 4 Group: 4 (outermost electron) Period: 3 (number of shells ﬁlled with electron) 2. Potassium: 2, 8, 8, 1 Group: Period: I 4 3. Nitrogen: 2, 5 Group: Period: 5 2 elements with charge Ions donate or receive electron to achieve stable octet electro arrangement Kt Mg Br 1 negative because it needs one more 9 electron +1 +2 +3 土4 -3 -2 -1 IIII III II I Group I Elements Also known as alkali metals Consists of six elements: lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs) and francium (Fr). Have one outer shell electron Soft metal can be cut with knife Good conductor of heat and electricity Quite low density compared with other metals Reactivity increases when going down the group. Physical Properties of Group I Elements It Soft and easy to cut, getting even softer and denser as going down the group Have shiny silvery surfaces when freshly cut Conduct heat and electricity They all have low melting points and low densities compared to other metals, and the melting point decreases as you move down the Group. Trend in Group I Density MP/BP Reactivity Lithium Lowest Highest Lowest Sodium Potassium Rubidium Caesium Francium Highest Lowest Highest Trend going down the Increasing Decreasing Increasing group (same as group VII) (opposite from (opposite from group group VII) VII) Trends in Group I Elements When going down group I, The melting point and boiling point Decrease ______________. The density ______________. Increase The atomic size ______________. Increase The atomic mass ______________. Increase The reactivity______________. Decrease Li K k Chemical Properties of Group I Elements Form cation with +1 oxidation state. All group 1 salts are soluble in water. Very reactive and must be kept in paraffin oil. Avoid handling with bare hands and must wear gloves and safety goggle. React vigorously in water to form metal hydroxide and hydrogen. 2Li + 2H2O → 2LiOH + H2 Oxidize easily in the air to form metal oxide. 4Li + O2 → 2Li2O Try! 1. Potassium + water → Potassium hydroxide Hydrogen 2. Sodium + water → sodium h How to test for hydrogen gas? 3. Potassium + oxygen → Test: lighted wooden Result: split pop sound 4. Sodium + oxygen → Red flame General observations when group Yellow flame I elements is added into water: 1. Group I metals will float 2. Group I metals will dissolve and becomes smaller 3. Fizz/ gas bubbles 4. Burns with flame 5. Forms alkaline solution Lilac flame Reactivity of Group I Elements As the reactivity of alkali metals increases down the group, rubidium, caesium and francium will react more vigorously with air and water than lithium, sodium and potassium Lithium will be the least reactive metal in the group at the top, and francium will be the most reactive at the bottom. Group VII Elements Also known as halogen. They are poisonous non-metals which include fluorine, chlorine, bromine, iodine and astatine Halogens are diatomic, meaning they form molecules of two atoms All halogens have seven electrons in their outer shell They form halide ions by gaining one more electron to complete their outer shells Reactivity decreases when going down the group Trend in Group VII Element Melting and boiling point The melting and boiling point of the halogens increases as going down the group Fluorine is at the top of Group VII so will have the lowest melting and boiling point which exists as gas. Astatine is at the bottom of Group VII so will have the highest melting and boiling point which exists as solid Physical states The halogens become denser as going down the group Fluorine is at the top of Group VII so will be a gas Astatine is at the bottom of Group VII so will be a solid Colour The colour of the halogens becomes darker as going down the group Fluorine is at the top of Group VII so the colour will be lighter, so fluorine is pale yellow Astatine is at the bottom of Group VII so the colour will be darker, so astatine is black Summary Halogens Formula Physical State at Colour Room Temperature Fluorine F2 Gas Pale yellow Chlorine Cl2 Gas Greenish yellow Bromine Br2 Liquid Reddish brown Iodine I2 Solid Purple Black Astatine At2 Solid Black Molecule that consist two atoms All have 7 valence electron It has more shell as It goes down teh groupings gas liquid solid black purple Density of Halogens Chlorine is a pale yellow-green gas, bromine is a red-brown liquid and iodine is a grey-black solid This demonstrates that the density of the halogens increases as going down the group Trend in Group VII Density MP/BP Reactivity Colour Fluorine Lowest Lowest Highest Most pale Chlorine Bromine Iodine Astatine Highest Highest Lowest Darkest Trend going Increasing Increasing Decreasing Increasing down the group Trends in Group VII Elements When going down group VII, The