The Periodic Table Unit 4
I. History A. Dmitir Mendeleev Russian chemist, 19th century Arranged elements by their properties Arranged by increasing atomic mass Groups: vertical groups-elements have similar properties Periods : horizontal rows Periodic Law: Properties of the element are a periodic function of their atomic mass O Now arranged by atomic number O Iodine and tellurium were out of order
B. Henry Mosely British physicist (1887-1915 years of accomplishment) Developed the modern periodic table Used x-rays to determine atomic number
II. Elements Arranged based properties 109 elements-mostly naturally occurring Any element greater than 83 is radioactive
III. Metals, Nonmetals, Metalloids A. Metals Make up 2/3rds of the periodic table Shiny Solids (not Hg) Malleable Ductile Good conductor of heat and electricity Mobile electrons Tend to lose electrons to become ions
B. Nonmetals Not shiny Gas, liquid, and solids Not malleable or ductile Brittle Poor conductors of heat and electricity
C. Metalloids or semi-metals In between in properties On stairs Be, Si, Ge, Sa, Sb, Te, Po, At
IV. Groups or Families A. Alkali Metals (group 1) Very active metals, reactivity increases as you go down a group React violently with water Always found in compound
B. Alkaline Earth Metals (group 2) Active but not as much as group 1 Reactivity of metals increases as you go down a group
C. Transition Metals (group 3-12) Can lose 1-3 electrons to become ions Multiple oxidation states High melting points O Hard solid at Standard Temperature and Pressure (STP: 0 C and 1 atm) Mercury (Hg) is the exception-liquid Form colored ions in solution Reactivity of metals increases as you go down a group
D. Halogens (group 17) salt-formers Tend to bond with group 1 and 2 Very active non-metals Only group containing all three states of matter at room temperature
E. Noble Gases (group 18) Non-reactive Inert gases 8 valence electrons Octet rule: all elements want 8 valance electrons He exception- only 2 valence electrons
Q1. Explain the placement of an unknown element in the periodic table Q2. Compare and contrast metals, metalloids, nonmetals. Q3. Why are noble gases non-reactive
V. Trends A. Electron Configuration Although arranged by atomic number, there are significant trends for electron configuration Groups: same number of valance electrons (valence electrons determine how an element will react with other elements/ compounds) O Draw Li, Na, K Periods: same number of principle energy levels O Draw Na, Mg, Al
Lewis Dot Diagrams Used to determine the type(s) of covalent bonds that an element may make in certain situations Used to predict the type of ion that an atom might make when it forms an ion. Each dot diagram consists of an elemental symbol, which represents the kernel of the atom, and a group of 1-8 dots which shows the configuration of the valence shell electrons (outer-most electron shell of the atom).
Order for placing dots, two dots can start on any side, continue either clockwise or counter clockwise, fill one dot at a time Remember that each side can only hold up to two dots The number of valance electrons can be determined using the group number
Q4. Explain how the number of energy levels containing electrons can be used to determine the Period the element would be found on the periodic table Q5. Draw a Lewis electron-dot structure for Na, Be, Al, Ne Q6. Distinguish between valence and nonvalence electrons, given an electron configuration the following electron configurations. 2-1, 2-8-7, 2-8-18-8
B. Atomic radii (Size)-Table S Measure of the size of the atom Atomic radii measured as half the distance between 2 nuclei r = ½ d Groups: increase the number of principle energy levels as you go down Radii decrease Period: same number of principle energy levels as you go across Number of protons increases More pull for electrons Radii decrease
C. Ions Octet Rule: all atoms in nature want to look like noble gases (8 valance electrons) Metals Nonmetals Few valence electrons Lose electrons become + ions Close to 8 Want to gain electrons become ions
D. Shielding The electron shielding effect is the effect where core electrons block valence electrons from the nuclear charge of the nucleus. If you increase the number of principle energy levels, shielding increases Positive and negative charges attract each other so the more effective charge the electrons gets, the more attraction there is between the nucleus and the outer electrons. So as the effective nuclear charge increases, the atom and it's radii becomes smaller As the shielding becomes stronger, the nuclear charge decreases and the size of the atom increases-more shielding, bigger atom
Q7. Explain the trends of periods, in terms of nuclear charge and electron shielding seen on the periodic table
E. Ionization Energy Energy to remove the most loosely bound electron from a neutral gaseous atom Trend in periods From left to right, there is an increase in the number of protons which results in the nuclear charge increasing, the electrons are more strongly attracted and more energy is needed to remove them from the atom
Trends in Group Ionization energy decreases because valance electrons in each successive element are at a higher energy level and farther from the nucleus
F. Electronegativity Electronegativity value of an atom is a measure of its attraction for electrons when bonded to another atom Table S Periods: from left to right shows an increase in electronegativity Group: the highest electronegativity value is found at the top. Attraction for bonded electrons is less towards the bottom of the group