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Chapter 1

Chapter 1. MATTER. 1.1 Atoms and Molecules. At the end of this topic, students should be able to:. Identify and describe proton, electron and neutron. Define proton no.,Z, nucleon no., A and isotope. Write isotope notation.

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Chapter 1

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  1. Chapter 1 MATTER 1.1 Atoms and Molecules

  2. At the end of this topic, students should be able to: • Identify and describe proton, electron and neutron. • Define proton no.,Z, nucleon no., A and isotope. Write isotope notation. • Define relative atomic mass and relative molecular mass based on the C-12 scale. • Sketch and explain the function of the following main components of a simple mass spectrum: • Analyze mass spectrum of an element. • Name cations, anions and salt according to the IUPAC.

  3. 1.1 Atoms and Molecules What is matter? • Matteris anything that has mass and occupies space. • All matters consist of tiny particles called atoms. • 3 states of matter are solid, liquid and gas. • For examples are metal, plastics, gas, etc.

  4. 1.1.1 Atoms

  5. Protons and neutrons are found in the nucleus of the atom, while electrons surround the nucleus to form ‘electron cloud’.

  6. The properties of these particles are summarized in the table.

  7. Nucleon • the particles that are found in the nucleus. • consist proton and neutron. • By proton we can identified the element.

  8. PROTON NUMBER AND NUCLEON NUMBER • Proton number (atomic number), Z. • The proton number = the number of protons in the nucleus of an atom. • Nucleon number (mass number), A. • The nucleon number = the total number of protons and neutrons in the nucleus of an atom.

  9. 1.1.2 Isotopic notation Nucleon number Atom or ion Proton number

  10. For neutral atom, number of protons = number of electrons. • For +ve ion, number of electron is less than number of protons. • For –ve ion, number of electrons is more than number of protons.

  11. Example: Number of protons = proton number = 13 Number of neutrons = nucleon number – proton number = 27 - 13 = 14 Number of electron = proton number – charge carried by species = 13 – ( +3) = 10

  12. Exercise: • Determine the number of proton, neutron and electron in the following species. a) b)

  13. Exercise: The atomic number of lead (Pb) is 82 and the mass number is 207. • Write out the atomic notation for lead. • Give the number of protons, neutrons and electrons.

  14. 1.1.3 ISOTOPES • Two or more atoms of the same element having same proton number but different nucleon number.

  15. Eg : 1) hydrogen isotope • 2) carbon isotope 1214 66 1 1 2 1 3 1 H H H C C ;

  16. Isotopes of an element have the same: • number of protons (proton number) • charge of nucleus of the atoms (ionization energy; electron affinity; size of the atom; electronegativity are the same) • number of electrons in a neutral atom • electronic configuration (the number of valence electrons) • chemical properties

  17. Isotopes of an element have different: • number of neutrons (nucleon number) in the nucleus of the atoms • relative isotopic mass • physical properties (e.g boiling point / melting point, density, effusion rate,…)

  18. 1.1.4 Molecule • A molecule consists of a small number of atoms joined together by covalent bond. • Diatomic molecule: contains two atoms (example: H2, Cl2, HCl, CO) • Polyatomic molecule: contains more than two atoms (example: H2O, NH3)

  19. 1.1.5 Ion • An ion is a charged species formed from a neutral atom or molecule when electrons are gained or lost as the result of a chemical reaction. • Cation: a positively charged ion (number e < number p) (example: Mg2+, Al3+) • Anion: a negatively charged ion (number e > number p) (example: Cl, OH)

  20. Monatomicion: ion contains only one nucleus (example: Fe3+, S2) • Polyatomic ion: ion contains more than one nucleus (example: H3O+, CN)

  21. 1.1.6 RELATIVE MASS • Isotopes carbon -12 as a reference or standard for comparing the masses of other atoms. RELATIVE ATOMIC MASS, Ar • Average mass of one atom of the element relative to 1/12 times the mass of one atom of carbon-12.

  22. Ar = average mass of one atom of the element 1/12 x mass of one atom of C-12 • The mass of C-12 is defined exactly 12.00 a.m.u. • 1 atomic mass unit : a mass unit equal to 1/12 the mass of a C-12 atom

  23. Example: Oxygen consists of three isotopes 16O ;17O and 18O in the ratio of 99.76 : 0.04 : 0.20. Calculate the relative atomic mass (to 4 decimal point) of oxygen. ANSWER : Relative Atomic Mass = (16 x 99.76) + (17 x 0.04) + (18 x 0.20) (99.76 + 0.04 + 0.20) = 16.0044

  24. Exercise: Determine the Ar for Ne consists of three isotopes 20Ne ; 21Ne and 22Ne in the ratio of 90.92 ; 0.26 and 8.82. Answer: 20.18

  25. RELATIVE MOLECULAR MASS, Mr • The mass of one molecule of the substance relative to 1/12 times the mass of one atom of carbon-12. • = sum of the relative atomic masses of all the atoms shown in the molecular formula. Mr = average mass of one molecule of the substance 1/12 x mass of one atom of C-12

  26. Example: The relative molecular mass of carbon dioxide, CO2. Answer: Mr CO2 = Ar C + 2 Ar O = 12 + 2(16) = 44

  27. Exercise: The relative molecular mass of ethanol, C2H5OH Answer: 46.00

  28. 1.1.7 Mass Spectrometer Mass spectrometer is used to determine: • Relative atomic mass of an element • Relative molecular mass of a compound • Types of isotopes that are found in the naturally occurring element including the abundance of the isotopes and its relative isotopic mass. • Recognize the structure of the compound in an unknown sample