melting point and boiling point ______________. The density ______________. The atomic size ______________. The atomic mass ______________. darker The colour ______________. The reactivity______________. Reactivity of Halogens Reactivity of Group VII increases as going up the group (this is the opposite trend to that of Group I) Each outer shell contains seven electrons and when the halogen reacts, it will need to gain one outer electron to get a full outer shell of electrons As going down Group VII, the number of shells of electrons increases This means that the outer electrons are further to the nucleus so there are weaker electrostatic forces of attraction, which help to attract the extra electron needed Usage of Group VII Elements Fluorine Active ingredient in toothpaste. Chlorine Disinfectants for drinking water and swimming pool. Bromine Production of insecticides and dyes. Iodine Used as antiseptics to clear wound. Group VIII Elements Also known as the noble gases They are non-metals and have very low melting & boiling points They are all monoatomic, colourless gases The Group 8 elements all have full outer shells This electronic configuration is extremely stable so these elements are unreactive and are inert Trends in Group VIII Elements When going down group VIII, The melting point and boiling point ______________. The atomic size ______________. The atomic mass ______________. Functions of Group VIII Elements Helium Used to fill weather balloons and airships Has low density and chemically unreactive. Neon Used in advertising light and television tubes Argon Fill electric lamp Radon Used in the treatment of cancer The Ionic Bonds Ionic compounds are formed when metal atoms react with non-metal atoms Metal atoms lose their outer electrons which the non-metal atoms gain to form positive and negative ions The positive and negative ions are held together by strong electrostatic forces of attraction between opposite charges This force of attraction is known as an ionic bond and they hold ionic compounds together Dot and cross diagrams are diagrams that show the arrangement of the outer-shell electrons in an ionic or covalent compound or element ○ The electrons are shown as dots and crosses In a dot and cross diagram: ○ Only the outer electrons are shown ○ The charge of the ion is spread evenly which is shown by using brackets ○ The charge on each ion is written at the top right-hand corner Electron arrangement of atoms Na: 2,8,1 Cl: 2,8,7 Electron arrangement of ions Na+: 2,8 Cl-: 2,8,8 Exercise 3 9 Try to draw the cross and dot diagram for lithium fluoride. I X Exercise Try to draw the cross and dot diagram for calcium sulfide. Giant Lattice Structure Ionic compounds have a giant lattice structure Lattice structure refers to the arrangement of the atoms of a substance in 3D shape In lattice structures, the atoms are arranged in an ordered and repeating fashion The lattices formed by ionic compounds consist of a regular arrangement of alternating positive and negative ions Properties of Ionic Compound Ionic compounds are usually solid at room temperature They have high melting and boiling points Ionic compounds are good conductors of electricity in the molten state or in solution due to the free moving ions. They are poor conductors in the solid state Melting and Boiling Point of Ionic Compounds Ionic substances have high melting and boiling points due to the presence of strong electrostatic forces acting between the oppositely charged ions These forces act in all directions and a lot of energy is required to overcome them The greater the charge on the ions, the stronger the electrostatic forces and the higher the melting point will be ○ For example, magnesium oxide consists of Mg2+ and O2- so will have a higher melting point than sodium chloride which contains the ions, Na+ and Cl- Model of Ionic Bonding Model is an idea that explains observation and helps in making predictions. Giant ionic structure is a model so it explains the physical properties of ionic compounds. - Ionic compounds have high melting point this is because the interaction between the oppositely charged ions are strong. - Ionic compound are brittle. If you drop a crystal onto an ionic compound, it will break between the rows of ions and another. The broken pieces have straight edges. Model has strength and limitations. It main strength of having models is that it can show the position of the oppositely charged ions. Limitations is that it is does not show that the ion vibrate on the spot. It also does not show how the electrostatic attractions act in all direction. This also does not explain why ionic compounds are brittle. Covalent Compounds Covalent compounds are formed when pairs of electrons