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5-1 Development of the Periodic Table Why have a table? • Chemists developed the Periodic Table to help organize and classify the elements. J.W. Dobereiner • Early 1800’s • Observed elements could be classified into sets of three – – – – Called Triads Ex. Li, Na, K Ex. Ca, Sr, Ba Cl, Br, I • Elements within a triad have similar chemical properties. – In each triad, the properties of the middle element are close to the averages of the properties of the 1st and 3rd element. – Ca Sr Ba (40 + 137) ÷ 2 = 88 40 88 137 J.A.R. Newlands • 1865 • By this time, 62 elements were known. • He observed that when elements were arranged in order of increasing atomic mass, the properties of the 8th element were like the 1st, 9th – Law of Octaves Dmitri Mendeleev • 1869 • Like Newlands, he saw a periodic repetition of properties when elements were arranged by increasing atomic mass – He developed the first periodic table • Elements in the same column have similar properties • He also knew that not all elements were known and left gaps in this table for these undiscovered elements – He predicted the properties of the undiscovered elements, and it turns out he was correct. • Example: Ekasilicon Germanium H.G.J. Moseley • Remember Mendeleev organized his periodic table in order of increasing atomic mass. – This was not quite correct, the elements are arranged in order of increasing atomic number. • Periodic Law – When elements are arranged in order of increasing atomic number, their physical and chemical properties show in a periodic pattern. • 5-1 Section Review • 1. What is the periodic law? • 2. How did Dobereiner and Newlands help to develop the periodic table • 3. How did Mendeleev demonstrate that his table was valid? What would have helped him develop the current PT? • 4. What did Moseley contribute to the development of the PT? • Do # 5 on p.163 5-2 Reading the Periodic Table Organizing Squares / Labeling and Naming Groups • Each square represents an element. It contains its symbol, atomic number and atomic mass. The table is read like a book, from left to right, top to bottom, in order of increasing atomic number. (Protons) • http://viewpure.com/DYW50F42ss8 • The stability of the atom depends on the electron configuration • The outermost electrons determine the chemical reactivity – These are called Valence Electrons Groups or Families • vertical columns with elements that have similar properties • 18 total. • The number with the “A” groups gives the number of valence shell electrons for all elements in the group. • As you move down a group, the boiling points and densities increase. Groups or Families Periods • Horizontal rows 1-7. • What does the row number tell you? Metals, Nonmetals and Semimetals (metalloids) • Metals-left side of PT. Characteristicsluster or shine, conductors of heat and electricity, usually solid at room temp, malleable, and ductile. Group 1 and 2. • Nonmetals-right side of PT. no luster, non conductors, non malleable or ductile, many are gases at room temp. Diverse properties • Metalloids (semimetals)-stair step from boron. Some properties of metals and nonmetals. Silicon is a metalloid, it is a semiconductor and therefore used in computer chips so you don't fry your unit. There are different “Blocks” on the Table S-Block • Groups IA and IIA • Outermost electrons of these elements are in sorbitals • Group IA 1 valence electron (Alkali Metals) • Group IIA 2 valence electrons (Alkaline Earth Metals) P- Block • Groups IIIA through VIIIA • Outermost electrons of these elements are in the p – orbitals • How many elements can each sublevel or block hold? • s= p= d= f= • How many electrons? • The shape of the periodic table is a result of the way the electrons fill the s, p, d, and f orbitals. 5-2 Section Review 1. Why do elements within a group have similar properties? 2. Color in a periodic table by orbitals (blocks), indicate the s, p, d, and f orbitals. 3. Describe the general difference between the elements on the right and the left side of the PT. 4. What information is present in each square of the PT? 5. Circle the ions most like the ions of Br? Be, Mg, F, O, C, B, I, Cr, F, K 6. Why would you expect the properties of those ions circled to be similar to Br ion? 5-3 Periodic Trends Atomic Radii one-half the distance between the nuclei of identical atoms that are bonded together. • Atom size increases going down a group because electrons are being added to higher energy sublevels. (don’t worry about memorizing exceptions). Atomic Radii • o Atom size decreases going across a periods left to right. – This is opposite what you would intuitively think. • Why? • As you move across a period, the atom gains electrons as well as protons. Because the effective nuclear charge has increased, the electrons will therefore be pulled closer to the nucleus, causing the atom to become smaller. Ionic Radii – the size of an ion. • Cation – a positive ion:Na+,Mg2+,Fe2+, Fe3+, – Formed by the loss of an electron – Cations are smaller than their neutral atom. • As an electron is lost, the effective nuclear charge is greater, causing the electron cloud to become pulled closer to the nucleus, making it smaller • Cations are pawsitive • Anion – a negative ion: Cl-, F-, N3-, … – Formed by the gain of an electron(s). – As an electron is added, the nucleus remains the same so the electrons are not drawn as close to the nucleus. Also, added electron causes repulsion among other electrons. Ionization Energy • Energy required to remove one electron from an atom • 1st ionization energy, 2nd ionization energy, 3rd ionization energy, … A + energy A+ + e- • Ion – an atom or group of bonded atoms that has a positive or negative charge. – Ex. Na+, Cl-, O2-, Fe3+, NH4+, CO32-, … – Monatomic ion – ion containing one atom – Polyatomic ion – ion containing atoms bonded together Ionization Energy • Noble Gases have the highest ionization energies because they have a full octet • Group IA have the lowest ionization energies – Ionization energy increases going across a period left to right and decreases down a group • Decreases down a group because electrons in outer energy levels are easier to remove • Increases across period because more electrons in an energy level make it more difficult to get to the magic number of 8 • It is possible to remove electrons from ions – Ex. Cu+ + energy Cu2+ + e– Each successive electron removal from an ion feels an increasingly stronger effective nuclear charge • Look at trend on fig 5-20 pg. 180 Electron Affinity • the energy change that occurs when an electron is acquired by a neutral atom • Most atoms release energy when they acquire an electron – A + e- → A- + energy • If energy is released, use negative sign • If energy is absorbed, use positive sign • Common unit is kJ/mol • Generally, the halogens gain electrons more readily; therefore, their values are most negative – This fact helps to explain their great reactivity • Notice electron affinities become more negative going across period • Generally, electron affinities become more negative going up a group – There are exceptions to this rule • It is also more difficult to add an electron to an already negative ion, so the second electron affinities are all positive. – Ex. Cl- takes less energy to form than Cl2- Electronegativity • a measure of the ability of an atom in a chemical compound to attract electrons – In a chemical compound, the more electronegative atom will attract electrons more – Ex. HCl : H-Cl : electrons are pulled closer to Cl • The most electronegative atom on the Periodic Table is Fluorine • Electronegativity tends to increase across a period • Electronegativity tends to decrease down a group