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Lecture 4 Ch.3, 4.1-4.3 Suggested HW Ch 3: 4, 28 Ch 4: 1, 4, 5, 12 Elemental Classifications • 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 Physical Characteristics of Nonmetals • Most nonmetals are gases • Non conductive of heat and electricity • Nonmetal solids are brittle, powdery • Low melting points O2 (g) S(s) He (g) All elements LEFT of the black line are metals, except Hydrogen. Chemical Groups and Diatomic Molecules • 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 H2 (Hydrogen gas) O O2 (Oxygen gas) N N2 (Nitrogen gas) F, Cl, Br, I F2, Cl2, Br2, I2 (Fluorine gas, chlorine gas, bromine gas and iodine gas) Nomenclature: Ionic Compounds • 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: Covalent Compounds • 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. Examples CO carbon monoxide CO2 carbon dioxide N2S P4Se10 dinitrogen sulfide tetraphosphorus decaselenide 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. H2S = 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. MgH2 = magnesium hydride; TeH2 = tellerium dihydride Group Work • Name the following: 1. 2. 3. 4. 5. SrO IF3 HBr CF4 NaH Chemical Reactions 𝐍𝐚 𝐬 + 𝐂𝐥𝟐 𝐠 → 𝐍𝐚𝐂𝐥 𝐬 ? • 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? Balanced Reactions • Mass can not be created or destroyed. This means that every element contained in the reactants must be accounted for in the product(s) 𝐍𝐚 𝐬 + 𝐂𝐥𝟐 𝐠 → 𝐍𝐚𝐂𝐥 (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) 𝐍𝐚 𝐬 + 𝐂𝐥𝟐 𝐠 → 𝟐 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. 𝟐 𝐍𝐚 𝐬 + 𝐂𝐥𝟐 𝐠 → 𝟐 𝐍𝐚𝐂𝐥 (𝐬) Coefficients And Subscripts 𝟐 𝐍𝐚 𝐬 + 𝐂𝐥𝟐 𝐠 → 𝟐 𝐍𝐚𝐂𝐥 (𝐬) • 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 Na Na Cl Cl NaCl NaCl • Do not confuse coefficients and subscripts. Do not alter subscripts when balancing. Tips For Balancing Reactions • 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 • Let’s balance the equation below using the rules from the previous slide. C3H8 (s) + O2 (g) CO2 (g) + H2O (L) • We’ll balance C first. C3H8 (s) + O2 (g) 3 CO2 (g) + H2O (L) • Now balance H. C3H8 (s) + O2 (g) 3 CO2 (g) + 4 H2O (L) • Now balance O. C3H8 (s) + 5 O2 (g) 3 CO2 (g) + 4 H2O (L) Group Work Balance the following: 1. Sulfur (s) + Oxygen gas (g) Sulfur trioxide (g) 2. Nitrogen gas (g) + Hydrogen gas (g) NH3 (ammonia gas) 3. C4H10 (L) + O2 (g) CO2 (g) + H2O (g) A B C n=4 n=3 n=2 n=1 Chemical Groups And Periodicity • 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 • 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 Decreasing metallic character 3 Li Highly reactive, highly conductive metal 4 Be 6 C Less reactive, less conductive metal Nonconductive, nonmetallic solid 9 F Highly reactive, diatomic, nonmetallic gas Decreasing metallic character 11 Na Highly reactive, highly conductive metal 19 K Highly reactive, highly conductive metal 14 Si 12 Mg Less reactive, less conductive metal 20 Be Less reactive, less conductive metal Slightly conductive semi-metal Decreasing metallic character 22 Ge Slightly conductive semi-metal 17 Cl Highly reactive, diatomic, nonmetallic gas 25 Br Highly reactive, diatomic, nonmetallic liq. 10 Ne Totally unreactive gas 18 Ar Totally unreactive gas 26 Kr Totally unreactive gas WHY? • 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? Explanation Of Elemental Groups • 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. Ionization Energy • 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 𝑀 → 𝑀+ + 𝑒 − 1st Ionization Energy 𝑀+ → 𝑀2+ + 𝑒 − 2nd Ionization Energy IE1 < IE2 < IE3 …….IEn Ionization Energy • 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. 1st Ionization Energy Shows A Periodic Trend For T very difficult to ionize very easy to ionize Ionization Energy • 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. Electron Arrangement (Electronic Structure) • 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. 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). 𝐵𝑒 𝑔 → 𝐵𝑒 + 𝑔 + 𝑒 − 𝐼𝐸1 = 1.49 𝑎𝐽 𝐵𝑒 + 𝑔 → 𝐵𝑒 2+ 𝑔 + 𝑒 − 𝐼𝐸2 = 2.92 𝑎𝐽 𝐵𝑒 2+ 𝑔 → 𝐵𝑒 3+ 𝑔 + 𝑒 − 𝐼𝐸3 = 24.7 𝑎𝐽 29.1 aJ Remember, energy is always in Joules (J). atto (a) = 10-18 Successive Ionizations • Let’s take a look at the electron configurations of Lithium (atomic # = 3) and Beryllium (atomic # = 4) Li 3 electrons Single electron that is easily removed Be 4 electrons Pair of tightly bound electrons Pair of electrons that are more easily removed Successive Ionizations Ne 10 electrons Na 11 electrons Eight electrons of similar attraction to the nucleus Same two tightly bound electrons 11th electron enters different “shell” 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. Na • Each shell holds 2n2 electrons, where the n=1 shell is the closest to the nucleus. Same Outer Electron Configuration Along A Group Leads to Similarities in Reactivity Na Li 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. Periodicity is Due To Repeating Valence Electron Configurations Li Be B C N O F Ne Na Mg Al Si P S Cl Ar 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 Full Lewis dot configuration Valence Lewis dot configuration • We use these representations to describe the electron configurations of an element.