are shared between atoms Only non-metal elements participate in covalent bonding As in ionic bonding, each atom gains a full outer shell of electrons, giving them a noble gas electronic conﬁguration When two or more atoms are covalently bonded together, we describe them as ‘molecules’ Dot-and-cross diagrams can be used to show the electronic conﬁgurations in simple molecules Electrons from one atom are represented by a dot, and the electrons of the other atom are represented by a cross The electron shells of each atom in the molecule overlap and the shared electrons are shown in the area of overlap The dot-and-cross diagram of the molecule shows clearly which atom each electron originated from ionic coyalist How to draw covalent bond? 1. Draw the atomic electron structure. 2. Centralise the atom which needs more electron to achieve stable electron arrangements. 3. Prioritise the needs of side atoms. 4. Rule: How many you need, how many you will give and how many you will get. Example 1: Oxygen, O2 Electron arrangement: 2,6 OO XX O X X O O X O O O X OO XX Covalent Bond Many simple molecules exist in which two adjacent atoms share one pair of electrons, also known as a single covalent bond (or single bond) Chlorine, Cl2 XX ii d i 12.6 11 Water, H2O Methane, CH4 Ammonia, NH3 1 Hydrogen Chloride, HCl Exercise Try to draw the cross and dot diagram for 1. PCl3 2. H2S Properties of Simple Molecular Compounds Small molecules are compounds made up of molecules that contain just a few atoms covalently bonded together They have low melting and boiling points so covalent compounds are usually liquids or gases at room temperature As the molecules increase in size, the melting and boiling points generally increase Small molecules have poor electrical conductivity Fluorine Lower MP & BP Higher MP & BP Iodine Melting and Boiling Point of Simple Molecular Compound Small molecules have covalent bonds joining the atoms together, but intermolecular forces that act between neighbouring molecules They have low melting and boiling points (high volatility) as there are only weak intermolecular forces acting between the molecules These forces are very weak when compared to the covalent bonds and so most small molecules are either gases or liquids at room temperature As the molecules increase in size the intermolecular forces also increase as there are more electrons available This causes the melting and boiling points to increase Electrical Conductivity of Simple Molecular Compound Molecular compounds are poor conductors of electricity as there are no free ions or electrons to carry the charge. Most covalent compounds do not conduct at all in the solid state and are thus insulators Common insulators include the plastic coating around household electrical wiring, rubber and wood Ionic (Example: MgO) Covalent (Example: O2) Metal + Non metal Non metal + Non metal High melting and boiling point due Low melting point and boiling point to strong electrostatic force between due to weak intermolecular force of oppositely charged ion attraction between particles Can conduct electricity in liquid due Cannot conduct electricity due to no to free moving ions free moving ions Ionic Covalent Transfer of electron Sharing of pairs of electron Between metal and non-metal Between non-metal and non-metal High melting and boiling point Low melting and boiling point Can conduct electricity in liquid state Cannot conduct electricity Soluble in water Insoluble in water State the differences between ionic and covalent bonds. Ionic Covalent Giant Covalent Structures (Macromolecules) Include diamond and graphite Show the same features as molecular covalent except the following features: Very high melting points High melting point and boiling point due to the strong bond which requires high amount of heat to break Variable conductivity Diamond and silicon dioxide does not conduct electricity because no free moving ions or electrons are present. Graphite can conduct electricity because it contains delocalized electrons. Types of Covalent Structures Simple Molecular Covalent Giant Covalent Structure (Macromolecules) Low melting and boiling point, due to Very high melting and boiling point, due to weak intermolecular force of attraction strong bond between particles Cannot conduct electricity due to no free Diamond moving ions Cannot conduct electricity due to no free moving ions Graphite: Can conduct electricity due to delocalized electrons All of these macromolecules do not have force of attraction but only with strong bonds. Silicon dioxide Allotropes of Carbon Diamond and graphite are allotropes of carbon which have giant covalent structures Both substances contain only carbon atoms but due to the diﬀerences in bonding