The Periodic Table Chapter 4 What information can be determined from the periodic table? How have elements been organized into the periodic table used today? Periodic Table with the f block in its proper location: Section 1: How Are Elements Organized? • 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. 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. 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. How are elements grouped on the periodic table? Yellow = nonmetals Blue = metalloids Green = metals 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). 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. Transition Metals • Groups 3-12. ▫ 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. 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. Other Properties of Metals • Varying melting points. ▫ Example: W = 4322oC and Hg = -39oC • 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. 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! 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. 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. 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. Noble Gases • Outermost energy level is completely filled with e-. ▫ s2p6 = 8 valence electrons Exception: He, which is 1s2. But the 1st 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 Hydrogen • Most common element in the universe. • Group by itself – very unique. ▫ Only 1 proton and 1 electron. ▫ Can gain or lose an electron. What trends can be found on the periodic table? 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 melting & boiling points IE, Atomic Radius, and E-neg • Ionization energy: energy needed to remove an electron (forms an ion- atom with a charge). • Atomic Radius (size): 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. IONIZATION ENERGY 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 Electronegativity • Ability of an atom to attract an e- when bonded with another atom. • 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 0 – 4.0. 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). ELECTRONEGATIVITY Preview: Shielding The Basics • Shielding Effect: inner electrons shield/block the valence electrons from the positive nucleus. This results in less attraction. ▫ Increases going down a group. Why? ▫ Stays the same going across a period. Why? • Nuclear Charge: positive charge of the nucleus. Increases as the number of protons increases. ▫ Increases down a group. Why? ▫ Increases across a period. Why? • Combine these two and you get effective nuclear charge. The Basics Cont. • Effective Nuclear Charge: how well valence electrons can feel the attraction to the nucleus’ positive charge. ▫ Takes shielding (inner e-) and nuclear charge (#protons) into account . More shielding = less effective nuclear charge. ▫ Decreases going down. Why? Shielding increases. ▫ Increases going across. Why? Shielding stays the same (because the number of inner electrons stays the same). Nuclear charge increases because protons increase. IE, Atomic Radius, e-neg Across a Period • Effective nuclear charge increases. SO… • Ionization energy INCREASES. • Atomic radius DECREASES. • Electronegativity INCREASES. IE, Atomic Radius, e-neg Down a Group • Effective nuclear charge decreases. SO… • Ionization energy DECREASES. • Atomic radius INCREASES. • Electronegativity DECREASES. Atomic Radius Cont. http://intro.chem.okstate.edu/1314f00/lecture/chapter7/ATRADIID.DIR_PICT0003.gif Explaining Reactivity & Demonstration • 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. ▫ Electron affinity is very similar to electronegativity- it indicates how well an atom can gain an electron. How are elements created? 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. 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 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! Supernovas http://en.wikipedia.org/wiki/Supernova_remnant 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.