Amount of Substance. http://www.avogadro.co.uk/definitions/elemcompmix.htm

of 14 Amount of Substance Key terms in this chapter are: Element Compound Mixture Atom Molecule Ion Relative Atomic Mass Avogadro constant Mole Isotope Relative Isotopic Mass Relative Molecular Mass Relative Formula Mass Simple molecular structure Giant ionic structure Molecular formula Empirical formula Structural formula Structural Isomer Before starting, you must check up on elements, compounds and mixtures. Write some brief notes to outline these ideas. The following link will help you. http://www.avogadro.co.uk/definitions/elemcompmix.htm Relative Atomic Mass and the Avogadro constant In this section you are going to learn: The meaning of the term relative atomic mass for elements containing atoms. The meaning of the term relative isotopic mass for elements with isotopes. That the relative atomic mass of an element when taken in grams is the mass of 1 mole of that element. About the Avogadro constant, its numerical value, the term 1 mol, and how these relate to elements. How to do simple calculations involving the mole. We start the topic Amount of Substance with a question Chemists don't deal with the masses of atoms and molecules measured in grams because they would be dealing with extremely small numbers. Instead, the masses of atoms and molecules are compared with that of the isotope carbon-12. 12 C is assigned a relative atomic mass of 12. As an example, consider an atom of magnesium. Two atoms of 12 C are required to be equal in mass to one atom of Mg therefore the relative atomic mass of Mg is 24. So, using the scales below, one atom of magnesium is balanced by two atoms of carbon-12. You should note that relative masses have no units because in their calculation the units have cancelled.

of 14 Q1. Use a chemistry textbook or Periodic Table to look up the relative atomic mass of silver. Now draw in the correct number of carbon atoms needed to balance one atom of Ag. Study the information in the table below. It refers to the atoms of elements. Element Ca Na Pb Fe Cl H He Ag Relative Atomic Mass, Ar Mass of 1 mol 40 23 207 56 35.5 1 4 108 40 g 23 g 207 g 56 g 35.5 g 1 g 4 g 108 g In each of the above masses of elements, there is the same number of atoms. That is, the relative atomic mass, Ar, of an element expressed in grams always contains the same number of atoms. This number can be determined experimentally. It is 6.023 x 10 23. The number is a constant, called the Avogadro constant, L (or N A ). The mass of this amount of atoms is called one mole (1 mol). Amedeo Avogadro was born in Turin, Italy, on 9th August 1776. It was in 1811 when he published his famous hypothesis, now called Avogadro's law. Avogadro died on the 9th July 1856. The postage stamp pictured was to remember Avogadro 100 years after his death. Avogadro's law is considered in a later part of this course. For an element with isotopes, the relative mass of the atoms of each of isotope of that element can be referred to as the relative isotopic mass. If an element has no isotopes then the term relative atomic mass is normally used. We will return to this idea in the next section. Note that chemists often use the term relative atomic mass when relative isotopic mass would be strictly more correct.

of 14 Relative Molecular Mass - Elements and compounds consisting of molecules In this section you are going to learn: That a molecule is a group of two or more covalently bonded atoms. That some elements and some compounds consist of molecules. That the relative molecular mass of an element or compound when taken in grams is the mass of 1 mole of that substance. About the units of the Avogadro constant. How to do simple calculations involving the mole. Study the information in the table below. It refers to the molecules of elements and compounds. Element or compound Relative Molecular Mass, Mr Mass of 1 mol CO 2 SO 2 N 2 H 2 Cl 2 CH 4 H 2 O NH 3 44 64 28 2 71 16 18 44 g 64 g 28 g 2 g 71 g 16 g 18 g 17 g A molecule is a group of two or more atoms that are bonded together by covalent bonding. A molecule is an entity in itself. A dot-cross diagram representing methane, CH 4, is represented below. The relative molecular mass, Mr, of an element or compound expressed in grams always contains the same number of molecules. Again, this number is the Avogadro Constant. The Avogadro constant has units Since the Avogadro Constant is the number of particles (atoms, molecules, ions) in 1 mole it has units they are mol -1 (particles mol -1, but the particles bit is omitted). The Avogadro Constant is 6.023 x 10 23 mol -1. Relative molecular mass, Mr, can be calculated by adding together the relative atomic masses of the constituent atoms, e.g. Mr (CO 2 ) = 12 + 16 + 16 = 44.

