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Unit 5 The Periodic Table The how and why Newlands -1865 Arranged known elements according to properties & order of increasing atomic mass Law of Octaves – pattern of chemical & physical properties repeated every 8 elements Mendeleev - 1869 1st periodic table (63 elements) Ordered by increasing atomic mass Predicted pattern of missing elements Started new rows and lined up columns to organize elements with similar properties Rearranged elements so similar properties would line up correctly Created The Modern Table Moseley- determined the atomic number for each known element. Elements are still grouped by properties Similar properties are in the same column Ordered by increasing atomic number Added a column of elements Mendeleev didn’t know about – noble gases Periodic Law When elements are arranged in order of increasing atomic number, elements with similar properties appear at regular intervals Horizontal rows are called periods There are 7 periods Vertical columns are called groups. Elements are placed in columns by similar properties. Also called families VIIIB IIB VIIB VIB VB 13 14 15 16 17 3A 4A 5A 6A 7A IB VIIIA VIIA VIA VA IVA IIIA IIIB 1 2 1A 2A IVB IIA IA Other Systems 3 4 5 6 7 8 9 10 11 12 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B 18 8A The elements in the s & p blocks are called the representative elements Transition metals D block elements These are called the inner transition elements and they belong here Three Classes of Elements Metals Nonmetals Metalloids Metals Metals Ductile – drawn into wires Malleable – hammered into sheets All solid at room temperature (except Hg- Mercury) Conductors of heat and electricity Group 1 are the alkali metals VERY reactive because one valence e• Found as compounds in nature • Not including H! Group 2 are the alkaline earth metals Still highly reactive but not as much so as alkali metals (2 valence e-) Transition Metals The weird ones… May lose different #s of valence electrons depending on the element with which it reacts Less reactive than alkali or alkaline earth metals Inner Transition Metals 1st row = lanthanides • Shiny metals similar in reactivity to alkaline earth metals 2nd row = actinides • Unstable nuclei – all radioactive Non-metals Non-metals Most are gases, some solid, and 1 liquid (Br) More variation than metals Group 7 is called the Halogens Most reactive non-metals – 7 valence React frequently with alkali metals Group 8 are the noble gases Low reactivity, very stable, inert Metalloids or Semimetals Metalloids Border the staircase between metals and nonmetals Properties – similar to metals and nonmetals Some used in semiconductors (silicon in electronics) Part 2 Periodic trends Identifying the patterns What we will investigate Atomic size • how big the atoms are Ionization energy • How much energy to remove an electron Electronegativity • The attraction for the electron in a compound What we will look for Periodic trends • How those things vary as you go across a period Group trends • How those things vary as you go down a group Why? • Explain why these variations exist Atomic Size Where do you start measuring? The electron cloud doesn’t have a definite edge. Scientists focused first on diatomic elements -- measured more than 1 atom at a time Atomic Size } Radius Atomic Radius = half the distance between two nuclei of molecule Atomic Size - Periodic Trends The positive nucleus pulls on electrons Periodic trend • As you move across a period, elements have more protons • The charge on the nucleus gets bigger • The outermost electrons of each element are in the same energy level • So there is more pull on the outermost electrons as you move across Periodic Trends As you go across a period, the radius gets smaller. Same outermost energy level More nuclear charge Pulls outermost electrons closer Na Mg Al Si P S Cl Ar Atomic Size – Group Trends The positive nucleus pulls on electrons Group Trend • As you go down a group, you add energy levels • Outermost electrons not as attracted by the nucleus Shielding Increasing numbers of electrons between the nucleus and the valence electrons tends to decrease the force between the nucleus & the valence electrons + Shielding The electron on the outside energy level has to look through all the other energy levels to see the nucleus + The Shielding electron on the outside energy level has to look through all the other energy levels to see the nucleus A second electron has the same shielding In the same energy level (period) shielding is the same + Shielding As the energy levels changes the shielding changes Moving down the group • More energy levels • More shielding • Outer electron less attracted + Three No shielding One Two shields shield shields Group trends As we go down a group • Each atom has another energy level • More shielding • The atoms get bigger H Li Na K Rb Rb K Atomic Radius (nm) Overall Na Li Kr Ar Ne H 10 Atomic Number Atomic size IONIZATION ENERGY It’s all about stability Alkali metals are more stable if they lose an electron Example • Sodium ([Ne] 3s1) • Getting rid of the 3s1 electron makes sodium more stable and creates a sodium ion (Na1+) Ionization Energy The amount of energy required to completely remove an electron from a neutral atom. The energy required for the 1st electron is called the first ionization energy Ionization Energy 2nd ionization energy is the energy required to remove the second electron Always greater than 1st IE The 3rd IE is the energy required to remove a third electron Greater than 1st or 2nd IE The Symbol First H He Li Be B C N O F Ne 1312 2731 520 900 800 1086 1402 1314 1681 2080 Second Third 5247 7297 1757 2430 2352 2857 3391 3375 3963 11810 14840 3569 4619 4577 5301 6045 6276 Group trends As you go down a group first IE decreases • Valence e- farther from nucleus • More shielding Periodic trends First IE increases from left to right across a period • Increased nuclear charge from added proton • Electron shielding not an issue b/c valence are all in same energy level Exceptions at full and 1/2 full orbitals • Lower IE b/c offer stability to atom How to remember? LO HI LO Ne First Ionization energy He N F Na has a lower IE than Li Both are s1 Na has more shielding H C O Be B Li Na Atomic number Atomic number First Ionization energy Electronegativity Electronegativity There’s an electron tug of war between atoms in a compound The tendency for an atom to attract electrons to itself when it is chemically combined with another element How “greedy” Large electronegativity means the atom pulls the electron towards itself Group Trend As you move down a group • More shielding • Less attraction for electrons • Lower electronegativity Periodic Trend As you move across a period from left to right, • Nuclear charge increases • Greater electronegativity How to remember? LO HI LO All 3 trends LO HI LO
