Compounds. I. Classify different forms of matter. A. Classification based on purity (of sorts)

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1 ompounds Reading: h 3 p.76-88, h 4 p , h 6 p.190 Problems: h 3.1, 3, 5, 7, 9, 11, 21, 23, 25, 27 (There is no charge on Ar.), 37, 41, 43, 45, 53, 55, 65a. h 4.3, 5, (Note: a line-bond structure is essentially a Lewis Dot Structure.) 7, 9, 11, 13, 15, 23, 25, 27, 29, 31, 33, 35, 37, 39. I. lassify different forms of matter. A. lassification based on purity (of sorts) B. Element: annot be separated into simpler substances by chemical means. 1. copper (u) 2. magnesium (Mg) 3. All those substances on the periodic table 4. Review: What determines the identity of an element? Note: With respect to elemental identity, an atom with 6 p + and 6 n in its nucleus is considered to be the same type of atom as one with 6 p + and 8 n in its nucleus. These two atoms are said to both be of carbon. 5. Sometimes we talk about a substance composed of only one type of atom as being in its elemental state. Elemental state means the state in which that element exists in nature. The elemental state of: a) e, helium, is as a gas, b) Fe, iron, is as a solid c) g, mercury, is as a liquid (Not many elements exist as liquids in nature.) Note: Seven elements are diatomic (two atoms of an elements bonded covalently) in their elemental state. I never order broiled clams from ardees.. ompounds: Pure substances formed when atoms combine in specific ratios. The combining is when the atoms form some kind of. 1. Sodium chloride, Nal 2. Water, 2 D. Mixtures: Two or more types of substances mixed together in varying ratios. They can be separated by physical means. 1. omogeneous: Uniform composition (also called a ) Example: 2. eterogeneous: Non-uniform composition. Ex:

2 II. Most matter on earth & in your body isn t in its elemental form. It is present as compounds or mixtures A. If we consider the lecture room (or lab?), can you identify any matter that is in its pure elemental form? B. What about outside of the lecture room? 1. the blue stuff is... elemental vs. compound pure vs. mixture homogeneous vs. heterogeneous Fig. from 2. the white stuff the green stuff.... Does classification of the stuff above depend on how big of a block of blue (white, green) stuff we chose to examine? III. Formation of two types of compounds from elements Intro (or intrusion?) of Sorts: Things at higher energy are less stable!! All living things are dependent on their ability to acquire energy from unstable things! The compounds in the food you eat must be relatively unstable for you to get useful energy from the food. Generally speaking, calorically rich food is made up of reduced (hydrogen containing) carbon compounds. The higher the hydrogen content, the more energy the food contains. (Ex.: Fat is more caloric than carbohydrate or protein. ne way to consider this: If something exists for a relatively long time, we say it is stable. If something exists for a brief time & then changes, it was unstable. What makes atoms etc. be stable? It depends on e. Focus: Lewis ctet Principle. Unstable atoms gain, lose or share electrons (forming ions or compounds) to end up with eight valence electrons (and become stable). For representative elements a full valence shell is stable. The valence e are those in the outermost s & p orbitals. We are going to look at ion and ionic compound formation first A. Elemental sodium (Na (s) ) reacting with chlorine (l 2(g) ) to form table salt (Nal (s) ), an : 2 Na (s) + l 2(g) 2 Nal (s)

3 Not safe to do, but see: & B. Important: 2 distinct things occur in these rxns. You must be able to recognize them as distinct. 1. First, e - are being transferred to convert atoms (or molecules) into ions: Na Na + + e - (ox) l e - 2 l - (red) metals tend to lose e - (give e - away), non-metals tend to gain e -. (Electronegativity: the relative ability of an element to attract electrons in a bond. Usually elements with high electronegativity tend to e -, and elements with low electronegativity tend to e -.) 2. Second, positively charged ions bind to negatively charged ions to form ionic compounds: Na + + l - Nal 3. Another example: Elemental magnesium (Mg (s) ) reacting with oxygen ( 2(g) ) to form magnesium oxide (Mg (s) ): 2 Mg(s) + 2 (g) 2 Mg(s) See: Two steps: 1. Mg(s) Mg e -, e Mg Mg positively charged ions are called negatively charged ions are called In the rxns shown above, a metal combined with a non-metal to form an ionic compound as the product a.k.a. a salt. The ions in the salt are held together by electrostatic attraction, + to -. IV. Ionic bonds, ionic compounds, ionic compound formulas A. In an ion: # of p + # of e -. B. Ions can be simple or polyatomic. 1. A simple ion consists of one charged atom (for example l - ).