of 14 Q2. The relative molecular mass of water is 18. Complete the diagram below by balancing the number of water molecules against carbon atoms. Relative Formula Mass - Compounds consisting of ions (ionic compounds) In this section you are going to learn: That Sodium Chloride, NaCl, is an example of an ionic compound. It has a giant ionic lattice structure. The terms 'relative formula mass' and 'formula unit' as they apply to ionic compounds. What is meant by 1 mol of an ionic compound and how this relates to the amounts in moles of the ions it contains. How to do simple calculations involving the mole. Giant ionic lattice structure (for NaCl) means that in the crystal structure of sodium chloride each Na + ion is surrounded by six Cl - ions and each Cl - ion by six Na + ions. This arrangement of ions can in theory go on and on indefinitely. This is why the structure is referred to as 'giant'. It is represented below. You can read more about this at: http://www.avogadro.co.uk/structure/chemstruc/ionic/g-ionic.htm Also, the structure is held together by electrostatic attractions between the oppositely charged ions. These are called ionic bonds. Sodium chloride therefore is not molecular, it is ionic. Also, the formula NaCl is an empirical formula (see later), and is sometimes called a formula unit. It is best to use the term relative formula mass (rather than relative molecular mass) for ionic compounds. For the formula unit Na + Cl - the relative formula mass is 23 +

of 14 35.5 = 58.5. One mole of Na + Cl - is 58.5 g. This mass contains 1 mol (23 g) of Na + ions and 1 mol (35.5 g) of Cl - ions, that is, two moles of ions in total. Note that sometimes the term relative molecular mass is used to refer to ionic compounds. Q3. Work out the relative formula mass of FeSO 4,7H 2 O(s). Q4. What is the mass of 1 mole of iron(ii) sulphate(vi)-7-water? Q5. You have 2 moles of FeSO 4,7H 2 O. How many moles of each type of ion and molecule are present in this amount? Mole Calculations In this section you are going to learn: More about simple calculations involving the mole concept. Think about the above ideas and you should be able to calculate the amount in moles of a given mass of a substance, e.g. 35.5 g of Cl atoms is 1 mol. 35.5 g of Cl 2 molecules is ½ mole of Cl 2 molecules. If you are given the amount in moles, then multiplying this by the mass of one mole gives the amount in grams, e.g. 1½ mole of SO 2 is 1½ x 64 = 96 g. Q6. Calculate the mass in grams of ¾ mole of sulphuric(vi) acid, H 2 SO 4. Q7. You have 145 g of propanone, CH 3 COCH 3, what amount of this is there in moles?

of 14 Balanced Chemical Equations In this section you are going to learn: That a balanced chemical equation gives information about the amounts of reactants that react and of products that form. How to do simple mole calculations involving chemical equations. Balanced chemical equations give information about the molar amounts of substances that react and the molar amounts of products that form. For example, 2Mg(s) + O 2 (g) 2MgO(s) The above equation tells us that 2 moles of magnesium atoms react with 1 mole of oxygen molecules to form 2 moles of magnesium oxide, an ionic compound. This means that, for example, 12 g (½ mole) Mg will require 8 g (¼ mole) of O 2 to react completely to form 20 g (½ mole) MgO. Q8. If 12.80 g of magnesium metal is completely burned in excess oxygen, what mass of magnesium oxide is obtained? Q9. Some granulated zinc metal was added to an excess of a solution of hydrochloric acid. The zinc reacted completely and 4 g of hydrogen gas was formed. Calculate the mass of zinc in grams added to the acid.

