Dr. Harris Suggested HW: ( Ch 3) 4, 28 ( Ch 4) 1, 4, 5, 12

1. Lecture 4Periodicity, Ionization Energy and the proposed “Shell” structure of the atom; Ch 3, 4.1-4.3 Dr. Harris Suggested HW: (Ch 3) 4, 28 (Ch 4) 1, 4, 5, 12

2. Chemical Reactions • When elements undergo a chemical reaction, the products may be quite different from the reactants • The simplest reactions are those between metals and nonmetals. The product of such a reaction is an ionic compound • Lets consider the reaction between sodium metal and chlorine gas

3. Stark Differences Between Reactants and Products • As you can see from the chemical equationabove, productscan exhibit physical characteristics that are vastly different from those of the reactants • Recall the law of conservation of mass. Based on this law, can you find a problem with the equation written above?

4. Balancing Reactions • Mass can not be created or destroyed. This means that every element involved in a reaction must be accounted for in a chemical equation. • As you can see, there are two chlorine atoms on the reactant side, and only one chlorine atom one the product side. To balance the chlorine atoms, we add a coefficient of 2 to the NaCl(s) • We have balanced the chlorine atoms, but the sodium atoms are now unbalanced. We add a coefficient of 2 to the Na (s). The reaction is now balanced.

5. Coefficients vs. Subscripts • The balanced equation above says that two Na atoms react with onechlorine gas moleculeto produce two molecules of NaCl • The coefficientof 2 means that there are two separate Na atoms • The subscriptof 2 indicates two Cl atoms bonded together in a single molecule • Do not confuse coefficients and subscripts Na (s) NaCl (s) Cl Cl Na (s) NaCl (s)

6. Balancing Equations • Before carrying out any calculations, it is imperative that you first confirm that a given chemical equation is balanced. • The rules for balancing a chemical equation are provided below. • First, balance those elements that appear only once on each side of the equation • Balance the other elements as needed. Pay attention to subscripts. • Include phases

7. Balancing Equations • Let’s balance the equation below using the rules from the previous slide. C3H8 (s) + O2 (g) C3H8 (s) + O2 (g) C3H8 (s) + O2 (g) C3H8 (s) + 5 O2 (g) CO2 (g) + H2O (L) 3 CO2 (g) + H2O (L) 3 CO2 (g) + 4 H2O (L) 3 CO2 (g) + 4 H2O (L) • We’ll balance C first. • Now balance H. • Now balance O.

8. Group Examples • Balance the following

9. Chemical Groups • As more and more elements were discovered, chemists began to notice patterns in the chemical properties of certain elements. • Consider the three metals Li, Na, and K • All 3 metals are soft • All 3 metals are less dense than water • All 3 metals have similar appearance and low melting points • The most interesting feature is that all 3 metals react with the same elements in a nearly identical manner • As you see in the periodic table, these elements are all listed in the same group. • Elements in a group behave similarly. Recognizing patterns allows us to predict reactions without memorizing every characteristic of every element

10. http://www.youtube.com/watch?v=qRmNPKVEGeQ&feature=related http://www.youtube.com/watch?v=MTcgo46nxNE

11. Periodicity • Dmitri Mendeleev created the periodic table in in 1869 by arranging the elements in order of increasing atomic mass. • In doing so, he observed repetitive patterns in chemical behavior across periods • This periodicityis described in the next slide.

12. Periodicity Totally unreactive gas 25 F 9 F 1 H 11 Na 20 Be 10 Ne 12 Mg 19 K 26 Kr 4 Be 3 Li 18 Ar Decreasing metallic character 17 Cl 6 C 2 He 22 Ge 14 Si Less reactive, less conductive metal Highly reactive, highly conductive metal Highly reactive, diatomic, nonmetallic gas Totally unreactive gas Nonconductive, nonmetallic solid Decreasing metallic character Highly reactive, diatomic, nonmetallic gas Totally unreactive gas Less reactive, less conductive metal Slightly conductive semi-metal Highly reactive, highly conductive metal Decreasing metallic character Less reactive, less conductive metal Highly reactive, highly conductive metal Highly reactive, diatomic, nonmetallic liq. Totally unreactive gas Slightly conductive semi-metal

13. Mendeleev’s Genius • At the time in which the periodic table was being constructed, not all of the elements had been discovered. • Based on the observed periodicity, Mendeleev realized that gaps in the initial periodic table belonged to undiscovered elements • For example, in 1869, the element following Zn on the periodic table was As. Yet, he knew to put As in group 15 rather than 13 because As behaved like P, and he knew that two undiscovered elements (Gaand Ge) would fill the gaps.

