Chemistry – Sublimation - Change either from solid to gas or gas to solid – Condensation - Change from gas to liquid – When a substance is heated, particles gain energy and moves faster then overcome the attraction forces between them, and move further apart – When a substance is cooled, particles lose kinetic energy and move slower, and come nearer to each other – The temperature is constant when heated because the energy supplied is used to overcome the forces between the particles – The forces of attraction between the particles are weakened enough so that the particles move far away from each other – Temperature o a substance is directly proportional to the kinetic energy of the molecules – When temperature of gas is increased the pressure of the gas will also increase – Pressure of the gas is increased when its compressed into a smaller volume, the molecules gets closer and hit the wall of the container more often – Diffusion - movement of particles (liquid, gas) from an area of high concentration to an area of low concentration – Particles that are lighter (lower molecular mass) diffuse faster – Particles diffuse faster at higher temperature – Diffusion is faster in gasses than in liquids – White solid ⽩⾊沉淀 tend to form closer to the heavier compound – Crystallisation: – 1. Heat to the point of crystallisation (evaporate) – 2. Cool to form crystals – 3. Filter through filter paper and funnel – 4. Dry the crystals between filter papers – Insert a glass rod into the hot solution and remove it. Crystals should form on the tip – Example: Obtain soluble salt rom aqueous solution – Calculate the percentage of different compound in a mixture: – (Initial Mass - X Compound to measure) / Total Mass of the mixture * 100% – To plan an experiment to investigate the percentage of copper (II) oxide in the mixture: – 1. Weight 10g of black solids using a balance – 2. Put the mixture in to a beaker – 3. Add access dilute sulphuric acid and stir with a rod – 4. Copper oxide dissolves, while carbon does not react – 5. Filter through a filter paper and funnel – 6. Carbon is collected as a residue – 7. Dry the carbon between the filter paper – 8. Weight the carbon – Use the formula: (Initial Mass - Carbon Mass) / 10g * 100% – Simple Distillation (salt water) – Thermometer - To measure the temperature – Flask - Store the impure solution to evaporate – Rubber Stopper (cork/bung) – Condenser - Condensation happens here – Flask/Beaker - Store the filtered liquid – Tripod Stand – Fractional Distillation (ethanol + water solution) – Fractionating Column - Glass bead to hinder the water vapour to move up quickly – Thermometer - To measure the temperature – Condenser - To condense the water – Beaker/Flask - To store the impure liquid (ethanol+water) and collect the filtered liquid – Tripod Stand An experiment of separating mixture of Ethanoic Acid & Chloroethanoic using Fractional Distillation: Provided Information: 1. - Ethanoic Acid, boiling point 118 Celsius - Chloroethanoic acid, boiling poin 190 Celsius Which liquid would be collected first ? Ethanoic Acid, because ethanoic acid will reach its boiling point before chloroethanoic acid. 2. How would we know when all of this liquid has been collected ? When the temperature start to increase 3. Suggest why small glass beads are used in the fractionating column instead of large glass beads. Small beads have larger surface area to prevent bumping of solution. Techniques used in the separation of mixtures include: – Pure water from sea water - Simple distillation – Ethanol from a mixture of ethanol and water - Fractional distillation – Calcium carbonate (insoluble solution) from a mixture of calcium carbonate and water Filtration – CuSO4 + 5H2O(s) from CuSO4(aq) - Crystallisation Chromatography – Ink/Food colouring solvent: use Water as solvent – Other substances tat does not dissolve in water: use ethanol as solvent – If the solvent does not move up (stay at baseline), spot x does not dissolve in that solvent – (try