Periodic Table Chapter 6 Periodic Table l Many

Periodic Table Chapter 6

Periodic Table l Many different versions of the Periodic Table exist l All try to arrange the known elements into an organized table

Alternate Periodic Tables

History l 1869 - Russian chemist and teacher, Dmitri Mendeleev proposed a table for organizing elements l Mendeleev arranged the elements in a table based on increasing atomic mass.

History l Mendeleev placed elements next to each other with similar chemical properties l He would leave elements out of order based on atomic mass if they lined up better based on chemical properties

History l Mendeleev left spaces for elements not yet discovered – l He predicted properties of elements that would fit in those spot He predicted very closely the properties of Ge, Ga, Sc, and 5 others

History l 1913 - British physicist, Henry Moseley, determined the atomic numbers for the elements l The modern periodic table is arranged in order of increasing atomic number.

Periodic Table

Arrangement l Columns are called Groups – Numbered 1 -18 l Rows are called Periods l Elements in the same group have similar properties

Group Names l Group 1 - Alkali Metals l Group 2 - Alkaline earth metals l Group 17 - Halogens l Group 18 - Inert or Noble gases.

Group Names l Groups 3 -11 – Transition Metals l Bottom 2 rows – Inner Transition

Phases at STP l Most elements are solids at STP l Hg and Br are liquids at STP l H, N, O, F, Cl and Noble Gases are all gases at STP

Periodic Law l Periodic Law – When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.

Valence Electrons l Electrons in outermost occupied energy level l Valence Electrons are responsible for most chemical properties – Elements in the same group have similar properties because they have the same number of valence electrons

Classifying Elements l Elements are classified into 3 groups based on their properties: l Metals – Left and Middle l Nonmetals – Right l Metalloids - Staircase

Metals l l Good conductors of heat and electrical current High luster or sheen Many are ductile, meaning they can be drawn into wires Most are malleable, meaning they can be hammered into thin sheets

Metals l Metallic Character increases as you move towards the lower left l Most Metallic Element is Francium, Fr

Nonmetals l Most are gases at room temperature, some are solids, and one is liquid l Most are poor conductors l Most solids are brittle

Nonmetals l Non-Metallic Character increases as you move towards upper right l Most nonmetallic element is Fluorine, F

Metalloids l B, Si, Ge, As, Sb, Te l Have properties of both metals and nonmetals, based on conditions l Exceptions: – – Al and Po are metals At is a nonmetal

Group Characteristics l Alkali Metals (Group 1) – – – H, Li, Na, K, Rb, Cs, Fr All have 1 valence electron, tend to form +1 ions Most reactive metals Not found in nature by themselves, always combined with someone else Have properties of metals but are softer and less dense

Group Characteristics (cont) l Alkaline Earth Metals (Group 2) – – – Be, Mg, Ca, Sr, Ba, Ra All have 2 valence electrons, tend to form +2 ions Harder and more dense than alkali metals, but also have higher melting and boiling points Highly reactive, but not as much as alkali metals Not found by themselves in nature

Group Characteristics (cont) l Halogens (Group 17) – – – F, Cl, Br, I, At All have 7 valence electrons, tend to form -1 ions Strongly non-metallic Most active nonmetals Have low melting and boiling points Combine readily with metals to form salts

Group Characteristics (cont) l Noble Gases (Group 18) – – He, Ne, Ar, Kr, Xe, Rn Colorless gases that are extremely non-reactive Full valence shell, non-reactive All are found in small amounts in our atmosphere

Group Characteristics (cont) l Transition Metals (Groups 3 -11) – – – Most are excellent heat and electrical conductors Most have high melting points and are hard, except Hg Less active than group 1 and 2 metals Many combine with Oxygen to form oxides (Chemical property) Many have more than one oxidation number Form compounds that are colorful

Reminder l l STP Standard Temperature and Pressure – l 1 atm, 0°C Reference Point for most measurements

Diatomics l Eight elements are diatomic molecules when alone in nature (exist as two atoms bonded together) l H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2, At 2

Diatomics l Hydrogen and the Magic 7

Coloring l Color in the specific groups with your own color choices

Coloring l l Color in the different classifications with your own color choices Metals, Nonmetals, Metalloids

Orbital Blocks s - block p - block d - block f - block

Periodic Trends l How a property changes either across a period or down a group – – – Atomic Number Atomic Mass Atomic Radius Ionization Energy Electronegativity

Trends l Atomic number increases across a period. – l Increasing number of protons Atomic number increases down a group – Increasing number of protons

Trends l Atomic mass generally increases across a period. – l Increasing protons, neutrons, and electrons. Atomic mass increases down a group. – Increasing protons, neutrons, and electrons.

Radius l Atomic Radius – measure of the size of the atom – l Half the distance between two nuclei Ionic Radius – measure of the size of an ion

Trends l Atomic Radius decreases across a period – l More protons to pull on the electrons Atomic Radius increases down a group – Increasing electrons into more energy levels (more shells)

Ions l Atom, or group of atoms, that has gained or lost electrons l Cation – positive ion Anion – negative ion l

Ions l When an atom loses an electron, it becomes positively charged – – The radius becomes smaller Metals tend to lose electrons

Ions l When an atom gains an electron, it becomes negatively charged – – The radius becomes larger Nonmetals tend to gain electrons

Trends l Ionic Radius decreases for positive ions across a period – l More protons to pull on the electrons Ionic Radius decreases for negative ions across a period – More protons to pull on the electrons

Ionic Radius +1 +2 +3 +4 -3 -2 -1

Trends l Ionic Radius increases down a group – Increasing electrons into more energy levels (more shells)

Ionization Energy (IE) l Amount of energy required to remove an electron from an atom – Ca Ca+ + e- 590 k. J/mol l First ionization energy is removing the first electron l Second Ionization energy is removing the second electron after having the first removed – Ca+ Ca 2+ + e- 1145 k. J/mol

IE Trends l Ionization energy tends to increase across a period – l More protons are able to hold on tighter to electrons Ionization energy tends to decrease down a group – Electrons are farther away from the protons (more shells)

Electronegativity (EN) l Ability of an atom to attract an electron from another atom when in a compound. – – – Noble gases are usually omitted since they don’t form compounds Fluorine, F, is the most electronegative element with a value of 4. 0 Francium, Fr, is the least electronegative element with a value of 0. 7

EN Trends l Electronegativity tends to increase across a period – l More protons are able to attract electrons better Electronegativity tends to decrease down a group – Electrons are farther away from the protons (more shells)

Trends Summary Property Atomic Number Atomic Mass Atomic Radius Ionization Energy Electronegativity Period (L R) Group (T B)

Reactivity l l Elements that are more reactive tend to either gain or lose electrons very easily Elements that lose electrons easily have low IE and low EN – l Lower left, Fr Elements that gain electrons easily have high IE and high EN – Upper right, F