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Atomic Weights At the end of this you should be able to… · Describe the mole as the SI unit for amount of substance · Relate amount of substance to relative.

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Presentation on theme: "Atomic Weights At the end of this you should be able to… · Describe the mole as the SI unit for amount of substance · Relate amount of substance to relative."— Presentation transcript:

1 Atomic Weights At the end of this you should be able to… · Describe the mole as the SI unit for amount of substance · Relate amount of substance to relative atomic mass · Describe relationship between the mole and Avogadro’s number · Conceptualise the magnitude of Avogadro’s number · Describe the relationship between molar mass and relative molecular mass · Calculate the molar mass of a substance given its formula

2 Chemical Change - Reactions Objectives: At the end of this section you should be able to:- Explain the difference between chemical and physical change. Use the periodic table to determine valency. Explain the significance of the law of constant composition for writing chemical reactions.*** (comeback to) Write and balance chemical formulae. Calculate reacting masses for given reactions. Explain energy changes in reactions. Explain Avogadro’s law. Work out reacting volumes of gases.

3 Chemical Change - Reactions Objectives: At the end of this section you should be able to:- Distinguish between synthesis and decomposition reactions Discuss how energy released by chemical reactions is essential for life Describe how metals react with oxygen Explain the reverse process in which metals are separated from their oxides Describe the reactions of non-metals with oxygen Write balanced equations for reactions involving non- metals.

4 Reacting Masses 1.2Na+Cl 2 -->2NaCl m(NaCl) = 2( ) = g 1.C+Cl 2 -->CCl 4 m( 1.2ZnS+3O 2 -->2ZnO + 2SO 2 2.FeS + 2HCl --> H 2 S + FeCl 2 3.SO 2 + 2H 2 S --> 3S + 2H 2 O Calculate the mass of each underlined compound either produced or required. (Balance the reactions first)

5 Balancing More Reactions 1.Na + H 2 O --> NaOH + H 2 2.H 2 + O 2 --> H 2 O 3.CaCO 3 --> CaO + CO 2 4.CaCl 2 + Na 2 SO 4 --> CaSO 4 + NaCl 5.Al(NO 3 ) 3 + K 2 CO 3 --> Al 2 (CO 3 ) 3 + KNO 3 6.Na 3 PO 4 + MgI 2 --> Mg 3 (PO 4 ) 2 + NaI

6 Relative Mass Atomic Certain products, such as paper for example, are sold by the ream. A ream is 500 sheets. Since it is impractical to actually count out 500 sheets, the weight (mass) of 500 sheets is determined; then each ream is packaged according to this mass. Atoms are even smaller than paper, so it is not possible to actually count them. However, it is possible to know the mass of an atom in respect to the mass of another atom. The Relative mass of an object is expressed by comparing it mathematically to the mass of another object. So the relative mass of an orange in relation to a grapefruit is.6. The relative mass of the grapefruit in relation to a grapefruit is 1.0. Atoms are compared to the lightest atom (hydrogen) which is 12 times lighter (1/12 of the mass of) one carbon atom. THE RELATIVE ATOMIC MASS IS THE NUMBER OF TIMES AN ATOM IS HEAVIER THAN 1/12 OF A C 12 ATOM.

7 The Mole The mole is defined as, “the amount of ………….. with the same number of ……………………… particles as ….. grams of carbon 12”. (n used as symbol for moles) Six hundred and two thousand, three hundred, billion billion ! 6.023x10 23 particles g C Symbol (….) Number of particles = no of moles x no. particles in a mole Particles = ……………..

8 The Carbon Standard Carbon-12 is the standard upon which the relative mass of other atoms is determined. It wasn’t always this way. At first hydrogen was used and it was assigned the atomic mass of one. If you have equal numbers of nitrogen atoms and hydrogen atoms, the nitrogen atoms are 14 times heavier than the hydrogen atoms. Therefore, nitrogen was assigned the atomic number of 14. oxygenLater oxygen was used as the standard with an atomic mass of 16. However, carbon-12 proved to be more convenient to capture and measure in pure form, so it became the standard. However, now even carbon-12 is slowly losing its position as standard, as sophisticated equipment makes it possible to give even more accurate measures of atomic mass. For this reason you will notice that on the periodic table the AMUs are not expressed as exact relative units to carbon-12.

9 Dozen & Particles particles 1 doz dozen  12 x ? 3

10 Moles & Particles particles 1 mol moles (n)  L x x x x ? 3

11 Avogadro’s No Egs If you have 6g of Hydrogen gas how many... A) Molecules B) atoms C) electrons do you have?

12 The Mole and Mass The mole is defined in such a way that the MASS NUMBER (A) of an element is equal to the relative atomic mass mass of one mole of the substance. (in grams) - THE MOLAR MASS Eg Na = 23g/mol, water(H 2 O)=18g/mol ZAXZAX Atomic Number (smaller) Mass Number (bigger) protons + neutrons Periodic Table Symbol Relative atomic mass or mass(g) of one mole

