Chapter 5 Notes Crash Course Chemistry - Periodic Table

History of the Periodic Table Who is credited with developing a method that led to the determination of standard relative atomic masses? Stanislao Cannizzaro Who discovered the periodic law and what is it? Dimitri Mendeleev. The physical and chemical properties of the elements are periodic functions of their atomic numbers. Who established atomic numbers as the basis for organizing the periodic table? Henry Moseley

Name 3 sets of elements that were added to the periodic table after Mendeleev’s time? Noble Gases, Lanthanides, Actinides Why are elements’ atomic masses not in strict increasing order, even though the properties of the elements are similar? Periodic Table is arranged by atomic # and some masses are out of order because they are weighted average of the isotopes.

Electron Configuration Into what 4 blocks can the periodic table be divided? s, p, d, f Name the following groups of elements in the periodic table and state what type of element they are: Group 1 Alkali Metals - Metals Group 2 Alkaline Earth Metals - Metals Groups 3-12 Transition Metals - Metals Group 17 Halogens – Non-Metals Group 18 Noble Gases – Non-Metals

What determines the length of each period in the periodic table? Total number of electrons that can fill the outer level of the element of that period.

Alkali Metals Alkaline Earth Metals Halogens Noble Gases Lanthanides Transition Metals Non-Metals Metals Metalloids Actinides

Ch 5.3 Electron Configuration and Periodic Properties

Atomic Size } Radius Measure the Atomic Radius - this is half the distance between the two nuclei of a diatomic molecule.

#1. Atomic Size - Period Trends Going from left to right across a period, the size gets smaller. Electrons are in the same energy level. Outermost electrons are pulled closer. Na Mg Al Si P S Cl Ar

#1. Atomic Size - Group trends H As we increase the atomic number (or go down a group), each atom has another energy level. Valence electrons get further from the nucleus. So the atoms get bigger. Li Na K Rb

Trend in Atomic Radius

Ions An ion is an atom (or group of atoms) that has a positive or negative charge Atoms are neutral because the number of protons equals electrons Positive and negative ions are formed when electrons are lost or gained between atoms

#2: Ionic Group trends Li1+ Na1+ K1+ Rb1+ Cs1+ Ions therefore get bigger as you go down, because of the additional energy level. Na1+ K1+ Rb1+ Cs1+

Ionic Period Trends Across the period from left to right, they get smaller. N3- O2- F1- B3+ Li1+ Be2+ C4+

#3. Trends in Ionization Energy Ionization energy is the amount of energy required to completely remove an electron (from a gaseous atom). The energy required to remove 1 electron is called the first ionization energy.

Ionization Energy The second ionization energy is the energy required to remove the second electron. Always greater than first IE. The third IE is the energy required to remove a third electron. Greater than 1st or 2nd IE.

Ionization Energy - Group trends As you go down a group, the first IE decreases because... The electron is further away from the attraction of the nucleus, and There is more shielding.

Ionization Energy - Period trends All the atoms in the same period have the same energy level. So IE generally increases from left to right.

Metals tend to LOSE electrons, from their outer energy level Ions Metals tend to LOSE electrons, from their outer energy level Sodium loses one electron. There are now more protons (11) than electrons (10), and thus a positively charged particle is formed = “cation” The charge is written as a number followed by a plus sign: Na1+ Now named a “sodium ion” Lost an electron, so a decrease in size from atom to ion.

Nonmetals tend to GAIN one or more electrons Ions Nonmetals tend to GAIN one or more electrons Chlorine will gain one electron Protons (17) no longer equals the electrons (18), so a charge of -1 Cl1- is re-named a “chloride ion” Negative ions are called “anions” Gained an electron so increase in size from atom to ion.

Valence Electrons The electrons available to be lost, gained, or shared in the formation of chemical compounds. Main group elements: valence electrons are in the outermost s and p orbitals.

Electronegativity is the tendency for an atom to attract electrons. It decreases as it goes down a group because electrons get farther away from the nucleus.

Electronegativity Period Trend Metals (left side) They let their electrons go easily Low electronegativity Nonmetals (right side). They want more electrons. Try to take them away from others High electronegativity.

Summary of Trends Ionization Energy and Electronegativity Atomic and Ionic Radius