The Periodic Table. Chapter 4

Transcription

1 The Periodic Table Chapter 4

2 What information can be determined from the periodic table?

3 How have elements been organized into the periodic table used today?

4 Periodic Table with the f block in its proper location:

5 Section 1: How Are Elements Organized? In 1869 Dmitri Mendeleev found that by placing elements in order of increasing atomic mass properties of elements were repeated. Each new row = properties repeated. This resulted in each column having elements with similar properties. Made the first periodic table! Able to predict missing elements using this repetition. Problem with ordering elements by atomic mass: some did not match properties of other elements in the same column. Needed to be switched around.

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8 Adjusting the Periodic Table About 40 years after Mendeleev s table, Henry Moseley made an important change to the periodic table: Organized elements by atomic number instead of atomic mass. Elements that had not previously fit into the correct column when ordered by atomic mass were fixed. This is the periodic table we still use today. Properties of elements repeat as a result of being ordered by atomic number. In other words, they exhibit periodicity. This is called the periodic law.

9 Additional Information Columns on the periodic table are called groups or families. Recall that these elements all have similar properties! This is because they have the same number of valence electrons, which means they will react in similar ways. Valence electrons: outermost electrons in an atom. We can easily determine the number of valence electrons by looking at group numbers in the s & p blocks. Rows are called periods (indicates the energy level). Recall that each row begins when properties begin repeating again.

10 Group Numbers & Valence e- 1 1 IA 1A Numbers in bold are the number of valence e- for each group VIIIA 8A

11 How are elements grouped on the periodic table?

12 Green = metals Blue = metalloids Yellow = nonmetals

13 Metals Most elements on the periodic table are metals. Conduct electricity and heat. Ductile Can be drawn into a wire. Malleable Can be hammered or rolled into sheets. Usually lustrous Look shiny. Dull in air or oxygen. Solids at room temperature (except Hg).

14 Nonmetals Opposite characteristics from metals: Do not conduct electricity and heat well. Not very ductile. Are not lustrous. Can be solids, liquids, or gases at room temperature.

15 Transition Metals Groups d block elements. Can lose a different number of valence electrons. Less reactive than other metals we will look at (alkali and alkaline earth metals). Some like Pd, Pt, and Au are very unreactive.

16 Rare Earth Metals f block- 2 rows at the bottom of the table. Fit into rows 6 & 7 (look for * or other symbol). Lanthanide & Actinide series Lanthanides = 4f Reactive (like alkaline earth metals we will look at). Actinides = 5f All of them are radioactive. Nuclei are unstable and break down.

17 Other Properties of Metals Varying melting points. Example: W = 4322 o C and Hg = -39 o C Used to make alloys. Alloys: Homogeneous mixtures of metals. New properties result from mixing metals. Example: Brass = copper and zinc. Harder than copper alone. More resistant to corrosion. Others include steel, stainless steel, sterling silver.

18 Groups Main group elements s & p block elements Groups 1,2 and 3-8 (or 13-18). Group 1(A) = Alkali Metals H is NOT included! Group 2(A) = Alkaline Earth Metals Group 7(A) (or 17) = Halogens. Group 8(A) (or 18) = Noble Gases. Remember: the group/column number tells you how many valence electrons those elements have!

19 Alkali Metals VERY REACTIVE! React with water to make alkaline/basic solutions. Stored in oil to keep them from reacting with air and water. Only 1 valence electron to lose- a filled valence shell is very stable. Not found pure in nature, but combined with other elements (as compounds). Soft can be cut with a knife. Usually lustrous but will dull in contact with air. Form an oxide layer.

20 Alkaline Earth Metals Also highly reactive. Less reactive than alkali metals. Have 2 valence electrons to lose. Also found as compounds, rather than pure substances. Harder and higher melting points than group 1. Often found as minerals and ores in the Earth s crust.

21 Halogens Most reactive nonmetals. 7 valence electrons. Only need to gain one more electron to have a full valence shell and be stable. Frequently react with alkali metals. Recall that alkali metals have 1 valence electron to lose. Ex: NaCl, KF, LiBr Compounds formed from halogens typically are called salts.

22 Noble Gases Outermost energy level is completely filled with e -. s 2 p 6 = 8 valence electrons Exception: He, which is 1s 2. But the 1 st energy level does not have a p sublevel, so it is filled. Low chemical reactivity very stable. They have no desire to gain or lose electrons! Example He used for blimps. Typically inert thought to be completely unreactive. Exception: 1962, chemists were able to make some compounds with Xe. Recall the Hindenberg

23 Hydrogen Most common element in the universe. Group by itself very unique. Doesn t fit perfectly into the alkali metals or halogens. Only 1 proton and 1 electron. Can gain or lose an electron- very reactive and forms compounds with many other elements. If it loses its 1 electron, only a proton remains! This is unlike any other element!

