Chemistry 3 Revision Cards

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) + H2O(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. H2SO4 (aq) + 2NH3 (aq) → (NH4)2SO4 (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. • H2SO4 2NH3 • 32x0.05 = 1.6 1.6 / 25 = 0.064 x 2 = 0.128m • If you were asked to work out the molarity of H2SO4 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 (Ca2+) – White Ppt • Magnesium (Mg2+) – White Ppt • Aluminium (Al3+) - White Ppt – this will dissolve when excess sodium hydroxide is added. • Iron (Fe2+) – Green Ppt • Iron (Fe3+) – Brown Ppt • Copper (Cu2+) – Blue Ppt • Ammonium (NH4+) 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 (CO32-) • Add an acid to the solution, if carbonates are present CO2 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 (SO42-) • Using Barium Chloride will produce a white Ppt if sulphate ions are present.

8. Non Metal Summary • Carbonates CO32- – Add Acid – CO2 – Limewater Milky • Sulphate SO42-– Barium Chloride - White Ppt • Nitrates - NO3- – 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 CO2 and 0.54g of H2O • Mr of CO2 is 44g. Mr of H2O is 18g. • 1.32 / 44 = 0.03 0.54 / 18 = 0.03 • Formula = CH2

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: • (transition elements usually) have same / similar number of outer / 4thshell • 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 2nd 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. • 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.

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