Deciphering the Code of the Universe:

1. Deciphering the Code of the Universe: Simple Keys to Understanding and Interpreting the Periodic Table of Elements

2. The Periodic Table of Elements • Element: atoms of a class of substances that cannot be separated into simpler substances by normal (natural) chemical processes. • Each element has the same number of protons (Atomic Number) in their nucleus. • For several reasons, atoms sharing the same atomic number (elements) almost always share the same chemical properties: a property or characteristic of a substance that is observed during a reaction in which the chemical composition or identity of the substance is changed. Ex. Sodium (Na-atomic number 11) has has combustibility- meaning that it tends to ignite fire in chemical reactions. • Atoms of the same element can have different bonding structures and physical properties. Ex. 1: Isotopes of Carbon (6) have different numbers of neutrons, but still share atomic numbers. This results in differing • In the 19th century, many scientists began isolating various different elements. The created a need for chemists to be able to better organize elements, resulting in Dmitri Mendeleev’s Periodic Table.

3. The Periodic Table of Elements • Periodic: Appearing or occurring at regular cyclic and repeating intervals. This means that the periodic table is based of patterns (cycles) of repeating intervals of elements with similar properties. • The columns on the Periodic table are called groups. Each group shares the same number of valence electrons, giving them similar properties • The rows on the periodic table are called periods. Atomic number increases across a period, and each period has the same number of orbitals.

4. PERIOD Atomic Number GROUP Symbol/ Name Atomic Mass

5. Atomic Radius • Atomic Radius: Electrons exist in a proportionately large “probability cloud” around the nucleus of every atom. The atomic radius describes the size of the atom’s cloud, and it decreases moving right across a period, and increases moving down a group. • What it means: Every one knows that opposite forces attract. The electrostatic attraction between the positive nucleus and the negative electrons decreases as the distance between the two increases. This means that elements with a higher atomic radii are more likely to give up their valence electrons because of reduced attraction with their nuclei.

6. Atomic Mass • Atomic Mass (IAUPC AMU Definition): The mass of an atom measured in AMUs (Atomic Mass Units). 1 AMU mass of 1 nucleon (proton or neutron). Therefore, to find the Atomic Mass in AMUS, simply add the number of protons (Atomic Number) to the number of neutrons. • Atomic masses of elements that have differing numbers of neutrons (isotopes) are averaged. • Atomic mass tends to increase across a period and down a row. • Atoms (esp. metals) with greater atomic masses tend to be highly radioactive, poisonous or contain other unique properties. • Atomic Masses with large repeating decimals have many differing isotopes.

7. Reactivity (Bonding) I • Atoms of elements can combine using electrons to produce larger chains or groups of atoms call compounds. Bonding generally results in one of two types: • 1.) Covalent Bonds=A covalent bondis is a chemical bond that results in the sharing of electrons between atoms. Ex. H2O • 2.) Ionic Bonds= Chemical bond in which electrons are stolen from atoms. Ex. NaCl

8. Reactivity II • Reactivity: behavior in which atoms forms new substances by addition or subtraction of atoms from another reactant or reactants. • Several Factors influence reactivity, including: atomic mass, atomic radius, atomic number, electronegativity, ionization energy and (sometimes) electron affinity. • Electronegativity: The tendency of an atom to attract electrons to itself. Increases across a period, and decreases down a group. Greater the electronegativity, the greater the reactivity. • Ionization Energy: The amount of energy required to remove an electron from an atom. Increases across a period, and decreases down a group. Greater the ionization energy, less reactive (perhaps ionic only?) the substance.

9. Metallic Character • Metallic Characteristics: Summary of Common Properties: They are/have generally: Shiny (Luster), Solids at room temperature (excp. Hg/80), High melting points, High densities, Low ionization energies, Low electronegativities Large atomic radii,, Malleable, Ductile, Thermal conductors Electrical conductors. • Metals are the most numerous elements on the periodic table. • Metalloids have semi-metallic properties, and often have technological uses. • Most non-metals are gaseous at room temperature, and bond to eachother to create a covalent bond. • Metals bond to nonmetals in ionic bonding. Ex. Fe2O3

10. Radioactivity • Radioactive decay: the process by which an atomic nucleus of an unstable atom loses energy by emitting ionizing particles. Ex. Uranium(92) is highly radioactive, and was used in the development of the atom bomb. • Nuclei of Elements with atomic numbers > 83 (Bismuth) are highly unstable, and begin radioactive decay by the emission of 2 protons and 2 neutrons (helium nucleus/alpha particle), an electron (beta emission) or a highly dangerous gamma ray. • Because of this emission, radioactive elements continue decaying until the reach a stable nucleus. Ex. When Polonium (Po/84) emits 2 beta particles, what element will it become? Is it stable? What is its new atomic mass?

11. The Key to the Code Part II:Groups and Families

12. Group 1/ Alkali Metals • All of the Group 1 elements except Hydrogen: Lithium, Sodium, Potassium, Rubidium, Cesium and Francium. • Hydrogen is often classified in its own family, because it is not truly a metal. • “Alkali Metals are fun to through in lakes,” because they react fiercely with water in an explosion. • Alkali metals are very soft silvery metals that must be stored and handled with caution due to their reactivity. • Readily give up their only valence electron (low ionization energy) to form cationic bonds (become more positive).

13. Group 2/ Alkaline Earth Metals • Beryllium, Magnesium, Calcium, Strontium, Barium, Radium. • React similarly to Alkali Metals, except slightly less violently. • Low ionization energy, readily give up two valance electrons to form cations.