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Presentation on theme: "Periodicity."— Presentation transcript:

1 Periodicity

2 General info. Periodicity is concerned with trends or patterns seen within elements on the periodic table. The patterns that will be covered are: Atomic radius Ionization energy Electronegativity Melting points (IB/SL only) Period 3 (IB/SL only

3 Atomic radius

4 Remember, the radius is determined directly by the electron cloud
Atomic radii Going down a group (column) The radius increases going down a group. This increase is due to the addition of energy levels. Each e.l. increases the electron cloud significantly.

5 Going across a period. The radius decreases going across a period.
The nucleus increases, which means there is a greater positive charge present. This increase in positive charge pulls the electrons in closer to the nucleus, thus, reducing the radius. No energy levels are being added.

6 Chart of atomic radii

7 Chart showing radius trends across several periods

8 Atomic radii with d block elements

9 Ionic radii The radius of an atom increases when an electron is added to make a (-) ion (anion). Anions have more e- than protons so the additonal e- is not pulled in close to the nucleus. This causes the e- cloud to expand.

10 Ionic radii continued The radius of an atom decreases when an electron is lost to make a (+) ion (cation). Cations have more protons than electrons, so the electrons now have a stronger pull on them from the nucleus, thus shrinking the e- cloud. Also, usually when atoms lose any e-, it is all of the valence e- they lose, so the entire energy level is lost. This really decreases the e- cloud. Transitional metals do not lose an e.l.

11 Metal ionic radii

12 Nonmetal ionic radii

13 Ionic radii charts

14 Ionization energy (IE)
Ionization energy- the energy required to remove an electron from a gaseous atom. Since the electrons are attracted to the nucleus, it takes energy to pull the e- from the atom. Electrons in the last e.l. are the electrons that are being removed. These electrons are called valence electrons. D-block elements have their valence e- in the last two energy levels. First ionization energy. The energy required to remove the first valence electron from a neutral atom.

15 Trends in the first ionization energy
Going down a group The first ionization energy decreases. The radius is increasing and the further away an e- is from the nucleus, the less pull (attractive force) there is on the electron. The further away an e- is from the nucleus, there are more e- (core e-) between the nucleus and the valence e-, thus weakening the pull from nucleus. (this is called shielding.) With the addition of each new energy level the valence electrons become further away from the nucleus.

16 Going across a period The first ionization energy increases
The atomic radius decreases. The valence e- become closer to the nucleus going from left to right , therefore there is a stronger pull (attractive force) from the nucleus on the valence e-. So, more energy is required to remove an electron.

17 3-d chart on first ionization energies

18 First ionization energy trends across several periods.

19 Successive ionization energies (IE)
The energy required to remove the first e- from an atom is the 1st ionization energy. The energy required to remove a 2nd e- is the 2nd ionization energy and so on. The energy required to remove an e- increases with each successive e-. For example: Al: IE1= 577 IE2= 1815, IE3=2740 all in kJ/mole

20 Succesive ionization energies cont.
Successive ionization energies increase because. After each e- is removed, the radius becomes smaller making the remaining e- closer to the nucleus. There are more protons than e-, making the attractive force from the nucleus stronger on the core e-.

21 Successive ionization energies cont.
After all of the valence e- have been removed from an atom, there is a huge increase in the IE. IE for core e- are extremely high due to them being in a lower e.l. that is closer to the nucleus. There are many more protons than e-.

22 Chart of succesive ionization energies ( IE)

23 Electronegativity A measurement of an element’s ability to attract an electron from another atom within a bond. When 2 atoms of different elements bond, one atom is better attracting the electrons in the bond. This attraction is electronegativity.

24 Electronegativity continued
An atom is more electronegative if: The atom has a small e- cloud. The small e- cloud allows its nucleus to get closer to another atom’s electrons and pull off e- Larger nucleus: more attractive force on another atom’s electrons. In summary: volume/mass ratio should be small.

25 Electronegativity continued.
The most electronegative element is Fluorine. Fluorine is given an electronegative rating of 4.0 which is the highest rating. All other elements are compared to it and are given a relative rating. The trends in electronegativity are the same as the ionization energy: increases across a period, decreases down a group.

