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Periodic Table And the Periodic Law Dmitri Mendeleev • Russian chemist • Created a table by arranging elements according to atomic masses • Noticed that chemical properties of the elements followed a repeating pattern Periodic Law – Henry Moseley who worked with Ernest Rutherford was the scientist who improved on Mendeleev’s periodic table by ordering the elements by the number of protons – the atomic number. • The physical and chemical properties of the elements are periodic functions of their atomic numbers. Mendeleev left spaces in his periodic table and predicted the existence of 3 elements and their 1. 2. 3. 4. Atomic numbers Colors Properties Radioactivity Arrangement of the table • The periodic table is an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column or group. Groups of the Periodic Table • 1st column (gold) alkali metals • 2nd column (purple) alkaline earth metals • 3rd-12th Transition Metals • 17th column – halogens • 18th – Noble Gases Electron Configuration and the Periodic Table • The periodic table can be divided into blocks which indicate which orbitals are filling. s Block – Groups 1and 2 • All are chemically reactive metals • Metals in group 1 (alkali metals) are more reactive than metals in group 2 (alkaline earth metals) • Neither are found in nature as elements since they react readily with water and nonmetals. Which is more reactive, magnesium or sodium? 1. Magnesium (Mg) 2. Sodium (Na) Name the element [Ar] 4s2 1. 2. 3. 4. Sodium Potassium Calcium Strontium Alkali Metals Video Hydrogen and Helium • Hydrogen is not an alkali metal. • Hydrogen is a unique element. • Helium is placed with the noble gases • Helium has a full outer shell of electrons. • Helium is a colorless gas, not a reactive metal. d Block – Groups 3-12 • Some of these elements don’t follow the diagonal rule exactly in electron distribution • Metals with typical metallic properties. – Good conductors of electricity – High luster • Called the transition elements. • Less reactive than s block metals Name the element [Ar]4s23d5 1. 2. 3. 4. Bromine Iron Magnesium Manganese p block – groups 13 – 18 except helium • p block and s block together are called the main group elements • Properties of this group vary – On the right side – nonmetals – On the left – metals – In between – metalloids (boron, silicon, germanium, arsenic, antimony, tellurium) Metals, Nonmetals, Metalloids • Metal- element that is shiny and malleable, and conducts heat and electricity • Nonmetal – Conducts heat and electricity poorly and is brittle • Metalloid – Element that has the properties of both metals and nonmetals, sometimes called a semiconductor An element which is shiny and brittle is likely a 1. 2. 3. 4. Metal Nonmetal Metalloid s-block element p-block • Group 17 – Halogens (fluorine, chlorine, bromine, iodine, astatine) – Most reactive nonmetals • Group 18 – Noble Gases (helium, neon, argon, krypton, xenon, radon) – Nonreactive and stable f-block: Lanthanides and Actinides • Lanthanides are shiny metals, with reactivity similar to the alkaline earth metals • Actinides are all radioactive, many are man-made Name the group 1. Alkali metals 2. Alkaline earth metals 3. Halogens 4. Noble Gases Name the group 1. Alkali metals 2. Alkaline earth metals 3. Halogens 4. Noble gases Name the section of the table 1. 2. 3. 4. S block P block D block F block What types of elements make up the p block? 1. 2. 3. 4. Metals Nonmetals Metalloids All of the above Periodic Trends • The arrangement of the periodic table shows directional trends for various properties of the atoms of each element. Nuclear Charge and Atomic Radius • Atoms decrease in size from left to right on the periodic table because of the increase in nuclear charge. IN A ROW IN THE PERIODIC TABLE, AS THE ATOMIC NUMBER INCREASES, THE ATOMIC RADIUS 1. Decreases 2. Remains constant 3. Increases 4. Becomes immeasurable Within a group of elements,as the atomic number increases, the atomic radius 1. 2. 3. 4. Increases Remains constant Decreases Varies unpredictably Ionization Energy • The energy required to remove an electron from a neutral atom of an element. – A large ionization energy shows that the electrons of an atom are bound more tightly to the nucleus and it is more difficult to remove the electron. Which group of elements has the highest ionization energies? 1. 2. 3. 4. Alkali metals Halogens Noble gases Alkaline earth metals The energy required to remove an electron from an atom is called the 1. 2. 3. 4. Electron affinity Electron energy Electronegativity Ionization energy Ionization Energy (IE) Trends • IE generally increases across each period. – As the nuclear charge increases, the electrons are held more tightly • IE decreases down a group – As the electrons reside farther from the nucleus, they can be removed more easily. Electron Affinity • The energy change that occurs when an electron is acquired by a neutral atom is called the atom’s electron affinity. • Halogens have high electron affinities because acquiring an electron will give them a full outer shell which increases the stability of the atom. • Half filled orbitals also give increased stability, so that the electron affinity of carbon is greater than the electron affinity of nitrogen. Electron Affinity Trend on the Periodic Table • Generally increases across a period. • Generally decreases down a group. When an electron is added to a neutral atom, a certain amount of energy is 1. Always absorbed 2. Always released 3. Either released or absorbed 4. Transferred to the more electronegative element Which represents a neutral atom acquiring an electron in a process where energy is released? 1. 2. 3. 4. A + e- + energy AA + e- A- + energy A + e- A- - energy A- + energy A + e- Ionic Radii • Cation – positive ion – Formation of a cation decreases the atomic radius • Anion – negative ion – Formation of an anion increases the atomic radius As the atomic number of the metals of Group 1 increases, the ionic radius25% 25% 25% 25% 1. Increases 2. Decreases 3. Remains the same 4. Cannot be determined 0 of 30 1 2 3 4 Valence Electrons • The electrons available to be lost, gained, or shared in the formation of chemical compounds are referred to as the valence electrons. • Valence electrons are those filling s and p orbitals. Valence electrons are those s and p electrons 1. Closest to the nucleus 2. In the lowest energy level 3. In the highest energy level 4. Combined with protons The number of valence electrons in Group 17 elements is 1. 2. 3. 4. 7 8 17 Equal to the period number The electrons available to be lost, gained or shared when atoms form compounds are called 1. 2. 3. 4. Ions Valence electrons d electrons Electron clouds Electronegativity • A measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound. Fluorine is assigned a value of 4 and all other elements have values relative to it. The element that has the greatest electronegativity is 1. 2. 3. 4. Oxygen Sodium Chlorine Fluorine