using ethanol) – Baseline must be above the solvent: or else it will move down instead of moving up – More soluble solvent will move up quicker than less soluble solvent Retention factor/Retardation Factor = distance traveled by the spot / Distance traveled by the solvent = X / Y Unit 3: Atomic Structure Atomic number/Proton Number/ Number of protons: It is the number of protons in the nucleus of an atom Mass Number/Atomic Mass/Nucleon Number: Total number of protons and neutrons in the nucleus When atom loose electron, it become positively charged (cation) When atom gain electron, it become negatively charged (anion) Atom – Number of proton is equal to number of electrons (neutral) – Number of proton is equal to number of electrons (neutral) Isotopes – Atoms of the same element having the same number of protons but different number of neutrons – Isotopes of an element have the same chemical properties because they have the same number of electrons in the outer shell Relative atomic mass – An average of the mass number of the different isotopes in the ratio of their occurrence in nature relative to the mass of an atom E.g. Chlorine has two main isotopes: Cl-35: 75%; Cl-37: 25% – Relative atomic mass of chlorine = (35 * 75) + (37 * 25) / 100 = 35.5 E.g The table shows the percentage composition by mass of a sample of silicon Isotope Percentage (%) 28Si 92.2% 29Si 4.70% 30Si 3.10% Calculate the relative atomic mass of this sample of silicon. Give your answer to one decimal place. Relative atomic mass of this sample of silicon: (28 * 92.2) + (29 * 4.7) + (30 * 3.1) / 100 = (2581.6 + 136.3 + 93) / 100 = 28.1 Unit 4: Ionic, covalent, and metallic bonding Ionic Bonds – The bond between metals and non-metals – Non-metals gain electrons to form negative ions They bond together due to the strong electrostatic attraction forces between positive and negative ions. Ionic Crystal Lattice – ionic crystal is a crystalline form of an ionic compound. – solids consisting of ions bound together by their electrostatic attraction into a regular lattice – giant three dimensional structure – regular arrangement of alternating positive (e.g Na+) and negative ions (Na-) – Exam definition: Each positive ion is surrounded by 6 negative ions and each ion is surrounded by six positive ions – The structure is held together by strong electrostatic attraction forces between the oppositely charged ions – Negative ions are bigger than positive ions (electrons are bigger than protons) Characteristics of Ionic compounds – Have high melting point and boiling point, because of strong attraction forces between positive and negative ions in the giant crystal lattice which need high temperature to break them and separate the ions – Increased charges on the ions will have stronger attraction forces which needs more energy to be broken – Less volatile: Not easy to evaporate – Dissolve in water – unable to dissolve in organic solvents such as Ethanol – Able to conduct electricity when molten or in solution, because positive and negative ions are free to move – Do not conduct electricity when solid because the ions are not free to move (both positive and negative ions are conductive) Valency = Valance Electron = Outer Shell = Group Number E.g Valency of scandium is three = The outer shell of scandium is three Covalent Bond – Formed between two non-metals – Sharing of one or more pair of electros to obtain full outer shell – Strong attraction forces between the bonding pair of electrons and the positive nuclei of the atoms – low melting point and boiling point, because of weak attraction forces between molecules – More volatile (evaporate easily) – Insoluble in water but dissolvable in organic solution (ethanol) – Do not conduct electricity, absence of free ions – Usually liquids or gases at room temperature, except bromine which is a reddish brown liquid Simple Molecular Structure – Low melting point due to weak attraction forces between molecules – Strong covalent bods, weak intermolecular