26 Chart of electronegativities

27 Electronegative trends for the first 5 periods.

28 Applying periodic trends to chemical reactivity
Metals are reactive if: (lose e- in reactions) Larger radius Lower ionization energy Fewer valence e- Nonmetals are reactive if: (gain e- in reactions.) Smaller radius High electronegativity More valence e-

29 Melting point trends Down a group: The melting points tend to decrease. The large radius decreases the attraction between atoms, thus, making a substance easier to melt. Across a period: the melting points increase until group 14. Then they decrease.

30 Melting points for the group 1 elements

31 States of matter for all the elements at 3507 C

32 Melting and boiling point trends for elements 1-95

33 Chemical and physicial properties
Alkali metals: Very soft metals Not very dense and so float on water Low melting points: Li=454K to Cs=302 K One valence e- with a low ionization energy. Readily react with nonmetals React with water to form a strong base. Example:2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g)

34 Continued: Halogens Halogen Fluorine Chlorine Bromine Iodine Color
Pale yellow Yellow green Red brown Black or purple State at room temp. Gas Liquid solid

35 Continued: Halogens All exist as diatomic molecules in their pure state. F2, Cl2 Br2 I2 Are nonpolar and have only van der waal forces when in their pure form. Slightly soluble in water. When in water they dissociate slightly: X2 + H2O  H+ + X- + HOX, The HOX is a weak acid but a strong oxidant due to the oxygen which reacts with other materials and bleaches them. Example is HOCl is in bleach. These acids are also toxic to microbes, and so make disinfectants and are used in water treatment.

36 Continued: halogens They combine with metals to produce ionically bonded salts called hallides. Ex. NaCl, Most are soluble in water, exception, AgCl, PbCl2 The reactivity decreases going down the group. (This is the opposite of the Alkali metals.)The radius increases which reduces the ability of the elements to attract electrons.

37 Trends in period 3 Left hand side: groups 1 and 2 alkali and alkaline metals. Large radii, low IE, and so react as metals, losing e- React with nonmetals to make solid ionic compounds. These solids are crystaline solids with high melting and boiling points. Ex. NaCl The solids easily dissolve in water to form ions. Oxides, such as MgO, K2O dissolve in water to form bases. Ex. MgO + H2O  Mg(OH)2 and can neutralize acids. ex MgO + 2HCl  MgCl2 + H2O

38 Middle of period 3 Group 3:Aluminum is a metal Its oxide is amphoteric, that is it can dissolve in either an acid or a base. Al2O3 + 6HCl  2AlCl3 + 3H2O, Al2O3 + 2OH- + 3H2O 2Al(OH)4- Makes crystalline solids like AlCl3 with high melting/boiling points. Group4: Si, the ionization energy too great to behave as a metal and so the first nonmetal appears. Its oxide will react with water to produce a weak acid. Ex. CO2 + H2O  H2CO3 (H+ + HCO3-) Make large crystalline solid networks with very high melting a boiling points. Ex. Graphite, diamond, silicon chips.

39 Group 5: Phosphorus: Makes primarily covalently bonded molecules.
Weak forces are between its molecules, with chlorine or anything else. Very low melting and boiling points. Oxides form acid in water (versions of phosphoric acid.) H3PO4

40 Oxides dissolve to form acid in water:
Group 6 Sulfur. Makes covalent compounds with weak intermolecular forces, low melting/boiling points. Oxides dissolve to form acid in water: SO2 + H2O  H2SO3 or SO3 + H2O  H2SO4

41 Group 7:Halogens Are nonmetals that exist as diatomic molecules. F2, Cl2, are gases, Br2 liquid, I2 solid. Only van der waal forces exist between molecules. These elements combine with metals to make water soluble ionic compounds that have high melting and boiling points. (Silver compounds are not water soluble.) Reactions are in previous slide.

42 Summary of period 3: trends seen.
Metal or nonmetal: Metallic  nonmetallic Physical state in pure form: solidgas. Pure element combines with water: bases  to acids Oxides react with water to form: bases amphoteric  acids, oxidants or bleaches. Compounds: crystalline solids (groups 1,2)  strong networks (group 4)  weak molecular solids (groups 5,6)  gases (group7) exception is iodine

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