November 07, 2014 Periodic Trends Remember from the "Periodic Table" Notes... • The periodic table is a tabular display of the chemical elements, organized by their atomic number, electron configuration, and recurring properties. • Periodic law: There is a periodic repetition of chemical and physical properties of the elements when they are arranged by increasing atomic number November 07, 2014 Atomic Radius Graph • What are some initial observations about the atomic radius data/graph? • What is atomic radius? November 07, 2014 Go finish the rest of the worksheet with your group! You have 20 minutes. November 07, 2014 Atomic Radius Trend Discussion • What happens to atomic radius as you go across the period? Why? • What happens to atomic radius as you go down the group? Why? November 07, 2014 Periodic Trends Notes Get your handout out! November 07, 2014 Why is it called a periodic table? • The properties of the elements in the table repeat in a "periodic" way (specific pattern). • Periodic law: There is a periodic repetition of chemical and physical properties of the elements when they are arranged by increasing atomic number • The modern periodic table is arranged by > atomic number = # of protons > properties > electron configuration November 07, 2014 Periodic Law • Now lets look at some properties of elements > We looked at some of these in "Meet My Family"! Alkali Metals Halogens November 07, 2014 Periodic Trends • Chemical properties of elements are determined by their electron configuration. • Properties are periodic because the number of valence electrons is periodic. November 07, 2014 Electron Configuration and the Periodic Table • Remember electrons are found in atomic orbitals > Principle energy level (n, shells) tells us the relative size and energy of atomic orbitals. > Each shell can hold a certain number of electrons. > # of e- = # of p+ for a neutral atom > Valence electrons = electrons in outermost shell 32 18 8 2 November 07, 2014 Electron Configuration and the Periodic Table • If we draw these orbital diagrams for H, He, Li, Be Na, Mg, 32 18 8 2 November 07, 2014 Electron Configuration and the Periodic Table • Periods: Period equals the highest principle energy level of electrons (shell of the valence electrons) • Groups: Elements in the same group have the same number of valence electrons November 07, 2014 Now lets look at two trends 1. Atomic radius 2. Ionization energy November 07, 2014 1. Atomic Radius The electron cloud surrounding a nucleus is based on probability and does not have a clearly defined edge. • Atomic radius: half the distance between nuclei of adjacent identical atoms > Adjacent nuclei in a crystal > Nuclei of two atoms bonded together http://en.wikibooks.org/wiki/High_School_Chemistry/Atomic_Size November 07, 2014 Atomic Radius • Atomic radius is determined by the strength of attraction between the valence electrons and the nucleus > # of protons in nucleus > # of shielding electrons that "block" protons from valence electrons November 07, 2014 Atomic Radius Trend • What happens to atomic radius as you go across the period? • What happens to atomic radius as you go down the group? November 07, 2014 So why does the atomic radius... • Decrease across a period? • Increase down a group? http://employees.csbsju.edu/cschaller/Principles%20Chem/atoms/atomperiodic.htm November 07, 2014 Atomic Radius • Moving down a group: Atomic radius increases > Principal energy levels increase (bigger shell) = increased orbital size, electrons are farther from nucleus http://employees.csbsju.edu/cschaller/Principles%20Chem/atoms/atomperiodic.htm http://www.chem.ox.ac.uk/vrchemistry/Machinery/html/page02.htm November 07, 2014 Atomic Radius • Moving across a period: Atomic radius decreases > Principal energy level remains the same > # of protons increases: electrons pulled in more http://chemwiki.ucdavis.edu/ http://chemwiki.ucdavis.edu/ http://employees.csbsju.edu/cschaller/Principles%20Chem/atoms/atomperiodic.htm November 07, 2014 Example 1: Order the following atoms from smallest to biggest atomic radius: Ga, Se, K I, F, Cl November 07, 2014 2. Ionization Energy • Ionization energy = energy required to remove an electron from a gaseous atom. (kJ/mol) When you remove an electron from an atom, what do you get? November 07, 2014 Ionization energy • Ionization energy is an indication of how strongly an atom's nucleus holds onto its valence electrons > Greater IE = harder to ionize • Ionization energy is always a positive value • You can take off more than one electron! > IE1 (1st ionization energy) = energy required to remove 1st valence e- > IE2 (2nd ionization energy) = energy required to remove a 2nd e- Na Na+ + e- H=495.8 kJ/mol Mg Mg+ + e- H=737.7 kJ/mol Mg+ Mg2+ + e- H=1450.6 kJ/mol November 07, 2014 Ionization Energy • Is an atom with a high ionization energy more likely or less likely to form a positive ion? • What about one with low ionization energy? What is the periodic trend for ionization energy? What factors affect ionization energy? November 07, 2014 Ionization Energy What pattern do you see? November 07, 2014 Ionization Energy • Moving down a group: IE decreases > There are more electrons in between the nucleus and the valence electrons (shielding effect) > Electrons are not as tightly bound to nucleus. > Electrons are further away, attraction decreases. • Moving across a period: IE increases > Increased nuclear charge (more p+, greater attractive force) November 07, 2014 Ionization Energy • What about successive IE (removing more than one e-)? > Successive IE increases • There is a big "jump" in some IE required. Why? > W/in a sublevel: small increases > Between sublevels: greater increase > Between energy levels: greatest increase http://www.avon-chemistry.com/p_table_lecture.html November 07, 2014 Example 2: In the following pairs, which atom is more difficult to ionize/Has the higher IE? Na or Li Na or Mg C or Na Cl or Al November 07, 2014 Using these trends, can you explain the alkali metal reactions?.