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Atomic Theory

Atomic Theory . Unit 7. The Atom. An atom is the smallest part of any element that can take part in a chemical reaction. An atom is the smallest particle of an element which still retains the properties of that element. Atoms are extremely small. Structure of the Atom.

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Atomic Theory

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  1. Atomic Theory Unit 7

  2. The Atom • An atom is the smallest part of any element that can take part in a chemical reaction. • An atom is the smallest particle of an element which still retains the properties of that element. • Atoms are extremely small.

  3. Structure of the Atom

  4. Elements and Compounds • An element is a substance that cannot be split into a simpler substance by chemical means. • A compound is a substance which is made up of two or more different elements combined together chemically. • Most elements from the periodic table are found in nature as compounds very few are found in elemental form.

  5. Molecules • A molecule is a group of atoms joined together. • It is the smallest part of an element or compound that can exist independently. • There are two types of molecules: • Molecules of elements: (All the same atoms in the molecule) e.g. Cl2,S8. • Molecules of compounds: (Different atoms in the molecule) e.g. H2O, HCl, H2SO4.

  6. Diatomic Elements • Elements that end in –ine and –gen exist as diatomic molecules in their natural form.

  7. Atoms, Molecules, Elements & Compounds Atoms (Only 1) of Elements of (all the same) Molecules (2 or more) of Compounds (different)

  8. Summary

  9. The Periodic Table • The Periodic table is the chemist’s compass to a vast array of compounds. • In the modern periodic table the elements are listed in order of increasing atomic number, and arranged in order of the numbers of electrons in their outer main energy level. • These are referred to as the group number and the period number. • The periodic table is divided into two main sections metals and non – metals. A memory aid to where this division occurs is : Being Silly Assures Teachers Attention.

  10. P.T.E The Periodic Table Unit 7, 10

  11. P.T.E Some of the Named Groups The Named Groups Unit 7, 11

  12. Group Number • The group number of an element is equal to the numbers of electrons in the outer main energy level (shell) of an atom of that element. • The named groups are: • Group one – Alkali metal. • Group two – Alkaline earth metals. • Group seven – Halogens. • Group eight – Noble gases.

  13. Period • The period number of an element is the number of main energy levels (shells) in an atom.

  14. Solids, Liquids and Gases Solids, Liquids and Gases Solids Gases Liquids Unit 7, 14

  15. Ions • Ions are atoms or groups of atoms whichhave lost or gained electrons and hence have a charge. • A positive ion (cation) is formed when an atom(s) loses electrons, they are denoted by the following notation Xn+. • A negative ion (anion) is formed when an atom(s) gains electrons, they are denoted by the following notation Xn-. • Simple ions (ions of 1 atom only) can be worked out from their valency (and from the periodic table). • The valencyof an element is equal to the number of electrons which an atom must lose or gain in order to attain noble gas configuration.

  16. P.T.E Simple Ions Simple Ions Unit 7, 16

  17. Transition Metal Ions

  18. Polyatomic Ions

  19. Making Formulas for Compounds • Get the correct ions from: • Periodic Table of Elements. • Transition metals table. • Polyatomic ions table. • Put the positive ion first. • If the charges are different cross over the magnitude (number part) of the charges. • If you require more than one of a polyatomic ion make sure to put it in brackets before adding the subscript. • There should be no charges on a compound.

  20. P.T.E Unit 7, 20

  21. P.T.E Unit 7, 21

  22. P.T.E Unit 7, 22

  23. P.T.E N.B. If you require more than one of a polyatomic ion you must put it in brackets. Unit 7, 23

  24. P.T.E N.B. When we require only one of an item there is no need to fill in a one, it is assumed Unit 7, 24

  25. P.T.E Try Some! Unit 7, 25

  26. Naming Compounds Summary

  27. Atomic Number • The modern periodic table has the elements arranged in order of increasing atomic number. (Smaller of the two) The atomic number of an element in the Periodic Table tells us the number of protons in the nucleus. • i.e. sodium is the eleventh element in the periodic table and therefore it has 11 protons in its nucleus.

  28. Atomic Mass Number The atomic mass number of an element is the sum of the numbers of protons and neutrons in the nucleus of an atom. • For example, sodium has 11 protons and 12 neutrons in the nucleus. Therefore its mass number is 23 or we say that the mass of the atom is 23 atomic mass units. • A special unit called the atomic mass unit is used as the masses of the atoms are so small.