24 What trends can be found on the periodic table?

25 Section 3: Trends in the Periodic Table Periodic trends exist since properties of elements repeat in the table. We will look at the following trends: ionization energy (IE) atomic radius electronegativity (e - neg) ionic size electron affinity MP/BP, density

26 Ionization Energy (IE) Ionization energy: energy needed to remove an electron from an atom (forms an ion- atom with a charge).

27 IONIZATION ENERGY

28 Atomic Radius (size) Atomic Radius: Half the distance between two bonded atoms nuclei. Hard to measure with only one atom due to e - cloud. How do we determine where it ends? Bond distance is easier to measure- then cut in half.

29 Atomic Radius Diagram Where should we consider the outside of the atom to be? Measured in picometers (pm) or Angstroms (Å). distance between two bonded atoms nuclei 2

30 Electronegativity Ability of an atom to attract electrons in a bond. Electrons from each atom are involved when atoms bond. Each atom s ability to attract e - is different. Linus Pauling invented a scale to indicate how well an atom can attract an e - in a bond. No units, just numbers. Ranges from F assigned 4.0 (highest value- has the greatest ability to attract e - when bonded). Noble gases don t have a value (don t need to form bonds- they are stable).

31 ELECTRONEGATIVITY

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33 Preview: Shielding

34 The Basics Shielding: inner electrons shield/block the valence electrons from the positive nucleus. j Li Going down a group. Li Going across a period. Q1: What happens to shielding as you move down a group? Why? Q2: What happens to shielding as you move across a period? Why? A1: Increases going down a group because more shells/energy levels are being added. A2: Stays the same going across a period because you re in the same shell/energy level.

35 Li The Basics Nuclear charge: positive charge in the nucleus. Increases as atomic number/protons increase. Li Be B j Na K Going down a group. +3 Going across a period Q1: What happens to nuclear charge as you move down a group? Why? Q2: What happens to nuclear charge as you move across a period? Why? A1: Increases going down a group because atomic number increases. A2: Increases across a period because atomic number increases.

36 The Basics Effective nuclear charge: how well valence e- can feel the positive nucleus based on shielding & nuclear charge. Greatest effective nuclear charge. +3 Going down a group. Li +3 Be B Li Na K Going across a period. Q1: What happens to effective nuclear charge as you move down a group? Why? Q2: What happens to effective nuclear charge as you move across a period? Why? A1: Decreases going down a group because shielding increases. A2: Increases across a period because atomic number increases & shielding stays the same.

37 IE, Atomic Radius, e - neg Down a Group 9p+ Effective nuclear charge decreases. SO 17p+ Atomic radius INCREASES. Ionization energy DECREASES. Electronegativity DECREASES. 35p+

38 IE, Atomic Radius, e - neg Across a Period 7p+ O 8p+ F 9p+ Effective nuclear charge increases. SO Atomic radius DECREASES. Ionization energy INCREASES. Electronegativity INCREASES.

39 Atomic Radius Cont.

40 Explaining Reactivity Recall that groups 1 and 7 are the most reactive metals and nonmetals. As we move down group 1, the alkali metals become more reactive- this is because of the trend seen in ionization energy! As we move down group 7, the halogens become less reactive- this is because of the trend seen in electron affinity (very similar to electronegativity). Electron affinity is how well an atom can gain an electron.

41 How are elements created?

42 Section 4: Where Did the Elements Come From? Only 93 of the elements are found in nature. 3 of these are not found on Earth. Technetium, Promethium, Neptunium Found in stars. Most living things contain C,H,N,O,P, & S. Compounds that contain carbon are called organic compounds. Found in living things. Big Bang Theory: elements were created when universe was formed in a violent explosion.

43 Big Bang Theory Cont. VERY high temperatures existed after the big bang. This form of energy cooled and formed matter (e -, p +, n). Further cooling allowed subatomic particles to join together to form H. Gravity pulled H clouds together and formed stars. Stars worked as nuclear reactors to form He (under high temperature and pressure). 4 H à 1 He + energy (gamma radiation) Other elements were formed as He and H combined (fusion) to form even heavier elements. Supernovas formed all elements heavier than iron. Star collapses and blows up, releasing heavier elements into space. This can emit more energy than the sun does in its life span!

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45 Supernovas

46 Synthetic & Superheavy Elements Transmutations: type of nuclear reactions that change one element into another element All elements greater than number 93 (except 61) are not naturally occurring synthetic elements. Particle accelerators can be used to create them. Different types exist. Nuclei collide and fuse together. Superheavy elements are those that have an atomic number greater than 100. Only exist for fractions of a second.