of 14 Chemical Formulae The meaning of empirical formula, molecular formula and structural formula and how these are related. About structural isomers and how to draw these for simple alkanes. This section refers to compounds that are molecular. With a little knowledge of organic chemistry you would know that C 4 H 10 is the formula for two alkanes. It tells us that in each of these there are four atoms of carbon and ten atoms of hydrogen bonded together. This type of formula is called a molecular formula. It gives the actual number of each type of atom in the molecule but does not tell you the arrangement in which they are bonded. The simplest whole-number ratio of C:H is 2:5. This gives the formula C 2 H 5, and it is called the empirical formula. (Remember that NaCl, for example, is an empirical formula because of the nature of its giant ionic lattice structure.) Q10. The molecular formula of glucose is C 6 H 12 O 6. What is its empirical formula? Q11. The molecular formula of sucrose is C 12 H 22 O 11. What is its empirical formula? So, the empirical formula of a covalent compound is the simplest whole-number ratio of atoms in its molecules. Or, you could say the simplest whole-number ratio of moles of atoms in 1 mole of the compound. If you know the empirical formula of a compound and also its relative molecular mass, then the molecular formula can be found. This is illustrated below. Q12. A hydrocarbon consists of 82.8% carbon and 17.2% hydrogen. Its relative molecular mass is 58. Calculate the empirical formula and the molecular formula of the compound.

of 14 A structural formula shows the arrangement in which the atoms are bonded. For C 4 H 10 these are shown below and are referred to as structural isomers: Some additional questions for you to try... 1. An atom of titanium has four times the mass of an atom of 12 C. What is its relative atomic mass? 2. An atom of helium has one-third the mass of an atom of 12 C. What is its relative atomic mass? 3. A molecule of water has 1½ times the mass of an atom of 12 C. What is its relative molecular mass? 4. What mass of 12 C consists of 12.04 x 10 23 atoms? 5. The relative atomic mass of lead is 207. What is the molar mass (mass of 1 mole) of Pb? 6. What is the mass of 2 moles of molecules of ethanol, CH 3 CH 2 OH (Ar (H) = 1; Ar (C) = 12; Ar (O) = 16)? 7. What mass of sulphur contains the same number of atoms as there are present in 1.5 moles of magnesium (Ar (Mg) = 24; Ar (S) = 32)?

of 14 8. What mass of ammonia (NH 3 ) contains the same number of molecules as there are atoms present in 2 moles of sodium (Ar (H) = 1; Ar (N) = 14; Ar (Na) = 23)? 9. How many moles of oxygen are there in 128 g of O 2 (Ar (O) = 16)? 10. How many moles of sodium sulphate(vi), Na 2 SO 4, are contained in 71 g (Ar (O) = 16; Ar (Na) = 23; Ar (S) = 32)? 11. How many moles of Al 3+ are there in 2 moles Al 2 (SO 4 ) 3? 12. In the following reaction, what mass (in grams) of SO 2 (g) is formed when ½ mole of S reacts with oxygen (Ar (O) = 16; Ar (S) = 32;)? S(s) + O 2 (g) SO 2 (g) 13. In an experiment to find the formula of copper iodide, 0.17 g of Cu reacts with 0.34 g of iodine. What is the formula of copper iodide (Ar (Cu) = 63.55; Ar (I) = 126.90)?

of 14 14. In an experiment to find the formula of a lead oxide, 11.95 g of the oxide, on reduction (removal of oxygen), gave 10.35 g of Pb. What is the formula of the oxide (Ar (0) = 16; Ar (Pb) = 207)? 15. What is the percentage mass of S in SO 2 (Ar (O) = 16; Ar (S) = 32)? 16. What is the percentage mass of water of crystallisation (H 2 O) in FeSO 4.7H 2 O(s) (Ar (H) = 1; Ar (O) = 16; Ar (S) = 32; Ar (Fe) = 56)? &KHPLFDO,GHDVH[SODLQVKRZUHODWLYHPDVVHVRIDWRPVDQGPROHFXOHVFDQEHPHDVXUHGXVLQJD PDVVVSHFWURPHWHU,WDOVRPHQWLRQVLWVXVHLQWKHILQGLQJRXWDERXWWKHVWUXFWXUHRIPROHFXOHV The Mass Spectrometer In this section you are going to learn: How the mass spectrometer works. How relative isotopic mass and relative atomic mass are related, and how to calculate a relative atomic mass from mass spectral data. How the mass spectrometer is used in the analysis of molecular structure. The existence of isotopes of elements was discovered using an instrument called a mass spectrometer. The mass spectrometer was invented by the English physicist Francis William Aston (1877-1945) when he was working in Cambridge with J. J. Thomson. Aston eventually discovered many of the naturally occurring isotopes of non-radioactive elements. He was awarded the Nobel Prize for Chemistry in 1922.