  29. Nuclear Formula

  30. Nuclear Formula • It is called a nuclear formula as it only gives information about the nucleus of the atom or ion, you have to deduce the information for electrons. • Remember: • All atoms are neutral. • Positive ions have lost electrons. • Negative ions have gained electrons.

  31. Unit 7, 31

  32. Isotopes • In 1919 an English chemist called Francis William Aston built an instrument called a mass spectrometer to measure the masses of atoms. • He started work with a sample of neon gas and discovered something very unusual. He found that that neon gas consisted of two varieties of neon atom. • One type of neon atom had a mass number of 20 and the other type had a mass number of 22. • He concluded that neon gas consisted of atoms of neon that differed in the number of neutrons in the nucleus.

  33. Isotopes of Chlorine & Carbon

  34. Isotopes of Lithium

  35. Isotopes of Hydrogen

  36. % of Each Isotope Present • For his work on the discovery of the isotopes Aston received the 1922 Nobel Prize in chemistry. • Not only did Aston detect the presence of isotopes, but he also determined the percentages of each of the isotopes present. • For example in his study of chlorine gas he found that there were approximately three times as many chlorine – 35 atoms as there were chlorine – 37 atoms. He was then able to work out the average mass of an atom of chlorine. • The periodic table contains this figure for the average mass of a atom of the element concerned. This is why the values given in the periodic table are all decimals. • His method of calculation is shown in the next example.

  37. Unit 7, 37

  38. Relative Atomic Mass • If you locate chlorine in the periodic table you will find that 35.45 is the number given under the symbol for chlorine. • This ‘average mass’ of an atom is measured relative to the mass of the carbon – 12 isotope. • For this reason it is called the relative atomic mass. The symbol for relative atomic mass is Ar. • Since the relative atomic mass is the ratio of two masses, it has no units. Therefore, we say that the average mass of an atom of chlorine is 35.5 a.m.u but its relative atomic mass is 35.5.

  39. Unit 7, 39

  40. The response of the ion detector (intensity of lines on photographic plate) is converted to a scale of relative numbers of atoms. (i.e. % of each isotope present) Unit 7, 40

  41. Mass spectrum for mercury. The percent natural abundances of the mercury isotopes are:

  42. Average Mass of Mercury

  43. Calculating Molecular Mass

  44. Unit 7, 44

  45. The Arrangement of Electrons • Up to this we said that electrons orbit the nucleus in shells (Bohr’s theory). • Each shell contains electrons with a certain amount of energy and electrons normally occupy the lowest available energy level. • The ground state of an atom refers to its state when all of its electrons are in their lowest available energy levels.

  46. No. of Electrons in each Shell • In this situation the first shell n = 1 could hold 2n2 = 2 electrons. • In this situation the second shell n = 2 could hold 2n2 = 8 electrons. • In this situation the third shell n = 3 could hold 2n2 = 18 electrons. • In this situation the fourth shell n = 4 could hold 2n2 = 32 electrons. • This works well to explain some topics in chemistry, but a more detailed arrangement is sometimes necessary.

  47. Bohr Model of the Atom • In a Bohr atom, the electron is a particle that travels in specific, fixed orbits, but never in the space between orbits. • This arrangement expresses energy quantization, and accounts for the atomic emission spectra.

  48. A Ladder for Distance, Bohr Model for Energy • Bohr orbits are like steps in a ladder. It is possible to be on one step or another, but it is impossible to be between steps. • Unlike the ladder Bohr orbitals do not have equal spacing between the orbits.

  49. Energy Absorption and Emission • When a hydrogen atom absorbs energy, an electron is excited to a higher energy level. • The electron is then in an unstable and temporary level. • The electron falls back to the lower energy level and emits a photon of light (or some other form of radiation).

  50. Evidence for the Existence of Energy Levels • When atoms are excited (i.e. given energy) by heating them or subjecting them to electrical discharge, they usually emit light or some other form of radiation. • A spectrum of light from a bulb would produce a continuous spectrum from red through to the various colours to violet. • The spectrum from excited atoms are not continuous but consist of a number of distinct lines each corresponding to a definite frequency of light. (Each colour of light has a different frequency)

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