1. Chemistry 3 Revision Cards Name________________________

2. Important information on Acids / Bases Definitions: Acid (H + Ions Present) – Proton Donor Base (OH - Ions Present) – Proton Acceptor Neutral – Same number of H + and OH - Ions. Strong – Complete Ionisation Weak – Partial Ionisation Distinguish between Strong and Weak. Either: Universal Indicator: Strong acid red and weak acid yellow. Strong alkali purple and weak alkali blue. Conductivity test – Strong will conduct better than the weak. More Ions! Titration – (4 marks) Use the word Burette. Add amount slowly. Wait until indicator changes Record amount. Wash equipment Repeat for reliability. Arrhenius, Bronsted and Lowry Arrhenius was first (student)– molecules ionise in water. B&L more experienced recognised scientists.

3. Indicators!! Phenolphthalein: Strong Alkali – Weak Acid titrations. Methyl Orange: Strong Acid – Weak Alkali titrations. Any acid / base indicator: Strong acid – Strong Alkali titrations Common question: Dry hydrogen chloride is not acidic – why? Hydrogen and chlorine are bonded. No Hydrogen Ions present.

4. Titration Calculations 1. NaOH(aq) + HCl (aq)  NaCl(aq) + H 2 O(l) 15cm3 of HCl of 2.5M was used to neutralise 10cm3 of NaOH. What is the concentration of NaOH? HClNaOH 15 x 2.5 = 37.5 37.5 / 10 = 3.75m 2. H 2 SO 4 (aq) + 2NH 3 (aq) → (NH 4 )2SO 4 (aq) The student found that 25.0 cm3 of ammonia solution reacted completely with 32.0 cm3 of sulfuric acid of concentration 0.050 moles per cubic decimetre. H 2 SO 4 2NH 3 32x0.05 = 1.61.6 / 25 = 0.064 x 2 = 0.128m If you were asked to work out the molarity of H 2 SO 4 instead of x 2 you would have to divide by 2. Check your work booklet for examples.

5. Metal Ions – Flame Test Remember all Metal Ions are positively charged (+) Flame tests – Using a splint or metal loop dipped in an ionic solution through a flame, will identify which ions are present by the colour of the flame. Lithium (Li+) – Red Flame Sodium (Na+) – Orange Flame Potassium (K+) – Lilac Flame Calcium (Ca2+) – Brick Red Barium (Ba2+) - Green

6. Metal Ions - Sodium Hydroxide Test Other metal Ions can be identified by using Sodium Hydroxide. This method of analysis will normally produce a precipitate (ppt). Calcium (Ca 2+ ) – White Ppt Magnesium (Mg 2+ ) – White Ppt Aluminium (Al 3+ ) - White Ppt – this will dissolve when excess sodium hydroxide is added. Iron (Fe 2+ ) – Green Ppt Iron (Fe 3+ ) – Brown Ppt Copper (Cu 2+ ) – Blue Ppt Ammonium (NH 4 + ) This is a non-metal Add sodium hydroxide – No Ppt will be formed. When the solution is heated, use damp red litmus paper – this will turn blue and a strong smell of ammonia will be present.

7. Non Metal Ions (-) Carbonates (CO 3 2- ) Add an acid to the solution, if carbonates are present CO 2 will be released. This will turn Limewater Cloudy. Carbonates will change colour when heated. Copper Carbonate will change black as it turns into copper oxide. Zinc Carbonate will change from white to yellow as it changes into zinc oxide. Sulphate Ions (SO 4 2- ) Using Barium Chloride will produce a white Ppt if sulphate ions are present.

8. Non Metal Summary Carbonates CO 3 2- – Add Acid – CO 2 – Limewater Milky Sulphate SO 4 2- – Barium Chloride - White Ppt Nitrates - NO 3 - – Add Aluminium Powder and sodium hydroxide, gas produced – damp red litmus will turn blue. Halides – Group 7 – Add Silver Nitrate. – Chlorides – White ppt – Bromide – Cream ppt – Iodide – Yellow ppt

9. Organic Chemicals Saturated and Unsaturated Hydrocarbons Identified using Bromine Water Saturated – Bromine water will stay orange / brown. Unsaturated will de-colourise. Instrumental Analysis Quick, Accurate, sensitive and can detect very small amounts. Expensive and need special training to use them. The chemical analysis we have just looked at can not detect very small amounts. There are lots of different methods of instrumental analysis – if in doubt use Mass Spectrometer. This is used to test for elements and + ions.

10. Instrumental Analysis Atomic Spectrometers- Analyse energy (usually light) using a line emission spectrum. Patterns are then matched to existing frequencies on a database. Infrared Spectroscopy- detects specific bonds in organic chemicals. Ultraviolet Spectroscopy- analyses levels of nitrates + phosphates in water. Nuclear magnetic resonance spectroscopy- used to detect some organic molecules. Gas-liquid chromatography- separates gases through a column- identified by the distance moved. Mass spectrometer- used for elements + ions. Masses of fragments used to identify what is in a sample.