of 14 Time-in-flight Mass Spectrometer,QWKLVVHFWLRQVRPHLGHDVDQGGLDJUDPVDUHDGDSWHGIURP KWWSZZZNRUHFRXNWXWRULDOKWP The following description of a time-of-flight mass spectrometer is very simplified. We assume all ions produced leave the source at the same time, with the same kinetic energy. A time-of-flight mass spectrometer identifies sample atoms (or molecules) by measuring their flight time. This diagram above shows the working principle of a linear time-of-flight mass spectrometer. It is necessary to put a charge on the atoms (or molecules) by bombarding them with electrons emitted from a filament. Each atom (or molecule) loses one electron to form an ion with a single positive charge. M(g) + e - M + (g) + 2e - The sample ions (M + ) are then accelerated by an electric field to give each of them the same kinetic energy. They then enter into a part of the mass spectrometer, called the flight path, which has no electric field. All of the ions travel the same distance through this part of the instrument to reach the ion detector. The speed of an ion is dependent upon its mass, with heavy ions having a lower velocity than light ones, and so the time it takes for an ion to reach the detector is related to its mass. The mass spectrometer measures the relative abundance (relative number) and relative mass of each type of ion reaching the detector. This data is represented as a mass spectrum. The mass of an atom or molecules is expressed relative to that of an atom of the isotope carbon-12 ( 12 C = 12).

of 14 Where an element has two or more isotopes, the mass spectrometer measures the relative isotopic masses and their relative numbers (relative abundances) and plots a mass spectrum, as shown below. The relative isotopic mass of a chlorine atom is either 35 or 75. However, the relative atomic mass of chlorine is given as 35.5 in a chemical data book. This is because, for an element with isotopes, the relative atomic mass of an element takes into account the relative abundance and relative mass of each isotope present. In a typical sample of naturally occurring chlorine, 100 atoms will consist of 75 of relative mass 35 and 25 of relative mass 37. Using the mass spectrum on the left, the relative atomic mass of Cl is calculated as follows: Ar = (75/100 x 35) + (25/100 x 37) = 35.5 In the spectrum on the right, the most abundant isotope is assigned a relative abundance of 100% and each of the other isotopes is assigned a relative abundance relative to that. Q13. Use the mass spectrum on the right above to calculate the relative atomic mass, Ar, of chlorine. Q14. The mass spectrum of antimony shows it has two naturally occurring isotopes at relative mass 121 (relative abundance 57.3%) and relative mass 123 (relative abundance 42.7%) respectively. Draw a diagram to represent this mass spectrum and calculate the relative atomic mass, Ar, of antimony.

of 14 The mass spectrometer in the analysis of molecular structure When the molecules of a substance enter the mass spectrometer to become ionised by loss of an electron, the 'molecular ion' (M + ) is formed. Some of these pass through the instrument and remain intact, to produce a peak in the mass spectrum at the relative molecular mass. Others are likely to fragment (break up) in various ways to produce smaller positively charge ions, which produce corresponding peaks in the mass spectrum. These numerous peaks can reveal how the atoms of the molecule were bonded together, that is its molecular structure. There is more about these ideas in a later module. Some additional questions about mass spectrometry This chapter provides some additional questions to help you assess your learning. When you have answered these and have checked your answers against the correct ones, refer to the table that follows to find out what you may need to do to further your progress. 1. An atom of titanium has four times the mass of an atom of 12 C. What is its relative atomic mass? 2. An atom of helium has one-third the mass of an atom of 12 C. What is its relative atomic mass? 3. Using a mass spectrometer, it is possible to determine the number of A protons in an atom. B energy levels in an atom. C atoms in 1 mole of an element. D isotopes of an element. E neutrons in an atom. 4. Iron has four isotopes with relative numbers of atoms as follows: 54 Fe (5.8%), 56 Fe (91.7%), 57 Fe (2.2%), 58 Fe (0.3%). Calculate the relative atomic mass of iron.

of 14 5. The mass spectrum for naturally-occurring magnesium is shown below. Singly charged Mg + ions have produced the peaks. What is the relative atomic mass of magnesium? A 24.2 B 24.3 C 24.4 D 24.7 E 24.8 6. Which one of the following particles has seven protons, eight neutrons and nine electrons? A 7 8 N - B 7 14 N 2- C 7 15 N - D 8 15 N 2- E 7 15 N 2-