Periodic Table and Periodicity. BHS Chemistry

Transcription

1 Periodic Table and Periodicity BHS Chemistry

2 In 1869, Dmitri Mendeleev, a Russian chemist noticed patterns in certain elements. He discovered a way to arrange the elements so that they were organized by their chemical and physical properties.

3

4 Moseley s Contribution Henry Moseley is credited for further arranging the elements on the periodic table in order of the number of protons they contained.

5

6 Circle Periodic Table

7 Attention: New Additions to Periodic Table WOMANIUM (Wo) Physical properties: Boils at nothing and may freeze any time. Chemical properties: Very active and highly unstable. Possesses strong affinity with gold, silver, platinum, and precious stones. Violent when left alone. Turns slightly green when placed next to a better specimen. Usage: An extremely good catalyst for dispersion of wealth. MANIUM (Xy) Physical properties: Solid at room temperature but gets bent out of shape easily. Chemical properties: Becomes explosive when mixed with Childrium for prolonged period of time. Usage: Possibly good methane source. Caution: In the absence of WO, this element rapidly decomposes and begins to smell.

8 Periods The table is arranged in horizontal rows (going across) called periods. There are 7 periods. The period tells you how many electron energy levels the atom has.

9

10 Group The table is also arranged in vertical columns (going down) called groups. There are 18 groups. Members of each group have similar physical and chemical properties.

11

12 III. Properties of Metals, Nonmetals and Metalloids A. Properties of Metals 1. They are malleable, and have luster. 2. They reflect heat and light 3. Good conductors of heat and electricity. 4. Typically solids at STP 5. High melting points 6. They lose electrons in chemical reactions to become cations (+)

13 B. Properties of Nonmetals 1. They are dull and brittle. 2. Don t conduct heat and electricity well. 3. Low boiling and freezing points. 4. They exist in all three phases at STP, but most are gases. 5. They gain electrons in chemical reactions to become anions (negative ions)

14 C. Properties of Metalloids (B, Si, Ge, As, Sb, Te) 1. They possess intermediate properties between metals and nonmetals. 2. They are semiconductors at higher than room temperatures. 3. They are all solids at STP.

15 IV. Special Groups A. Alkali Metals Group 1 (IA) on the periodic table. 1. They are soft and easily cut. 2. They are highly reactive. (especially with H 2 O) Highly Reactive Video 3. All have an electron configuration ending in s 1 4. Gives up 1 electron in bonding (+1) 5. Has 1 valence electron

16 B. Alkaline Earth Metals Group 2 (IIA) on the periodic table. 1. They are less reactive than the alkali metals. 2. They have an electron configuration ending in s 2 3. Gives up 2 electrons to form a (2+) charge 4. Has 2 valence electrons

17 C. Halogens Group 17 (VIIA) on the periodic table. 1. They are highly reactive and react violently with hot metals. 2. They form diatomic molecules 3. They have an electron configuration ending in s 2 p 5 4. Accepts 1 electron to have a (-1) charge 5. Has 7 valence electrons

18 D. Noble Gases Group 18 (VIIIA) on the periodic table. 1. Extremely stable and unreactive 2. They exist as single atoms 3. They have an electron configuration ending in s 2 p 6 4. Has 8 valence electrons

19 E. Transition Metals Groups 3-12 (B groups) on the periodic table. 1. They possess characteristics of active metals to varying degrees. 2. They form compounds that are usually brightly colored.

20 F. Inner Transition Metals two rows at the bottom of the periodic table. 1. Many of the inner transition metals are radioactive. 2. Many of the actinides are synthetic.

21 The Periodic Table

22 The Periodic Law says: When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties. Horizontal rows = periods Vertical column = group (or family) Similar physical & chemical prop.

23 ALL Periodic Table Trends Influenced by three factors: 1. Energy Level Higher energy levels are further away from the nucleus. 2. Charge on nucleus (# protons) More charge pulls electrons in closer. (+ and attract each other) 3. Shielding effect

24 Shielding The electron on the outermost energy level has to look through all the other energy levels to see the nucleus. This effect decreases the attraction of the nucleus for the outer electrons..

25 What do they influence? Energy levels and Shielding have an effect on the GROUP ( ) Nuclear charge has an effect on a PERIOD ( )

26 Atomic Size } Radius Measure the Atomic Radius - this is half the distance between the two nuclei of a diatomic molecule.

27 #1. Atomic Size - Group trends As we increase the atomic number (or go down a group)... each atom has another energy level, so the atoms get bigger. H Li Na K Rb

28 #1. Atomic Size - Period Trends Going from left to right across a period, the size gets smaller. Electrons are in the same energy level. But, there is more nuclear charge. Outermost electrons are pulled closer. Na Mg Al Si P S Cl Ar

29

30 Ions Some compounds are composed of particles called ions An ion is an atom (or group of atoms) that has a positive or negative charge Atoms are neutral because the number of protons equals electrons Positive and negative ions are formed when electrons are transferred (lost or gained) between atoms

31 Ions Metals tend to LOSE electrons, from their outer energy level Nonmetals tend to GAIN one or more electrons

32 Ions Here is a simple way to remember which is the cation and which the anion: + + This is Anion. She s unhappy and negative. This is a cat-ion.

33 The size of an ion Ionic Radius Cations are smaller (lost e-) and anions are larger than the atoms they cam from.

34 Ionic radius Group trends Each step down a group is adding an energy level Ions therefore get bigger as you go down, because of the additional energy level. Li 1+ Na 1+ K 1+ Rb 1+ Cs 1+

35 Ionic radius Period Trends Across the period from left to right, the nuclear charge increases - so they get smaller. Notice the energy level changes between anions and cations. Li 1+ B 3+ N 3- O 2- F 1- Be 2+ C 4+

36 Trends in Ionization Energy Ionization energy is the amount of energy required to completely remove an electron.

37 Ionization Energy - Group trends As you go down a group, the first IE decreases because... The electron is further away from the attraction of the nucleus, and There is more shielding.

38 Ionization Energy - Period trends IE generally increases from left to right. Same shielding. But, increasing nuclear charge

39 The arrows indicate the trend: Ionization INCREASE in these directions

40 Trends in Electronegativity Electronegativity is the tendency for an atom to attract electrons to itself. An element with a big electronegativity means it pulls the electron towards itself strongly!

41 Electronegativity Group Trend The further down a group, the farther the electron is away from the nucleus, plus the more electrons an atom has. Thus, more willing to share. Going down a group, EN decreased

42 Electronegativity Period Trend Metals They want to lose electrons Low electronegativity Nonmetals. They want more electrons. Going across a period, the EN increases

43 The arrows indicate the trend: Electronegativity INCREASE in these directions

44 Textbook: Atomic Radius: pg. 141 Ionization Energy: pg. 143 Ionic Radius: pg. 149 Electronegativity: pg. 151