14. Transition Metals • Transition metals span the region where the transition from metal to nonmetal occurs. • Transition metals are very denseand have very high melting points. transitions metals Semiconductors

15. Intro to Ch 4 • In ch. 4, we begin to answer many questions about chemical reactivity • Why is it that some atoms join together and form molecules, while others don’t? • Why is there such wide variation in the reactivity and physical properties of elements? • Why is there periodic repetition (periodicity) of the chemical/physical properties of elements as we move across the periodic table?

16. How to Interpret the Findings of Mendeleev • As previously discussed, Mendeleev noticed that chemical behavior was repeated periodically when elements were sorted by increasing atomic number • The existence of periodicity proves a very important point: Atomic number, and therefore, atomic mass, has no effect on chemical behavior. Otherwise, chemical behaviors would never repeat. Therefore, the chemical behavior of an element must be due to the configuration of electrons around the nucleus.

17. Ionization Energy • A direct indication of the arrangement of electrons about a nucleus is given by the ionization energiesof the atom • Ionization energy (IE) is the minimum energy needed to remove an electron (form a cation) completely from a gaseous atom • Ionizations are successive. • As you remove one electron, it becomes increasingly difficult to remove the next because of the increasing attraction between the remaining electrons and the protons in the nucleus 1st Ionization Energy 2nd Ionization Energy IE1 < IE2 < IE3 …….IEn

18. What Can Ionization Energy Tell Us About Chemical Behavior? • By measuring the energy required to remove electrons from an element, you can gain an idea of how “willing” an atom is to lose an electron, and relate this to its reactivity • In the next slide, you will see data from an experiment in which the 1st ionization energies of elements are plotted against atomic number.

19. 1st Ionization Energies Show A Periodic Trend

20. Trends in 1st Ionization Energies • It is relatively easy to remove electrons from group 1 metals. • It becomes increasingly difficult as you move right across the periodic table, and up a group. • It takes a very large amount of energy to ionize a noble gas. • Like chemical properties, ionization energies are also periodic. The lower the ionization energy of an element, the more METALLICand REACTIVEit is.

21. Electron Arrangement • The closer an electron is to the nucleus, the harder it would be to pull the electron away. • By carrying out multiple ionizations, we can gain insight into the arrangementof electrons around the nucleus of the element.

22. Example • Using the table of ionization energies in the previous slide, calculate the energy required to ionize Be to Be3+ • In order to go from Be to Be3+, you must LOSE 3 electrons. This will require 3 ionization steps (see pg 107 in book). 29.1 aJ Remember, energy is always in Joules (J). atto (a) = 10-18

23. Graph of Successive Ionization Energies • Let’s take a look at the electron configurations of Lithium (atomic # = 3) and Beryllium (atomic # = 4) Li 3 electrons Be 4 electrons Pair of tightly bound electrons Pair of electrons that are more easily removed Single electron that is easily removed

24. Larger Atomic Numbers Ne 10 electrons Na 11 electrons Same two tightly bound electrons Eight electrons of similar attraction to the nucleus 11th electron enters different “shell”

25. Electrons reside in “shells” of different distances from the nucleus • From these plots, Niels Bohr derived the Bohr model of the atom. In it, electrons reside in shells that orbit at different distances from the nucleus. • Each shell has a finite number of electrons that it can hold • The two electrons closest to the nucleus are the hardest to remove • Each shell holds 2n2 electrons, where the n=1 shell is the closest to the nucleus. Na atom

26. Same Outer Electron Configuration Along A Group Leads to Similarities in Reactivity Na Li K All group 1 metals have a lone electron in the outermost shell. Chemical properties of an element are determined by the outer electron configuration.

27. Periodicity is Due To Repeating Valence Electron Configurations Be Mg Ar Ne Si Cl C P O B N F S Al Na Li

28. Noble Gas Configurations • The inner-most electrons of an element comprise what is known as a noble gas core. • At the close of each shell, you have a noble gas configuration. Noble gases are chemically inactive because they have completely filled shells. • Lithium, for example, has a two electron core, which we call a Helium core, and one outer, or valence electron. Sodium has a 10-electron, Neon core, and one valence electron; and so on. • The electron configuration of an element can be represented with a Lewis dot formula

29. Full Lewis dot configuration Valence Lewis dot configuration We use these representations to describe the electron configurations of an element.