13 Relative Masses Relative atomic(A r ) - The mass of the atom relative to 1/12 of the mass of a C 12 atom. (Number of times heavier than…) O - 16 one atom of oxygen is 16 times heavier than 1/12 of the mass of a C 12 atom, Na - 23 one atom of sodium…, H - 1 etc. Formula mass (M r ) - The sum of all the atomic masses of the atoms in a molecule. Water H 2 O one molecule of water has a relative mass of (2x(1)+16) = 18 - that is the molecular or formula mass of water. M r (H 2 O) = 18 ( Times heavier than…)

14 Relative Atomic Mass ZAXZAX Atomic Number (smaller) Mass Number (bigger) protons + neutrons Relative atomic mass or mass(g) of one mole Periodic Table Symbol Calculate: The mass in grams - 1.of one mole of copper chloride (CuCl 2 ) 2.one mole of carbon dioxide (CO 2 ) 3.One and a half moles of oxygen (O 2 ) 4.TWO moles of methane (CH 4 ) 5.Four moles of water. m = n x M r mass of substance = number of moles x mass of 1 mole

15 Relative Masses - examples Calculate the Relative Atomic Mass of: O 2 (oxygen gas) Cl 2 (chlorine gas) NaCl (sodium chloride - table salt) H 2 SO 4 Sulphuric acid CaCO 3 (calcium carbonate) (NH 4 ) 2 Cr 2 O 7 (ammonium dichromate)

16 Isotopes Isotopes - Atoms of the same element which have different numbers of neutrons. Eg: 6 13 C & 6 12 C Relative atomic mass is (actually) the average mass (of all the isotopes in a random sample) of the atoms of an element relative to 1/12 of the mass of a carbon-twelve atom C 6 protons 6 electrons Neutrons? 6 12 C 6 protons 6 electrons Neutrons?

17 Isotopes Chlorine has two isotopes Cl & Cl Cl(35) has 35-17=18neutrons Cl(37) has 20 neutrons! 37 Cl (25%) & 35 Cl (75%) - exist in the ratio 1:3 Calculate the average mass of a Cl atom. (Two methods)

18 ISOTOPES SymbolPROTONSELECTRONSNEUTRONS Carbon C Carbon C Boron B Boron B Hydrogen 1 Hydrogen 2 Chlorine 35 Chlorine 37

19 Isotopes Boron’s two isotopes 5 10 B & 5 11 B 5 10 B (19.8%) & 5 11 B (80.2%) - exist in the ratio 1:4 Calculate the average mass of a Boron atom. (Two methods) In 100 atoms – 19.8 have a mass of 10 and 80.2 have mass 11! Average A r (B)= total mass = (10x19.8)+(11x 80.2) = 10.8 no of atoms 100 Or 5 atoms – 4 are 10 and 1 is 11! Av A r (B) = (10x1)+(11x4) =

20 The Mole - moles --> Mass m = n x Mr Calculate the mass of 2 moles of copper oxide (CuO) 0.5 moles of copper (II) sulphate (CuSO 4 ) 0.01 moles of calcium carbonate 5 moles of ammonium carbonate mass = moles x relative mass

21 The Mole - Mass --> Moles n = m/Mr Eg calculate the number of moles of water that would have a mass of 100g. How many moles of iron (II) chloride are in 50 g of iron (II) chloride? Calculate the number of moles needed to get1kg of calcium carbonate. How many moles of CuSO 4.5H 2 O would give you 0.1g of water?

22 The Mole & Mass --> Relative Mass Mr = m/n = 5.56 mol Eg Calculate the relative mass of a compound for which moles have a mass of 0,0056 g. What is the relative mass of a compound for which 0.01 mols has a mass of 0.18g Identify the element for which 0.05 moles has a mass of 0.16 g ? Mr (X)= m/n = /0.001 = 5.6 g/mol

23 The Mole - Reactions Sodium reacts with water to form hydrogen and sodium hydroxide according to the equation. Na + H 2 O  H 2 + NaOH If 46g of sodium are reacted with excess water what mass of hydrogen would be formed? 1.Balance the reaction 2.Work out moles of substance GIVEN. 3.Go through the equation to find out the number of moles reacting and being formed. (Molar ratio). 4.Work out quantity asked for.

24 Steps 1. Balance equation 2. Calculate moles given 3. Use Molar Ratio to find moles asked 4. Calculate quantity asked. The Mole - Reactions GIVEN ASKED 2. Moles given (m/m r ) 1 & 3 MOLAR RATIO 4. Moles asked (nxM r/v )  2H 2 + O 2  2H 2 O 4g of O 2 ? g H 2 O  n(O 2 ) = m/Mr  M:R O 2 :H 2 O 1:2.: n(H 2 O) = 2xn(O 2 )  m(H 2 O) = nxMr Amount GIVEN Amount ASKED

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26 The Mole - mass calculations C + O 2  CO 2 Carbon reacts with oxygen to form carbon dioxide as shown. If 0.12g of carbon are reacted with excess oxygen what mass of carbon dioxide would be formed? 1.Balance the reaction 2.Work out moles of reactant(mass given). 3.Go through the equation to find out the number of moles being formed 4.4.Work out quantity asked for.

