1. 1 Topic 3: Periodicity 3.1: The periodic table Essential Idea: The arrangement of elements in the periodic table helps to predict their configurations Nature of Science: Obtain evidence of scientific theories by making and testing predictions based on them – scientists organize subjects based on structure and function; the periodic table is a key example of this. Early models of the periodic table from Mendeleev, and later Moseley, allowed for the prediction of properties of elements that had not yet been discovered. (1.9)

2. 2 Topic 3: Periodicity 3.1: The Periodic Table Understandings: 1. The Periodic Table is arranged into four blocks associated with the four sub-levels: s, p, d, and g 2. The Periodic Table consists of groups (vertical columns) and periods (horizontal rows) 3. The period number (n) is the outer energy level that is occupied by electrons 4. The number of the principal energy level and the number of the valence electrons in an atom can be deduced from its position on the periodic table. 5. The periodic table shows the positions of metals, non- metals, and metalloids.

3. 3 Topic 3: Periodicity 3.1: The Periodic Table Applications and Skills: 1. Deduction of the electron configuration of an atom from the element’s position on the Periodic Table and vice versa.

4. 4 3.1 Background Information Development of the Periodic Table Johan Dobereiner Grouped similar elements into groups of 3 (triads) such as chlorine, bromine, and iodine. (1817-1829). John Newlands Found every eighth element (arranged by atomic weight) showed similar properties. Law of Octaves (1863).

5. 5 3.1 Background Information Development of the Periodic Table Dmitri Mendeleev Arranged elements by similar properties but left blanks for undiscovered elements (1869). Crash Course Video: https://www.youtube.com/watch? v=0RRVV4Diomg

6. 6 3.1 Background Information Development of the Periodic Table Henry Mosley Arranged the elements by increasing atomic number instead of mass (1913) Glen Seaborg Discovered the transuranium elements (93-102) and added the actinide and lanthanide series (1945)

7. 3.1 U2 Groups and Periods increasing atomic number Elements arranged by increasing atomic number into periods (rows) 1-7, which relate to energy levels groups or families (columns), which share similar properties 7

8. 3.1 U5. Metals, Non-metals, Metalloids 8 www1.whsd.net

9. 9 3.1 U5. Metals, Non-metals, Metalloids Metals –Left side of the periodic table (except hydrogen). –Good conductivity of heat and electricity –Luster (shiny) –Ductile (drawn into wires) –Malleable (hammered into sheets) –Lose electrons in chemical reactions (oxidized)

10. 10 3.1 U5. Metals, Non-metals, Metalloids Alkali metals: Group 1 Alkaline earth metals: Group 2 Transition metals: Group 3- 12, lanthanide & actinide series

11. 11 3.1 U5. Metals, Non-metals, Metalloids Nonmetals –Right side of the periodic table –Poor conductors of heat and electricity –Non-lustrous –Gain electrons in chemical reactions (reduced)

12. 12 3.1 U5. Metals, Non-metals, Metalloids –Halogens: Group 17 –Noble gases: Group 18

13. 13 3.1 U5. Metals, Non-metals, Metalloids Metalloids –Between metals and nonmetals, along the stair step line –Properties intermediate between metals and nonmetals - Some are semi-conductors

14. 14 3.1 U5. Metals, Non-metals, Metalloids Metalloids –Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As), Antimony (Sb), Tellurium (Te), Astatine (At)

15. 3.1 The Periodic Table 15 http://www.rsc.org/periodic-table Amazing Periodic Table Resource

16. 3.1 U4. Valence Electrons Valence Electrons: electrons in the outermost (highest) energy level of s and p sublevels –Group 1 elements s 1 = 1 –Group 2 elements s 2 = 2 –Group 13 elements s 2 p 1 = 3 –So on and so forth –Group 18 s 2 p 6 = 8 (except for helium, which has 2) 16

17. 3.1 U4. Valence Electrons Drawing Valence Electrons: (Lewis Dot Diagrams) Write the element symbol, starting at the top, add dots 1 at a time clockwise, separate to 1 per side up to 4, then double up. No more than 8 dots total. Group 1 = 1 dot Group 2 = 2 dots 17

18. 18 3.1 U4. Valence Electrons Lewis (electron) dot diagrams Group 1: 1 dotXGroup 15: 5 dotsX Group 2: 2 dotsXGroup 16: 6 dots X Group 3: 13 dotsXGroup 17: 7 dotsX Group 4: 14 dotsXGroup 18: 8 dots (except He)X

19. 19 IB Topic 3: Periodicity 3.2: Physical properties Essential Idea: Elements show trends in their physical and chemical properties across periods and down groups. Nature of Science: Looking for patterns – the position of an element in the periodic table allows scientists to make accurate predictions of its physical and chemical properties. This gives scientists the ability to synthesize new substances based on the expected reactivity of elements (3.1)

20. 20 IB Topic 3: Periodicity 3.2: Physical properties Understandings: 1.Vertical and horizontal trends in the periodic table exist for atomic radius, ionic radius, ionization energy, electron affinity, and electronegativity. 1.Trends in metallic and non-metallic behaviour are due to the trends above.