arrangements they are physically completely diﬀerent Giant covalent structures contain billions of non-metal atoms, each joined to adjacent atoms by covalent bonds forming a giant lattice structure Structure of Diamond In diamond, each carbon atom bonds with four other carbons, forming a tetrahedron All the covalent bonds are identical, very strong and there are no intermolecular forces Properties of Diamond It does not conduct electricity ○ All the outer shell electrons in carbon are held in the four covalent bonds around each carbon atom, so there are no freely moving charged particles to carry the current thus it cannot conduct electricity It has a very high melting point ○ four covalent bonds are very strong and extend in a giant lattice, so a very large amount of heat energy is needed to break the lattice It is extremely hard and dense ○ Strong covalent bonds in its tetrahedral structure which makes it diﬃcult to slide past each other. So, it is very useful to make extremely tough material such as cutting tools Structure of Graphite Each carbon atom in graphite is bonded to three others forming layers of hexagons, leaving one free electron per carbon atom which becomes delocalised The covalent bonds within the layers are very strong, but the layers are attracted to each other by weak intermolecular forces Properties of Graphite Each carbon atom is bonded to three others forming layers of hexagonal-shaped forms, leaving one free electron per carbon atom These delocalised electrons exist in between the layers and are free to move through the structure and carry charge, hence graphite can conduct electricity ○ So, it is used to make non-reactive electrodes for electrolysis The covalent bonds within the layers are very strong which makes graphite to have very high melting point The layers are connected to each other by weak forces only, hence the layers can slide over each other making graphite slippery and smooth ○ Graphite is used in pencils and as an industrial lubricant, in engines and in locks Diamond Vs Graphite Diamond Graphite 1 carbon bonded 4 other carbon to form 1 carbon bonded 3 other carbon to form tetrahedral structure hexagonal layered structure Hard due to strong bond in tetrahedral structure Soft and slippery due to weak force of attraction which is not easy to slide past each other between hexagonal layers which is easy to slide past each other Cannot conduct electricity due to no free moving Can conduct electricity due to delocalized ions/ electrons electron. Diamond Vs Graphite Diamond Graphite Functions of Macromolecules Diamond Used as drilling tool because it is very strong and hard Used as jewelry and decoration due to its shiny and bright appearance. Graphite Used as lubricant and pencil because it is soft and slippery Used as electrode because it is a good conductor of electricity Metallic Bond Interaction of regularly arranged structure Positive metal ions and Sea of delocalized electron Properties of Metallic Bond Metals have high melting and boiling points ○ There are many strong metallic bonds in giant metallic structures between the positive metal ion and delocalised electrons ○ A lot of heat energy is needed to break these bonds Metals conduct electricity in solid ○ There are delocalized electrons available to move through the structure and carry charge Metals are malleable and ductile ○ Layers of positive ions can slide over one another and take up diﬀerent positions ○ Metallic bonding is not disrupted as the outer electrons do not belong to any particular metal atom so the delocalised electrons will move with them ○ They can be hammered and bent into diﬀerent shapes or drawn into wires without breaking e e e e e 2+ 2+ 2+ 2+ e e e e e 2+ 2+ 2+ 2+ e e e e e 2+ 2+ 2+ 2+ Summary Name of Giant or Particles in Elements or Typical state at Example structure simple? structure compound? 20ºC Ionic Giant Positive and Compound Solid NaCl negative ions Metallic Giant Positive ions Element Solid Iron and negative electrons Giant covalent Giant Atoms Element or Solid Diamond Compound Simple covalent Simple Atoms Element or Gas or liquid H2O compound Summary Name of Giant or Particles in Elements or Typical state at Example structure simple? structure compound? 20ºC Ionic Metallic Giant covalent Simple covalent Structure and the Periodic Table Periodic table shows the type of element at 20ºC. The metal elements have giant metallic structures. Most non-metal exist as molecules. A few non-metal elements have giant covalent structure. Question 3 Giant ionic Giant metallic Similarities Differences

Cambridge Lower Secondary Year 9 Science Chapter 2 Properties of Materials PDF

Document Details

Tags

Related

Summary

Full Transcript

Upgrade to continue