Atoms and Reactions (Module 2.1)

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nucleus and that it was very small and positively charged. concluded that atom is mainly empty space in a 'negative cloud. Gold leaf experiment: positive a-particles fired at thin gold leaf. most went through (mainly empty space). small number deflected back (hit positive nucleus) • Niels Bohr (1913): discovered problem with Rutherford's model: cloud of electrons could collapse into positive nucleus. proposed that electrons were in fixed energy shells. PROOF: when EM radiation is absorbed, electrons move between shells and emit this radiation when moving down to lower energy shells. •Model Today: electrons don't have the same energy in shells. they have subshells. this explains ionisation trends. FORMING COMPOUNDS o ions are created when electrons are transferred from one atom to another. They attract each other to form compounds. group 1 form +1 ions group 2 form +2 ions molecular ions: group 7 form 1-ions group 3 form 3+ions group 5 form 3-ions group 6 form 2-ions water molecules can exist within crystal structures. this is called water of crystallisation for 0 salts are examples of ionic compounds. they can be hydrated (with water) or anhydrous (without water) H>0 - Na* Cl every = 0.0083 0.0083 = 1 0 mole of salt there are a certain number of moles of water of crystallisation ammonium (NH4) hydroxide (OH) • sulfate (504) water is polar and is attracted to ions in a salt. the & Oxygen attracts positive ions and Hydrogen attracts negative ions. = 0.0416 0.0083 : 5 = H H- NOTE: Zinc is Zn nitrate (NO₂) WATER OF CRYSTALLISATION: 1.88g of hydrated Ca504.xH₂0 was heated until there was no more water of crystallisation was left in the sample. mass of anhydrous is 1.13g. Work out the value of 'x'. 1. write molecules involved: Ca504 H₂0 2. Write masses 1.13g 0.75g 3. divide by Mr to get number 136 18 of moles 4. divide by smallest no. moles x=5 24 " 1.88-1.13-0.75 • written as .xH₂0. carbonate (CO₂) Silver is Ag CaSO4.5H₂0 IONIC EQUATIONS: • ionic equations show the ions that are formed in solution and show which particles are reacting. H₂SO4 + 2KOH 2H+ 50 2K 20H→→→→→2K+ SO4+H₂O + + Final ionic equation: 2H+ 20H→→ 2H₂O SO and K* are spectator ions - don't get involved in the reaction. K₂SO4 + 2H₂O not an ion THE MOLE-AVOGADRO'S NUMBER mol-unit for amount of substance- Number of Particles Avogadro's number x Number of moles 8 Number of moles - Concentration (moldm³³) ÷ Volume (dm³) 8 p = pressure (Pa) V= volume (m³) Number of moles Mass (g) ÷ Molecular mass IDEAL GAS EQUATION: this is the number of moles in a specific volume of a gas PV = nRT 1 mol contains 6.02 x 10 atoms/molecules [Avogadro's number] n= moles (mol) R= gas constant (8.314 JK"mol) NOTE: -Standard Conditions: 298K (25°C) 100 kPa T-temperature (k) ATOM ECONOMY: shows how efficient a reaction is. • atom economy = Mr of desired product ÷ Mr sum of all products x 100 • high atom economies produce less waste- benefits the environment higher atom economy means that raw materials are used more 0 companies will try to use reactions that tend towards 100%. atom -more sustainable efficiently. economy • higher atom economy = less by products = less money spent seperating products. 0 PERCENTAGE YIELD: • percentage yield = actual yield ÷ theoretical yield x 100 ACIDS AND BASES • acids are proton donors, bases are proton acceptors. an alkali is a soluble base. produces OH" ions when added to water. • o when acids and bases react with each other a reversible reaction is formed. Acids: HA+H₁₂0= H₂0* + A¯ o when acids react with bases it forms salts which are pH neutral. the H' ions produced from acids react with OH ions produced from alkalis makes water which is neutral. salts are made from the metal of the base (or NH4+) and the non-metal of the acid. • ammonia reacts with acids to make ammonium salts but no water. e.g→ 2NH3 (aq) + H₂SO4 (aq) = (NH4)₂ 504 (aq) • ammonia doesn't directly produce OH ions. it needs to react with water to produce NH4 ions and OH ions. 0 Bases: B+H₂0BH*+ OH" •WEAK ACID: e.g. ethanoic acid Equation: CH₂COOH=CH₂C00+ H* WEAK BASE: e.g. ammonia Equation: NH₂ + H₂O⇒ NH,₂” + OH” backwards reaction favoured so not many H*(protons) produced. → backwards reaction favoured so not many OH ions produced. STRONG ACID: e.g. hydrochloric acid, sulfuric acid, nitric acid Equation: HCI⇒ H*+CI™¯ · forward reaction favoured strongly. lots of Hª produced. 0 0 H²(aq) + OH (19₂) = H₂0 (1) STRONG BASE: .g. sodium hydroxide, potassium hydroxide Equation: NaOH⇒ Na*+OH¯ → forward reaction strongly favoured. lots of OH ions produced. ACIDS AND METALS: 1. metal + acid→→→→salt + hydrogen 2. metal oxides + acid 3. metal hydroxides + acid salt + water TITRATIONS: used to work out the concentration of an acid or alkali. 1. have an acid/alkali in burette with known concentration. salt water 4. metal carbonates + acid salt + carbon dioxide + water 2. I have an acid or alkali of an unknown concentration but known volume in conical flask. add a few drops of indicator. 3. add chemical in the burette to the conical flask until indicator changes colour. this is known as the end point. →add drop by drop near end point. NOTE: read how much chemical was added from the burette to neutralise the solution in the conical flask. → record results to 2 decimal places and repeat until you get two concordant titres MAKING STANDARD SOLUTIONS: weigh out exact amount of solid using a balance and weighing boal → transfer solid from weighing boat to beaker. wash remaining solid into beaker using deionised water. → dissolve solid fully using deionised water. stir so everything is fully dissolved. → keep filling solution with deionised water until it reaches the graduation line. do not go above this line. REDOX • reduction and oxidation takes place when electrons are transferred. oxidation is the loss of electrons. reduction is gain of electrons. ] • reduction is the decrease in oxidation number, oxidation is the increase metals are oxidised when they react with acids to form a salt and hydrogen gas. OXIDATION NUMBERS: → elements can have different oxidation states. the oxidation number changes according to these rules: uncombined elements →→→ oxidation number state: 0 oxidation number /state: same as charge on ion oxidation number /state: +1 e.g. KCI ions group 1 aluminium oxidising agents are reduced (they gain electrons) reducing agents are oxidised (they lose electrons) oxygen e.g. Cl₂, Fe, 0₂ e.g. Cl¯=-1, group 2 oxidation number /state: +2 e.g. Cao hydrogen oxidation number Istate: +3 e.g. Al ₂03 oxidation number/state: +1,-1 in hydrides e.g. HF (+1), NaH(-1) chlorine →→→→ oxidation number Istate: -1, has positive value for Fand 0 e.g. KCI (-1), CIF, (+3) fluorine oxidation number state: -1 → e.g. KF → oxidation number /state: -2,-1 in peroxides, +2 in OF₂ e.g. Li₂0 (-2), H₂0₂ (-2) Ca²+ = +2