11. Empirical Formula of an Organic Hydrocarbon 0.42g of a hydrocarbon burns in oxygen to form 1.32g of CO 2 and 0.54g of H 2 O Mr of CO 2 is 44g. Mr of H 2 O is 18g. 1.32 / 44 = 0.03 0.54 / 18 = 0.03 Formula = CH 2

12. Development of the periodic table – Newland and Mendeleev Newlands, and then Mendeleev, attempted to classify the elements by arranging them in order of their atomic weights. PROBLEM WITH NEWLANDS Newland attempted to arrange the elements with his law of octaves. (repeating pattern of similar properties every eighth element) e.g. 7 elements between fluorine and chlorine and 7 between sodium and potassium on his table. His work was ignored because it had inconsistencies, he did not leave gaps for undiscovered elements, some boxes had 2 elements in and metals and non metals were mixed up.

13. Mendeleev Classified elements in terms of atomic weight. He left gaps for undiscovered elements He did reverse some elements to allow them to be in groups with similar properties. Problems with Mendeleev His work was not accepted at first because: Some boxes had two elements in them Copper and Silver were in group 1 and were unreactive unlike the others There are metals and non metals in the same group He left spaces / Gaps Some chemists thought there were no more elements to discover

14. Modern Periodic Table Elements in the same group have similar properties – same number of electrons in the outer shell. Reactivity?

15. Different Groups Group 1 – Alkali Metals – Add to water with Universal indicator in it you will see: – Fizzing, flame, UI will turn purple (alkali) hence the name Transition Metals -Have coloured compounds, good conductors of heat and elec, And coloured compounds. Similar to each other because: 1.(transition elements usually) have same / similar number of outer / 4 th shell electrons 2. inner (3rd ) shell / energy level is being filled No transition metals between Magnesium (Grp 2) and Aluminium (Grp 3) because they are in period 2 which only have 2 levels and the 2 nd level will only contain 8 electrons. Compared to Alkali Metals − Have higher melting points (except for mercury) and higher densities − Are stronger and harder − Are much less reactive and do not react as vigorously with water or oxygen.

16. Water Water is made safe to drink by removing solids and killing bacteria. Filtering and bubbling with chlorine. Tiny particles which are still suspended in the water are allowed to settle out and fall to the bottom. This is called sedimentation. Chlorine is bubbled through the water to kill bacteria. Chlorine Water filters containing carbon, silver and ion exchange resins can remove some dissolved substances from tap water to improve the taste and quality. Pure Water can be produced by DISTILLATION. However Drinking distilled water removes ions from the body which may lead to health problems.

17. Hard WaterSoft Water Contains Magnesium and Calcium Ions – from rock. Calcium is good for health. Hard water reacts with soap to form scum and so more soap is needed to form lather. Scum is created when calcium ions form an insoluble precipitate. Using hard water can increase costs because more soap is needed. When hard water is heated it can produce scale that reduces the efficiency of heating systems and kettles. Soft water readily forms lather with soap. Water that doesn't have enough calcium and magnesium. Soft water is treated water in which the only positively charged ion is sodium. Sodium is bad for health – high blood pressure. Soft water, on the other hand, may taste salty and may not be suitable for drinking.

18. Hard to Soft Making Hard Water Soft by removing dissolved calcium and magnesium ions Adding sodium carbonate which reacts with the calcium and magnesium ions forming a precipitate of calcium carbonate and magnesium carbonate. This can then be filtered. Using an ion exchange column containing hydrogen ions or sodium ions which replace the calcium and magnesium ions when hard water passes through the column.

19. Solubility of Solids Solute – Solid being dissolved, Solvent – Liquid Solution – Dissolved Solute in Solvent Solubility of a solute increases as the temperature of the solvent increases. Once no more solute will dissolve the solution is said to be SATURATED.

20. Solubility of Gases Many gases are soluble in water. Their solubility increases as the temperature decreases and as the pressure increases. Gases as a solute dissolve and remain dissolved better under colder temperatures. Think - as the temperature decreases, so does the kinetic energy of any particle and just the opposite. Consider a can of coke. If it gets hot - the pressure inside increases, as it does if you shake it. If you then open it, there is a sudden rush of carbon dioxide coming out of solution and soda shoots everywhere. If you open an almost frozen soda can, there is little release of gas from solution. Dissolved oxygen is essential for aquatic life. If the temperature of the water increases, the amount of oxygen that is dissolved decreases.

21. Chemical Reactions BONDS BREAK (need energy) requires the Activation Energy (catalysts reduce this) BONDS FORM (releases energy) EXO – More energy is released from the forming of new bonds than was required to break existing ones. ENDO – More energy is required to break existing bonds than is released when new bonds form. Calculating bond energies If the bond energies of the reactants are greater than the products – the reaction is endothermic If the bond energies of the reactants are lower than the products – the reaction is exothermic.

22. Endothermic and Exothermic Reaction Graphs Exothermic A = Activation Energy – amount of energy required to break existing bonds B = Total amount of energy released from the production of new bonds. C= Amount of energy given out into the surroundings Endothermic A = Activation Energy – amount of energy required to break existing bonds B = Total amount of energy released from the production of new bonds. C= Amount of energy absorbed from the surroundings