1. Lecture 4 Ch.3, 4.1-4.3 Suggested HW Ch 3: 4, 28 Ch 4: 1, 4, 5, 12

2. Elements can be metals, nonmetals, or semiconductors (we will discuss semiconductors later) Physical Characteristics of Metal Malleable Ductile Conductive of electricity Conductive of heat Have luster and shine Very High Melting Points Elemental Classifications

3. S(s) O 2 (g) Elemental Classifications He (g) Physical Characteristics of Nonmetals Most nonmetals are gases Non conductive of heat and electricity Nonmetal solids are brittle, powdery Low melting points

4. All elements LEFT of the black line are metals, except Hydrogen.

5. Certain elements are unstable, and hence, do not commonly exist as individual species, but as diatomic molecules These include H, O, N, and all of the halogens (group 17) H  H 2 (Hydrogen gas) O  O 2 (Oxygen gas) N  N 2 (Nitrogen gas) F, Cl, Br, I  F 2, Cl 2, Br 2, I 2 (Fluorine gas, chlorine gas, bromine gas and iodine gas) Chemical Groups and Diatomic Molecules

6. Molecules are formed by the atomic bonding. There are two types of bonds: IONIC and COVALENT. Special rules exist for naming molecules of each type. Ionic bonds form between metal ions and nonmetal ions (further detail provided in chapter 6) To name an ionic compound, you do the following 1.Write the name of the metal 2.Follow it with the ionic name of the nonmetal Example KF Potassium Fluoride Nomenclature: Ionic Compounds

7. Covalent bonds form between two nonmetals (further detail provided in chapter 7) To name an covalent compound: 1.Write the name of the first nonmetal. For non-unity subscripts, use greek prefixes (shown on right) 2.Follow that with the ionic name of the second nonmetal. Again, include greek prefixes. Only use mono- for oxygen containing molecules. Nomenclature: Covalent Compounds Examples CO carbon monoxide N 2 S dinitrogen sulfide CO 2 carbon dioxide P 4 Se 10 tetraphosphorus decaselenide

8. Nomenclature: Hydrogen Hydrogen is strange. It’s a nonmetal, but tends to react in a manner similar to that the metals in group 1. 1.If hydrogen is listed first (not including halogens), use ionic rules to name the molecule. Ex. H 2 S = hydrogen sulfide 2.Hydrogen halides (HX where X is a halogen) are acids and are named as such. We drop –gen and end the second nonmetal with the suffix “–ic acid” HCl = hydrochloric acid ; HF = hydrofluoric acid 3.If hydrogen is listed last, it is a hydride anion (H - ) Ex. MgH 2 = magnesium hydride; TeH 2 = tellerium dihydride

9. Group Work Name the following: 1.SrO 2.IF 3 3.HBr 4.CF 4 5.NaH

10. As you can see from the chemical equation shown to the left, products typically exhibit vastly different characteristics the reactants Also recall our discussion on the law of conservation of mass. Based on this law, can you find a problem with the equation written shown? Chemical Reactions

11. Mass can not be created or destroyed. This means that every element contained in the reactants must be accounted for in the product(s) 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. Balanced Reactions

12. The balanced equation above says that two Na atoms react with one chlorine gas molecule to produce two molecules of NaCl The coefficient of 2 means that there are two separate Na atoms The subscript of 2 indicates two Cl atoms bonded together in a single molecule Do not confuse coefficients and subscripts. Do not alter subscripts when balancing. Na Cl NaCl Coefficients And Subscripts

13. 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. 1.First, balance those elements that appear only once on each side of the equation 2.Balance the other elements as needed. Pay attention to subscripts. 3.Include phases Tips For Balancing Reactions

14. Let’s balance the equation below using the rules from the previous slide. C 3 H 8 (s) + O 2 (g) CO 2 (g) + H 2 O (L) We’ll balance C first. Now balance H. C 3 H 8 (s) + O 2 (g) 3 CO 2 (g) + H 2 O (L) C 3 H 8 (s) + O 2 (g) 3 CO 2 (g) + 4 H 2 O (L) Now balance O. C 3 H 8 (s) + 5 O 2 (g) 3 CO 2 (g) + 4 H 2 O (L) Tips For Balancing Reactions

15. Balance the following: 1.Sulfur (s) + Oxygen gas (g)  Sulfur trioxide (g) 2.Nitrogen gas (g) + Hydrogen gas (g)  NH 3 (ammonia gas) 3.C 4 H 10 (L) + O 2 (g)  CO 2 (g) + H 2 O (g) Group Work

16. As scientists first began to discover and classify the elements, patterns and similarities were observed in chemical behaviors of certain groups of 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, or vertical column. Chemical Groups And Periodicity

17. Dmitri Mendeleev created the periodic table in in 1869 by arranging the elements from left to right in order of increasing atomic number, and vertically according to their behavior (groups) In doing so, he observed repetitive patterns in chemical behavior across periods (horizontal rows) This periodicity is described in the next slide. Chemical Groups And Periodicity

18. Totally unreactive gas 3 Li 4 Be Highly reactive, highly conductive metal Less reactive, less conductive metal 6C6C 9F9F 10 Ne Nonconductive, nonmetallic solid Highly reactive, diatomic, nonmetallic gas Decreasing metallic character Totally unreactive gas 11 Na 12 Mg Highly reactive, highly conductive metal Less reactive, less conductive metal 14 Si 17 Cl 18 Ar Slightly conductive semi-metal Highly reactive, diatomic, nonmetallic gas Totally unreactive gas 19 K 20 Be Highly reactive, highly conductive metal Less reactive, less conductive metal 22 Ge 25 Br 26 Kr Highly reactive, diatomic, nonmetallic liq. Decreasing metallic character Slightly conductive semi-metal Chemical Groups And Periodicity

19. We must now answer many questions about chemical reactivity. – Why is it that some atoms join together and form molecules, while others can’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? WHY?

20. 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: The number of protons in the nucleus has no effect on chemical behavior. If it were so, chemical behaviors would never repeat given that no two elements have the same atomic number. The chemical behavior of an element must be dictated by the configuration of electrons around the nucleus. Explanation Of Elemental Groups

21. A direct indication of the arrangement of electrons about a nucleus is given by the ionization energies of 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 1 st Ionization Energy 2 nd Ionization Energy IE 1 < IE 2 < IE 3 …….IE n Ionization Energy

22. 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 1 st ionization energies of elements are plotted against atomic number. Ionization Energy

23. 1 st Ionization Energy Shows A Periodic Trend For T very difficult to ionize very easy to ionize

24. 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 METALLIC and REACTIVE it is. Ionization Energy

25. 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 arrangement of electrons around the nucleus of the element. Electron Arrangement (Electronic Structure)

26. Using the table of ionization energies in the previous slide, calculate the energy required to ionize Be to Be 3+ In order to go from Be to Be 3+, 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 Example

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

28. Na 11 electrons Ne 10 electrons Same two tightly bound electrons Eight electrons of similar attraction to the nucleus 11 th electron enters different “shell” Successive Ionizations

29. 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 2n 2 electrons, where the n=1 shell is the closest to the nucleus. Na

30. Same Outer Electron Configuration Along A Group Leads to Similarities in Reactivity Li Na K All group 1 metals have 1 lone electron in the outermost occupied shell (valence shell). Higher energy shells exist, but are empty! Chemical properties of an element are determined by the outer electron configuration.

31. Periodicity is Due To Repeating Valence Electron Configurations Li Be B C NOF Ne Na Al Si PS Cl Ar Mg

32. 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

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