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Chemistry UNIT 3

Chemistry UNIT 3

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Chemistry UNIT 3

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  1. ChemistryUNIT 3 PERIODICITY

  2. 3.1 The periodic table

  3. 3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number

  4. Metalloids Metalloids are the elements found along the stair-step line that distinguishes metals from non-metals. This line is drawn from between Boron and Aluminum to the border between Polonium and Astatine. The only exception to this is Aluminum, which is classified under "Other Metals".

  5. 3.1.2 Distinguish between the term group and period • Group • These are the numbers represented by Roman Numerals at the top of the periodic table. • This number tells us the number of valence electrons in an atom. • Valence electrons are important because they determine the chemical reactivity of elements. • Period: Represent the number of energy levels /Shells

  6. 3.1.1 Apply the relationship between the electron arrangement of elements and their position in the periodic table up to Z = 20

  7. The electron configuration of the first 20 elements 3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table.

  8. Do the following questions: 1-3 p. 75 students text book • Question from work book. Exercise 3.1 P. 97-99

  9. 3.2 Physical Properties • Elements in the same group therefore have different physical properties • Physical properties include: • Effective nuclear charge • Atomic radius • Ionic radius • Ionization energy • Electronegativity • Melting point

  10. The trend in the physical and chemical properties are governed by the effective nuclear charge. • The nuclear charge of the atom is given by the atomic number. As you go across the periodic table the atomic number increases by one, as one proton is added to the nucleus. • The outer electron (that determine many of the physical and the chemical properties) do not experience the full attraction of this charge (protons) as they are shielded from the nucleus and repelled by the inner electrons.

  11. The presence of the inner electrons reduces the attraction of the nucleus for the outer electrons. • The effective charge experienced by the outer electron is less than th efull nuclear charge. • As you go across a period, one proton is added to the nucleus and one electron is added to the outer electron shell. The effective charge increases with the nuclear charge as there is no change in the number of the inner electrons

  12. Chemical property • Group I Metals(Alkali Metals) • These metals react with water to form alkaline solutions • They have low melting points • Low boiling points • Soft • Low density • Very reactive

  13. 3.2 Physical properties: Trends in Melting and boiling points        You will see that both the melting points and boiling points fall as you go down the Group.

  14. Group II Metals The alkaline earth elements are metallic elements found in the second group of the periodic table. All alkaline earth elements have an oxidation number of +2, making them very reactive. Because of their reactivity, the alkaline metals are not found free in nature. The Alkaline Earth Metals are: Beryllium Magnesium Calcium Strontium Barium Radium

  15. Extracted from oxides in the earth’s crust • Less reactive than the alkali metals

  16. Metalloids have properties of both metals and non-metals. Some of the metalloids, such as silicon and germanium, are semi-conductors. This means that they can carry an electrical charge under special conditions. This property makes metalloids useful in computers and calculators The Metalloids are: Boron Silicon Germanium Arsenic Antimony Tellurium Polonium

  17. Properties of Metals - They have high melting and boiling points • Properties of Non-metals • Have low densities • Brittle

  18. Group VII THE HALOGENS The halogens are five non-metallic elements found in group 7 of the periodic table. The term "halogen" means "salt-former" and compounds containing halogens are called "salts". The Halogens are: Fluorine Chlorine Bromine Iodine Astatine

  19. The halogens exist, at room temperature, in all three states of matter: • Solid- Iodine, Astatine • Liquid- Bromine • Gas- Fluorine, Chlorine

  20. The Halogens are: Fluorine Chlorine Iodine Bromine Astatine

  21. Group VIII The Noble Gases • The six noble gases are found in group VIII (8) of the periodic table. • These elements are very stable and do not react with anything else. • They are sometimes referred to as inert gases • Helium • Neon • Argon • Krypton • Xenon • Radon • These gases have low melting points and boiling points.

  22. PERIODS • Period Number indicates the number of occupied shells. • Elements in the same period share a gradual change in their physical and chemical properties.

  23. Trends in the Periodic Table • The trends in the periodic table are - atomic radius - first ionisation energy - electronegativity - melting and boiling points - density.

  24. 3.2.2 Trends in Atomic Radius • ATOMIC SIZE The radius of an atom is found from the distance between the nuclei in a molecule of two touching atoms, and then halving that distance.

  25. Trends in atomic radius in Periods 2 and 3

  26. Trends in atomic radius down a group • It is fairly obvious that the atoms get bigger as you go down groups. The reason is equally obvious - you are adding extra layers of electrons.