forces between molecules – E.g. Iodine, Carbon Dioxide – Weak attraction forces between molecules = weak van der waals forces between molecules Giant Macromolecular Structure – a large three dimensional structure in which many atoms are strongly bonded together – High melting point due to strong bonds between atoms in the giant structure – E.g. Ionic crystals, Metallic structure, diamond & graphite, silicon dioxide – Ionic Crystal Lattice – Three dimensional structure with regular arrangement of alternating positive and negative ion – Strong electrostatic attraction forces between oppositely charged ions – Can conduct electricity when molten or in aqueous – Soluble in water – High melting and boiling point – Metallic Structure – Thee dimensional structure with rows of closely packed positive ions surrounded by a sea of delocalised electrons (free to move electrons) – Have high melting and boiling points, due to the strong electrostatic attraction forces between the positive ions and negative electrons Physical properties of all metals – Can conduct electricity both solid and molten due to the presence of free moving electrons – Malleable (can be spread into sheets) and ductile (can be stretched into wires) because the layers of positive ions can slide over each other when hit with a hammer – Shiny Metal is malleable: the layers of positive ions can slide over each other when hit with a hammer, especially when the metal is hot since the atoms have more energy and are vibrating faster Titanium is a good conductor: they have closely packed positive ion surrounded by free moving electrons Typical Metal: has strong electrostatic attraction forces between the closely packed positive ions. and delocalised (free to move) electrons Attractive forces are present in a metallic structure: strong electrostatic attraction forces between closely packed positive ions and delocalised electrons Diamond & Graphite – Allotropy i the existence of an element in two or more different forms in the same physical state – The different forms are called allotropes – e.g. Carbon exists. in two forms (allotropes): Diamond &. Graphite Diamond – Has giant three-dimensional structure – Carbon atom is covalently bonded to four other carbons in a rigid tetrahedral structure – Hard because of the strong covalent bonds between carbon atoms in the rigid threedimensional tetrahedral structure – Does not conduct electricity because of the absence of free electrons – Used in cutting, drilling equipment Graphite – Has giant three-dimensional structure – Carbon atom is covalently bonded to three other carbon atoms in layers of 6-sided rings – Weak attraction forces between the layers – Soft and used as lubricating agent, due to the weak attraction forces between the layers that allows them to slide past each other – Can conduct electricity due to the presence of free moving electrons, since each carbon atom is covalently bonded to three other carbon atoms and the fourth electron in each carbon is free to move between the layers Silicon Dioxide (sand) SiO2 – Has giant three dimensional tetrahedral structure (same as diamond) – Each silicon is covalently bonded to four oxygen atoms – Each oxygen is covalently bonded to two silicon atoms – Solid – Have high melting point Cation Formula of cation Sodium ion Na+ Potassium ion K+ Aluminium ion Al3+ Zinc ion Zn2+ Magnesium ion Mg2+ Iron(II) ion Fe2+ Iron(III) ion Fe3+ Copper(II) ion Cu2+ Calcium ion Ca2+ Sliver ion Ag+ Lead(II) ion Pb2+ Ammonium ion NH4+ Anion Formula of anion Oxide ion O2- Chloride ion Cl- Bromide ion Br- Iodide ion I- Hydroxide ion OH- Carbonate ion CO3 2- Nitrate ion NO3- Sulphate ion SO4 2- Phosphate ion PO4 3- Manganate (VII) ion MnO4 - Thiosulphate ion S2O3 2- Dichromate(VI) ion CR2O7 2- – Valency of Transition Metals are indicated between brackets: Zinc (II) – Valency of Zinc = 2 Name Formula Hydrochloric acid HCL Nitric acid HNO3 Sulphuric acid H2SO4 Phosphoric Acid H3PO4 Carbonic Acid H2CO3 (s) to indicate solid (g) to indicate gas (l) to indicate liquid (aq) to indicate dissolved in water Unit 6: The mole (amount of substances) – Mole: the relative atomic mass in grams – Molecular formula: number and type of different atoms in one molecule – Relative atomic mass (RAM): (atomic mass number) average of. the mass number of different isotopes in the ratio of their occurrence in nature – Relative molecular mass: (mass of one mole) is the sum of the relative atomic masses of the atoms in the molecule – E.g. SO2, CO2 To determine the mass of one mole (molecular mass) – Magnesium phosphate – 1. Get the compound formula 配平 Mg3(PO4)2 – 2. Get the mass number of Magnesium & Phosphate (Phosphorus & Oxygen) – Magnesium: 24 – Phosphorus: 31 – Oxygen: 16 – 3. Solve the equation Mg3(PO4)2 – 3(24) + [ (31 * 2 ) + ( 16* 2 ) ] – = 262g/mol To determine the empirical formula and molecular formula – Empirical formula: Simplest whole number ratio of the different atoms or ions in a compound – E.g. CH2O – Molecular formula: Number and type of different atoms in one molecule – E.g. C6 H12 O6 Question: A compound has 18.2% potassium, 59.4% iodine and the rest is oxygen. Determine its empirical formula 1. Find the Mass/Percentage of each element – Potassium (K): 18.2% – Iodine (I): 59.4% – Oxygen (O): 100% - 18.2% - 59.4% – 22.4% 2. Get the Nucleon Number (mass number) of each element – Potassium (K): 39 – Iodine (I): 127 – Oxygen (O): 22.4 3. Calculate the mole of each element – Potassium (K): 18.2 / 39 – 0.47 mole – Iodine (I): 59.4 / 127 – 0.47 mole – Oxygen (O): 22.4 / 16 – 1.4 mole 4. Find the simplest ratio – Take the smallest mole: 0.47 – Potassium (K): 0.47 / 0.47 – 1 – Iodine (I): 0.47 / 0.47 – 1 – Oxygen (O): 1.4 / 0.47 – 2.98 ~ 3 5. The empirical formula: – KIO3 Simplest ratio rules: – only 0.1 can be round off/up (e.g 2.9 ~ 3) – if the simplest ratio is 2.4, we’ll have to multiply it until it reaches a whole number or able to be rounded, in this case we need to multiply by 5 ( 2.4 * 5 = 12 ) – when one of the mole is multiplied, the other must be multiplied too Molecular formula of the oxide of nitrogen (NO2) ? – It has a relative molecular mass of 92 – The empirical formula is N1O2 – ( NO2 )n = 92 ( 14 + 32 )n = 92 46n = 92 n=2 – N(1*2) O(2*2) – N2O4 Number of moles (m/Mr) solid = mass (g) / Molecular Mass Number of moles (v/24) volume of gas at room temperature = volume(dm3) / 24 – 1. Find the number of moles – 2. Use the formula to find the volume Number of moles (C*V) = Concentration * Volume (dm3) Concentration (mol/dm3) = number of mole / volume What is the concentration, in mol dm-3, of a solution containing 7.84g of phosphoric(V) acid, H3PO4 in 400cm3 of solution 1. The mass of H3PO4 is 7.84g 2. Get the nucleon number (mass number) of H3PO4 – Hydrogen (H): 1 – Phosphorus (P): 31 – Oxygen (O): 16 3. Solve the equation – 7.84 / [ 1(3) + 31(1) + 16(4) ] – = 0.08 moles 4. Now that we have the moles, we can solve the concentration – concentration = number of mole / volume = n / v(dm3) – = (0.08 moles / 400) * 1000 – = 0.20 mole/dm Avogadro’s law – states that equal volumes of gases (at the same temperature and pressure) contain the same number of molecules (6.02 * 10^23) – Number of moles = number of molecules / 6.02 * 10^23 The Mole: Relative atomic mass in grams The Avogadro constant: Number of particles in one mole of substance Chemical Calculations – 2H2(g) + O2(g) —-> 2H2O Balanced equation tell us 1. The ratio of the number of mass of the reactants and products 2. Mole of hydrogen reacts with 1 mole of oxygen to produce 2 moles of water 3. The physical state of the reactants of products Water is formed when 48g of oxygen combine with 6g of hydrogen What mass of oxygen combines with 2g of hydrogen ? – 16g Thermal decomposition of magnesium carbonate (Mr = 84) Which mass of magnesium oxide is formed when 21.0g of magnesium carbonate are completely decomposed – MgCO3 —> MgO + CO2 – n = 21 / 84 – = 0.25 – 1 mole of MgCO3 —> 1 mole MgO – 0.25 * (24+16) = 10.0g