  29. A B 37Cl+ Vacuum 35Cl+ Vaporisation chamber Acceleration chamber (Electric field) Magnetic field Ionization chamber Ion detector There are five main stages:

  30. (a) Vaporisation Chamber ~Sample of the element is vaporised into gaseous atom (b) Ionisation Chamber ~A gaseous sample (atom or molecule) is bombarded by a stream of high-energy electrons that are emitted from a hot filament. Collisions between the electrons and the gaseous atom (or molecule) produce positive ions by dislodging an electron from each atom or molecule. M sample + efilament M+ sample + e sample + efilament M M+

  31. (c) Acceleration Chamber (Electric field) • The positive ions are accelerated by an electric field towards the two oppositely charged plates. The electric field is produced by a high voltage between the two plates. The emerging ions are of high and constant velocity.

  32. Magnetic field • The positive ions are separated and deflected into a circular path by a magnet according to its mass/charge (m/e) ratio. • Positive ions with small m/eratio are deflected most and appear near A. Ions with large m/e ratio are deflected least and appear near B (Figure 1.1.2).

  33. (e) Ion detector • The numbers of ions and types of isotopes are recorded as a mass spectrum. • In practice, the ion detector is kept in a fixed position. The magnetic field is varied so that the positive ions of different masses arrive at the detector at different times.

  34. Mass spectrum • the horizontal axis – • the m/e ratio • nucleon number • isotopic mass • relative atomic mass of the ions entering the detector. • The vertical axis - the abundance or detector current or relative abundance or ion intensity or percentage abundance of the ions.

  35. Information from a mass spectrum of an element • the isotopes which are present in the element • the relative isotopic mass of each isotope • the abundance of each isotope • Thus, the relative atomic mass of the element can be determined

  36. Relative abundance 63 9.1 8.1 0 24 25 26 m/e The mass spectrum of magnesium shows that naturally occurring magnesium consists of three isotopes: 24Mg, 25Mg and 26Mg.

  37. The height of each line is proportional to the abundance of each isotope. • In this example, 24Mg is the most abundance of the three isotopes. Ar of Mg: = 24.33

  38. Average atomic mass = where • Q = the abundance of an isotope of the element = the percentage of the isotope found in the naturally occurring element • m = the relative isotopic mass of the element

  39. Notes : • The height of each peak measures the relative abundance of the ion which gives rise to that peak. • The total number of peaks in the mass spectrum of an element shows the types of naturally occurring isotopes. 3) The ratio of mass/charge for each species is found from the value of the accelerating voltage associated with a particular peak. Many ions have a charge of +1 elementary charge unit, and the ratio m/e is numerically equal to m, the mass of the ion. (1 elementary charge unit = 1.60x10 19 C) 4) The ion with the highest value of m/e is the molecular ion, and its mass gives the molecular mass of the compound

  40. 1.1.8 IUPAC Nomenclature for Cations, Anions and Salts • Cations • A positive ion is formed by the loss of a certain number of electrons from the given atom • For example: • Mg 2e- + Mg2+ (2e- were lost by Mg atom) • K 1e- + K+(1e- were lost by K atom)

  41. Anions • A negative ion is formed by gain of one or more electrons by the given atom • For example: • S + 2e- S2-(2e- were gain by sulphur atom) • Cl + 1e- Cl-(1e- were gain by chlorine atom)

  42. Salts - Combination of cation and anion to form ionic compound • For example: • Na+ + Cl- NaCl • Mg2+ + 2Br- MgBr2

  43. a)Naming monoatomic ions • Monoatomic ions:ions consisting of single atomic nucleus • Cation: named by just adding ‘ion’ Ex: K+ : potassium ion Ca2+ : calcium ion • Anion : named by adding suffix ‘ide’ Ex: Cl- : chloride ion S2- : sulphide ion O2- : Oxide ion

  44. b) Naming ionic compounds When naming an ionic compound made up of two element (binary): • the name of the cation (metal ion) appear first, followed by the name of the anion(non-metal ion)

  45. System for naming cation of transition metals • Stock system/systematic system: - a Roman numeral that match the ionic charge is placed in bracket immediately after the element name of metal - for example: Fe 2+ iron(II) ion

  46. Common nomenclature system - suffix-ous indicates the lower ionic charge and the suffix-ic indicates the higher ionic charges - for example: Fe2+ ferrous ion Fe3+ ferric ion

  47. Systematic (stock) and common names for ion and copper ions

  48. Rules for writing for an ionic compound • The positive ion is always written first. • The ratio of positive ions to negative ions must be such that the total number of +ve charge equal the total number of -ve charge; the formula unit must be electrically neutral • The smallest set of subscripts that give electrical neutrality is always chosen. We always write empirical formulas for ionic compounds

  49. c)Naming polyatomic ions • Polyatomic ions are composed of 2 or more atoms bonded together • Most polyatomic ion consist to non-metal such as P, S, C, or N bonded to oxygen atoms • -ate: most common polyatomic ion • -ite : used for names of related ion that have one less oxygen atom

  50. Chapter 1 MATTER 1.2 Mole Concept

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