The Periodic Table Chapter 5

Dmitri Mendeleev Dmitri Mendeleev developed the periodic table in 1869 Dmitri Mendeleev developed the periodic table in He grouped elements with similar chemical and physical properties together 2. He arranged the elements in increasing atomic mass. 3. He predicted the properties of elements not yet discovered.

Periodic Table Today’s periodic table, the elements are arranged in increasing atomic number so elements with similar properties fall in the same group. Today’s periodic table, the elements are arranged in increasing atomic number so elements with similar properties fall in the same group. Note: Ar and K both in according to atomic number, not atomic mass

Periodic Law Periodic Law: When elements are arranged in increasing atomic numbers, there is a repeating (periodic) pattern to the properties. See tables p. 142, 144, 148, 152. Periodic Law: When elements are arranged in increasing atomic numbers, there is a repeating (periodic) pattern to the properties. See tables p. 142, 144, 148, 152.

Group Names Group 1 elements – Alkali Metals – Group 1 elements – Alkali Metals – Very, very reactive metals Very, very reactive metals Group 2 elements – Alkaline Earth Metals – Group 2 elements – Alkaline Earth Metals – Very reactive metals Very reactive metals Group 17 elements – Halogens – Group 17 elements – Halogens – Very reactive nonmetals Very reactive nonmetals Group 18 elements – Noble gases – Group 18 elements – Noble gases – Nonreactive elements Nonreactive elements

Transition Metals Group 3 to 12 – Transition Metals Group 3 to 12 – Transition Metals Typical metals Typical metals Inner Transition Elements : lower two rows detached from main table Inner Transition Elements : lower two rows detached from main table Lanthanides – atomic numbers from 58 to 71 Lanthanides – atomic numbers from 58 to 71 Actinides – atomic numbers from 90 to 103 Actinides – atomic numbers from 90 to 103

Write electron configuration of the following elements. Then answer these questions. How many valence electrons does each have? Which group are these elements found in? Write electron configuration of the following elements. Then answer these questions. How many valence electrons does each have? Which group are these elements found in?

LiNaKRbBeMgCaSrFClBrHeNeAr

Blocks of Elements S-block elements- valence electrons are filling the s orbitals- Groups 1-2 D-block elements-valence electrons are filling the d orbtials- Groups 3-12 P-block elements-valence electrons are filling the p orbitals- Groups except He F-block elements-valence electrons are filling the f orbitals- Lanthanide and Actinide

Practice p.133, 136, 138, 139 and section review p139 p.133, 136, 138, 139 and section review p139 Practice reading off the electron configuration directly from the periodic table. Practice reading off the electron configuration directly from the periodic table.

Periodic Trends Atomic Radii – ½ the distance between the nuclei of 2 of the same atoms bonded together. Atomic Radii – ½ the distance between the nuclei of 2 of the same atoms bonded together. Atomic radii trends: Atomic radii trends: Going across a period – atoms get smaller – caused by increasing positive charge. Going across a period – atoms get smaller – caused by increasing positive charge. Going down a group – atoms get larger – adding energy level Going down a group – atoms get larger – adding energy level Prac P.142

Ions Ion is an atom with a positive or negative charge. Cation formed by loss of electron(s) and has a positive charge. Anion formed by gain of electron(s) and has a negative charge. Ion is an atom with a positive or negative charge. Cation formed by loss of electron(s) and has a positive charge. Anion formed by gain of electron(s) and has a negative charge. Illustrate example of atom, cation, and anion. Illustrate example of atom, cation, and anion.

Ionic Radii Ionic radii – positive ions are smaller than neutral atom due to loss of electron – negative ions are larger than neutral atom due to gain of electron Ionic radii – positive ions are smaller than neutral atom due to loss of electron – negative ions are larger than neutral atom due to gain of electron Ionic radii trends: same as atomic radii patterns Ionic radii trends: same as atomic radii patterns Going across a period – atoms get smaller – Going across a period – atoms get smaller – Going down a group – atoms get larger – Going down a group – atoms get larger –

Ionization Energy Ionization energy (IE) – energy it takes to remove an electron from a neutral atom Ionization energy (IE) – energy it takes to remove an electron from a neutral atom IE trends: IE trends: Going across period – IE increases – greater positive to negative attraction, harder to remove electron Going across period – IE increases – greater positive to negative attraction, harder to remove electron Going down a group – IE decreases – easier to remove an electron because it is farther from the nucleus Going down a group – IE decreases – easier to remove an electron because it is farther from the nucleus

Electronegativity Electronegativity – measure of the ability of an atom to attract electrons to itself Electronegativity – measure of the ability of an atom to attract electrons to itself Electronegativity trends: Electronegativity trends: Going across a period – increase – more positive charge in nucleus in same energy level Going across a period – increase – more positive charge in nucleus in same energy level Going down a group – decrease – nuclear attraction decrease with distance from nucleus Going down a group – decrease – nuclear attraction decrease with distance from nucleus Prac P.152 and section review p.154