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28 Mole examples - B & J p & p Na + Cl 2  NaCl Calculate the mass of salt formed if 2.3g of sodium is reacted with XS chlorine. 2.Zn + HCl  ZnCl 2 + H 2 What mass of HCl is needed to produce 100g of hydrogen? 3.KClO 3  KCl + O 2 What mass of oxygen is produced from 1kg of potassium chlorate? 4.Fe 2 O 3 + H 2  Fe + H 2 O What mass of iron is produced if 3g of rust (Fe 2 O 3 ) is reacted with XS(100g )of hydrogen?

29 Percentage Composition Analysis of a compound by mass makes it possible to work out the % mass of each element. eg Table salt: NaCl mass analysis: One mole of NaCl would have a mass of = 58.5g The % composition can be found using the formula: Mass element X x100 Total Mass Compound %Na = [ ….. / (…..) ]x100 = …………..% (by mass) %Cl = ( ….. / (…….) )x100 = …………% % Mass Element X =

30 Percentage Composition from mass. Eg2 Calculate the % of oxygen in water.

31 Empirical and Molecular Formula. A compound consists of carbon, hydrogen and oxygen only. The % by mass are Carbon 40.0% and 6.7% hydrogen. Calculate the empirical and molecular formula of the compound if M r = 60g·mol -1 %(O) = 100 – (40+6.7) = 53.3 CHOCHO In 100g:…….g……..g….…g n= m / Mr : … / … 6.7 / … / …… ……………….. …….…………. Simplest:…………. Empirical Formulae: ……. ( = …..) Molecular Formula: 2(CH 2 O) ……… (Mr = …. X 30)

32 Empirical and Molecular Formula. MOLECULAR FORMULA: CH 3 COOH or C 2 H 4 O 2 Actual formula M r : 2(12)+4(1)+2(16)=60g.mol -1 %C: ( 24 / 60 )x100 = 40.0% %H: ( 4 / 60 )x100 = 6.7% %O: ( 32 / 60 )x100 = 53.3% EMPIRICAL FORMULA: CH 2 O (Simplest whole number ratio)

33 Empirical and Molecular Formulae. Eg3. If a compound consisting of nitrogen and oxygen only - contains 30.4% by mass of nitrogen. What is the molecular formula of the compound? >>

34 Molar Volumes One mole of an ideal (ANY) gas occupies a volume of …………. 3 at ………………………… temperature and pressure. (STP) STP: T= ….ºC, ……K P =1 atmosphere (……...kPa) Fe 2 O 3 + 3H 2  2Fe + 3H 2 O What volume of hydrogen reacts with 50g of Fe 2 O 3 Fe 2 O 3 : H 2 … : ….. n(H 2 ) =…..n(Fe 2 O 3 ) = …………………… v(H 2 ) = …………………………… dm 3 n(Fe 2 O 3 ) = m/M r = ………………….= ………………mol moles = volume / molar volume ==> n = …. / …..

35 Molar Volumes One mole of an ideal (ANY) gas occupies a volume of 22,4dm 3 at standard temperature and pressure. (STP)

36 ASKEDGIVEN Mole Calculations REACTANTS  PRODUCTS MOLES MASS VOLUME MOLAR RATIO

37 Volume - Volume Calculations 1.Balance the equation 2.Calculate the moles of the substance given. 3.Work through the molar ratio to find out the moles of the substance asked. 4.Calculate the quantity asked for. (Volume V = n x M v ) M v = 22.4dm 3 At STP EG:H 2 + N 2 --> NH 3 If 3.00 dm 3 of nitrogen are reacted to produce ammonia, what volume of hydrogen will be required? (At STP) H 2 + N 2 --> NH 3

38 Reactions – Limiting reagent The reagent that runs out first and stops the reaction is known as the LIMITING REAGENT. If 46g of sodium are reacted with excess water what mass of hydrogen would be formed? Na + H 2 O  H 2 + 2NaOH 46g  2 moles XS Na will run out first Na is LIMITING REAGENT What is the minimum amount of water needed to react completely with 46g of sodium??

39 Limiting reagent example Ammonia gas is made by reacting ammonium chloride with calcium hydroxide according to: NH 4 Cl + Ca(OH) 2  NH 3 + CaCl 2 + H 2 O If 32.1 g of ammonium chloride reacts with 7.5 g calcium hydroxide in solution, Show by calculation; which is the limiting reagent and what mass of ammonia is produced.

40 Mass Volume Calculations 1.KClO 3  KCl + O 2 What volume of oxygen is produced by the decomposition of 1kg of potassium chlorate? 2.H 2 + N 2 --> NH 3 How much nitrogen (in dm 3 ) would be needed to produce 46dm 3 of ammonia? 3.S + O 2 --> SO 2 What volume of sulphur dioxide could be produced from 20.0dm 3 of oxygen? 4.Zn + 2HCl  ZnCl 2 + H 2 What mass of zinc is needed to produce 100dm 3 of hydrogen? 5.Fe 2 O 3 + H 2  Fe + H 2 O If 3.00kg of iron oxide is reacted with 0.256dm 3 of hydrogen, what mass of water would be produced?


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