  27. Trends in atomic radius across periods • You have to ignore the noble gas at the end of each period. Because neon and argon don't form bonds. • Leaving the noble gases out, atoms get smaller as you go across a period.

  28. Explaining the increase in atomic radius The radius of an atom is governed by • the number of layers of electrons around the nucleus • the pull the outer electrons feel from the nucleus. • Compare lithium and sodium: • Li • Na

  29. In each case, the outer electron feels a net pull of 1+ from the nucleus. The positive charge on the nucleus is cut down by the negativeness of the inner electrons.

  30. The only factor which is going to affect the size of the atom is therefore the number of layers of inner electrons which have to be fitted in around the atom. • The more layers of electrons you have, the more space they will take up - electrons repel each other. • This means that the atoms are bound to get bigger as you go down the Group.

  31. Trends in First Ionisation Energy • First ionisation energy is the energy needed to remove the most loosely held electron from one mole of gaseous atoms.

  32. The state symbols - (g) - are essential. When you are talking about ionisation energies, everything must be present in the gaseous state. • Ionisation energies are measured in kJ mol-1 (kilojoules per mole). They vary in size from 381 KJ (which you would consider very low) up to 2370 KJ (which is very high). • All elements have a first ionisation energy. Example: Helium (1st I.E. = 2370 kJ mol-1) Large amount of energy that is needed to remove one of its electrons. (Break a complete shell)

  33. Patterns of first ionisation energies in the Periodic Table The first 20 elements

  34. Factors affecting the size of ionisation energy - Ionisation energy is a measure of the energy needed to pull a particular electron away from the attraction of the nucleus. A high value of ionisation energy shows a high attraction between the electron and the nucleus. The size of that attraction will be governed by: • The charge on the nucleus. - The more protons there are in the nucleus, the more positively charged the nucleus is, and the more strongly electrons are attracted to it. • The distance of the electron from the nucleus. - Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away.

  35. The patterns in periods 2 and 3

  36. Explaining the general trend for ionisation energiesacross periods 2 and 3 • The general trend is for ionisation energies to increase across a period. • The major difference is the increasing number of protons in the nucleus as you go from lithium to neon. • This causes greater attraction between the nucleus and the electrons and so increases the ionisation energies. • The increasing nuclear charge also drags the outer electrons in closer to the nucleus. • That increases ionisation energies still more as you go across the period.

  37. Removal of successive electrons • Second ionisation energy is defined by the equation: • It is the energy needed to remove a second electron from each ion in 1 mole of gaseous 1+ ions to give gaseous 2+ ions. More ionisation energies • You can then have as many successive ionisation energies as there are electrons in the original atom. :

  38. The first four ionisation energies of aluminium, for example, are given by 1st I.E. = 577 kJ mol-1 2nd I.E. = 1820 kJ mol-1 3rd I.E. = 2740 kJ mol-1 4th I.E. = 11600 kJ mol-1

  39. In order to form an Al3+(g) ion from Al(g) you would have to supply: 577 + 1820 + 2740 = 5137 kJ mol-1

  40. Why do successive ionisation energies get larger? • Once you have removed the first electron you are left with a positive ion. Trying to remove a negative electron from a positive ion is going to be more difficult than removing it from a neutral atom. • Removing an electron from a 2+ or 3+ (etc) ion is going to be progressively more difficult.

  41. Using ionisation energies to work out which group an element is in • A big jump between two successive ionisation energies is typical of suddenly breaking in to an inner level. (new shell) • You can use this to work out which group of the Periodic Table an element is in from its successive ionisation energies.

  42. Magnesium (2,8,2) is in group 2 of the Periodic Table and has successive ionisation energies:

  43. Here the big jump occurs after the second ionisation energy. • It means that there are 2 electrons which are relatively easy to remove (the 2 electrons in last shell), while the third one is much more difficult (because it comes from an inner level closer to the nucleus and with less screening).

  44. Silicon (2,8,4) is in group 4 of the Periodic Table and has successive ionisation energies: :

  45. Decide which group an atom is in if it has successive ionisation energies?

  46. Plotting graphs of successive ionisation energies • Chlorine has the electronic structure 2,8,7 • This graph plots the first eight ionisation energies of chlorine. • The green labels show which electron is being removed for each of the ionisation energies.

  47. .

  48. The seventeenth ionisation energy of chlorine is nearly 400,000 kJ mol-1, and the vertical scale has to be squashed to accommodate this. .