21. 21 IB Topic 3: Periodicity 3.2: Physical properties Applications and Skills: 1.Prediction and explanation of the metallic and non- metallic behaviour of an element based on its position in the periodic table. 1.Discussion of the similarities and differences in the properties of elements in the same group, with reference to alkali metals (group 1) and halogens (group 17).

22. 22 IB Topic 3: Periodicity 3.2: Physical properties Important Terms: Core Electrons: the inner non-valence electrons of the atom Nuclear Charge: the number of protons in the nucleus of the atom Shielding (screening): the core electrons shield (block) the valence electrons from the nucleus, reducing the nuclear charge

23. 23 3.2 U1. Atomic Radius Atomic Radii: size of the radius of the atom Group trend –Atomic size increases as you move down a group of the periodic table. –Reason: each row going down adds an energy level, increasing the size of the atom

24. 24 3.2 U1. Atomic Radius Atomic Radii: size of the radius of the atom Periodic trend across a period –Atomic size decreases as you move across a period. –Reason: The increase in nuclear charge (more protons in the nucleus) increases the attraction to the outer shell so the outer energy level progressively becomes closer to the nucleus decreasing the size of the atom

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26. 26 Atomic Radii- Your graph should look similar

27. 27 3.2 U1. Ionic Radius Ionic Radius: size of the ion’s radius Positive ions (cations) are smaller than their atoms. –Fewer electrons so nucleus attracts remaining electrons more strongly –One fewer energy level since valence electrons removed. Negative ions (anions) are larger than their atoms –More electrons so nucleus has less attraction for them –Greater electron-electron repulsion

28. 28 3.2 U1. Ionic Radius Ionic Radius: size of the ion’s radius Group trend –Ions get larger down a group –Reason: energy levels are added, electrons are farther away from nucleus

29. 29 3.2 U1. Ionic Radius Ionic Radius: size of the ion’s radius Periodic trend –Ions decrease as you move across a period. –Reason: This increase in nuclear charge increases the attraction to the outer shell so the outer energy level progressively becomes closer to the nucleus

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32. 32 3.2 U1. First Ionization Energy First Ionization Energy: The energy required to remove the first electron from a gaseous atom. *Second ionization removes the second electron and so on. Can be used to predict ionic charges. Group trend –Generally decreases as you move down a group in the periodic table –Reason: Since atomic radius increases down a group, the outermost electron is farther away from the nucleus and is easier to remove. The shielding effect of the core electrons also increases.

33. 33 3.2 U1. First Ionization Energy First Ionization Energy: The energy required to remove the first electron from a gaseous atom. Periodic Trend –Increases as you move from left to right across a period. –Reason: effect of increasing nuclear charge (more protons) makes it harder to remove an electron, stronger attraction between the nucleus and the outer electrons, the atomic radius also decreases, electrons are closer to the nucleus, more difficult to remove

34. 34 Filled n=1 shell Filled n=2 shell Filled n=3 shell Filled n=4 shell Filled n=5 shell

35. 35 3.2 U1. Electronegativity Electronegativity: The relative attraction an atom has for the shared pair of electrons in a covalent bond. *Helps predict the type of bonding (ionic/covalent). Group trend –Generally decreases as you move down a group in the periodic table. –Reason: atomic radius increases and shielding effect of core electrons, electrons are farther away, less likely to attract more –For metals, the lower the number the more reactive. –For nonmetals, the higher the number the more reactive.

36. 36 3.2 U1. Electronegativity Electronegativity: The relative attraction an atom has for the shared pair of electrons in a covalent bond. Periodic Trend –Increases as you move from left to right across a period. –Reason: nuclear charge increases, atomic radius increases creating greater attraction of electrons –Nonmetals have a greater attraction for electrons than metals.

37. 37 Electronegativity

38. 38 3.2 U1. Electron Affinity Electron Affinity: the energy required to detach an electron from a singly charged negative ion in the gas phase. Group Trend –Generally become less negative as you move down a group –Patterns vary by group, do not show a clear trend down a group

39. 39 3.2 U1. Electron Affinity Electron Affinity: the energy required to detach an electron from a singly charged negative ion in the gas phase. Period Trend –Generally become more negative as you move across a period from left to right –Reason: gaining electrons makes negative ions more stable –Trends are not as well highlighted as other trends

40. 40 3.2 A2 Similarities and Differences in Groups Group 1: Alkali Metals Have 1 valence electron Shiny, silvery, soft metals React with water & halogens Oxidize easily (lose electrons) Reactivity increases down the group –Atomic radius increases and ionization energy decreases going down group, easier to lose an e-. Images.fineartamericacom www.Britannica.com www.elementsales.com www.greatmining.com

41. 41 3.2 A2 Similarities and Differences in Groups Group 17: Halogens Have 7 valence electrons Colored gas (F 2, Cl 2 ); liquid (Br 2 ); Solid (I 2 ) Oxidizer (gain electrons) Reactivity decreases down the group –Atomic radius increases, more difficult to gain an e-