4 2. A polyatomic ion has multiple atoms held together by covalent bonds and carries a charge. (for example N 3 - ). Ionic bond refers to the electrostatic attraction between oppositely charged ions. What do like charged things do? D. Ionic compound formulas must have zero net charge. (No shock when you touch an ionic compound.) 1. Determine charge of simple ions in ionic components from Periodic Table. ow? 2. A few polyatomic ions like sulfate ion, S 4 2-, are listed in the lab manual appendix and on p.78 of textbook. 3. An ionic compound formula consists of the minimum number of each ionic component that results in equal amounts of (+) and (-) charge. Problem. Write ionic compound formulas for: 1. potassium chloride 2. calcium iodide 3. sodium phosphide 4. sodium phosphate 5. magnesium oxide 6. iron (II) nitrate Remember: Ionic compounds are not molecules! Although the ionic compound formula, Nal, shows a 1 to 1 ratio of the ions, it is incorrect to think of a specific, directionally-oriented interaction between one Na + and one l -. You can see from the figure that each Na + is actually interacting equally with 6 l - ions (and vice versa).

5 An interesting thing happens to ionic compounds that dissolve in water. The lowest energy state occurs when the ions separate and each is surrounded by water. The 2 molecules reduce the strength of attractive or repulsion between ions significantly.. In M 109 we are interested in ions and ionic bonds, but we will have even more interest in covalent bonds. V. ovalent Bonding A. A second type of compound results from combining B. ydrogen gas ( 2(g) ) reacting with oxygen ( 2(g) ) to form water ( 2 (g) ): 2 2(g) + 2(g) 2 2 (g) 1. The indenberg was a air ship (specifically, a dirigible) whose buoyancy was provided by 2 (g). A known risk associated with this technology was hydrogen reacting violently with oxygen. Is there a readily available source of the reactant 2 on earth? Viewer warning, people died 2. In the rxn. shown above in, two non-metals combine to form a covalent compound. The product is a molecule. The atoms in the molecule are held together by covalent bonds.. What is a covalent bond? A covalent bond is a shared e - pair! VI. Electronic onfiguration and e sharing. A. The Periodic Table s shape helps you understand outer- (and inner-) shell e configuration. Which e were of greatest interest to us in predicting ion formation? These e also determine covalent bonding patterns. B. Different elements show different tendencies to give, take, or share valence e. We use the term electronegativity to describe this. 1. Elements that tend to give up e have electronegativity numbers. Metals tend to have low numbers.

6 2. Elements that tend to take e have electronegativity numbers. Non-metals have high numbers. 3. Sharing a blanket analogy. Electronegativity Table (Pauling scale) p. 106 in textbook 1A A 3A 4A 5A 6A 7A 8A 2 e 3 Li Be 1.5 < Atomic number < Elemental symbol < Electronegativity 5 B N F Ne 11 Na Mg 1.2 3B 4B 5B 6B 7B < B > 1B 2B 13 Al Si P S l Ar 19 K a Sc Ti V r Mn Fe o Ni u Zn Ga Ge As Se Br Kr 37 Rb s Sr Ba Y La Zr f Nb Ta Mo W Tc (98) 75 Re Ru s Rh Ir Pd Pt Ag Au d g In Tl Sn Pb Sb Bi Te Po (209) 53 I At Xe 86 Rn. Electronegativity lets us make predictions about bonding. 1. A metal bonded to a non-metal results in an ionic bond (electrons given and taken; not shared). In general, an electronegativity difference of > 2.1 is associated w/ ionic bond. 2. A bond between two non-metals results in a covalent bond. In general, an electronegativity difference of < 1.9 is associated w/ a covalent bond. Examples: Indicate whether the compounds below are ionic or covalent. Nal N 2 2 K 3 N VII. Lewis Structures of Atoms (These help with e accounting.) A. ount valence e (these are the e in the outer most shell). B. Arrange the e (dots) on the 4 sides of the elemental symbol.. Don t worry too much about transition metals. D. Examples: Xe F S e Important: Nobel (or inert) gases are inert! (non-reactive, stable) in the atomic state (ie.when # p + = # e ) because these atoms have full valence shells. (Part of The Lewis ctet Principle) Although this seems very simple, you can do a lot of neat chemistry with the Lewis ctet Principle.

7 VIII. Molecules are atoms joined by covalent bonds. Examples: water β-d-glucose adenine IX. Seeing patterns in molecules. A. Atoms share electrons to fill their valence e shells. (Why?) B. For valence e purposes, we count the shared e pair as belonging to both atoms. Is the stable elemental form of hydrogen? Why/why not? an we show the electron sharing in 2 with pictures? Lewis Dot Flat structure Ball & stick Space filling X. Electronegativity and Bond Polarity (view nuclei as fixed reference points.) ovalent bonds involve e sharing, but how evenly is the negative charge arranged? ow you decide? Use the electronegativity chart! A. Non-polar covalent bonds have completely even sharing. Examples: 2 and 2 Electronegativity:

8 B. Polar covalent bonds occur when bonding atoms have different electronegativities. 2.1 I F 4.0 δ + δ δ? δ? I slightly polar bond. F exceedingly polar bond XI. Representing molecules & compound ions w/ Lewis Structures A. Lewis Structures can be viewed as a visual bookkeeping device for valence electrons. B. The Lewis ctet Principle tells us that atoms with full valence shells are more stable than atoms with unfilled valence shells.. Since only the inert gas (Group VIII) elements have full valence shells in their atomic state, all other elements must gain, lose, or share e to achieve full valence shells. D. Steps for writing Lewis Structures: 1. Add up number of valence e (group #) for each atom 2. Make adjustments for non-zero net charge a) Add one valence e for each negative charge b) Subtract one valence e for each positive charge 3. Write elemental symbols (which atom is attached to which?) 4. Form single bonds to connect bonded atoms 5. Fill in non-bonding e pairs 6. nly if there is not enough e to go around, form double or triple bonds Write Lewis structures for: a l atom a l ion 4 N 3 X. Using Lewis dot structures to determine Electron Pair Geometry, Molecular Shape, and Polarity A. VSEPR stands for valence shell electron pair repulsion. Use this model to determine 3-D shapes of molecules 1. Based on idea that electron pairs are arranged to be as far apart as possible. 2. Three electron pair geometries (discussed in this class) a) linear (2 electron pairs)

9 b) trigonal planar (3 electron pairs) c) tetrahedral (4 electron pairs) 3. Molecular Shapes (determined by location of atom centers) a) linear electron pair geometry one possible molecular shape linear b) trigonal planar electron pair geometry two possible molecular shapes trigonal planar bent c) tetrahedral electron pair geometry three possible molecular shapes tetrahedral trigonal pyramidal bent 4. Try some examples a) 4 b) 3 l c) 2 d) 2

10 B. Polarity of molecules 1. Polarity: uneven distribution of charge. 2. Since p + must stay in nucleus, uneven charge distribution refers to e distribution around specific nuclei or the molecule as a whole. 3. You must look at e pair geometry (molecular symmetry) to determine molecular Polarity 4. Symmetry, vectors or a tractor pull? Appendix A: rbital hybridization A. Introduction 1. The s and p orbitals we used for isolated atoms does not work for molecules (ie.when atoms have formed covalent bonds). 2. A model that is useful for understanding these systems is called orbital hybridization. B. rbital hybridization orbital food processor 1. Put 2s orbital w/ appropriate # of 2p orbitals 2. Blend to obtain a set of hybrid orbitals. 3. Number of hybrid orbitals obtained equals number of s + number of p orbitals 4. Unlike p orbitals that have two equal lobes per orbital, sp, sp 2, and sp 3 hybrid orbitals have one main lobe (but see Dr. Winter s rbitron at per orbital. 5. The geometries of the hybrid orbitals match the patterns predicted from VSEPR theory. 6. The 2 lobed pi bonds that are part of double & triple bonds accurately predict the limited rotation observed w/ these types of bonds. sp 3 hybridization Formation of 4 sp 3 hybrid orbitals from 1 s and 3 p orbitals energy 2p orbital hybridization sp 3 2s geometry s p x p y p z orbital hybridization Note: Each p orbital has two lobes. Note: Each sp 3 hybrid orbital (in blue) has 1 main lobe.

11 Molecules that contain sp 3 hybridized Period 2 atoms, and N 4, methane N 3, ammonia N sp 2 hybridization` Atoms involved in forming one double bond exhibit sp 2 orbital hybridization patterns: Formation of 3 sp 2 hybrid orbitals from 1 s and 2 p orbitals energy 2s 2p orbital hybridization sp 2 2p z geometry s p x p y p z orbital hybridization Note: ombine one s and two p orbitals. Note: Each sp 2 hybrid orbital (shown in green) has one main lobe. p z orbital is unchanged. Molecules that contain sp 2 hybridized Period 2 atoms 2, formaldehyde (a.k.a., ethanal) 2, elemental oxygen Although you might conclude from the structures shown above that the two bonds that make up a double bond are identical, that is not the case. (See below.)

12 rbital hybridization: Building 2 from sp 2 hybridized 2 and s p x p y p z Note: ombine one 2s with two 2p orbitals to obtain three sp 2 hybrid orbitals. Trigonal planar geometry gives maximum sp 2 orbital separation. Note: Each sp 2 hybrid orbital (shown in green) has one main lobe. The p z orbital is unchanged. Formaldehyde, 2. Imagine pushing an sp 2 hybridized 2 group and an sp 2 hybridized atom together. p p p p ybrid orbitals are used for sigma (F) bonds (black) & non-bonding e - pairs (green). Unhybridized orbitals are used for pi (B) bonds (red). Note that the B bond has two lobes. Rotation about - axis??? B B ne bond is formed from the direct overlap of sp 2 orbitals, a sigma bond. The other is formed from sideways overlap of p orbitals, a pi bond. When a double bond is formed, rotation about that bond is eliminated. It is the pi bond components of double (and triple) bonds that limit rotation about these bonds. ontrast this with the free rotation about the - single bond in ethane.

13 The limited rotation in double bonds is responsible for another group of cis/trans isomers. The structures of the cis and trans isomers of 2-butene are shown below. (The -ene ending indicates that there is a double bond, and the 2" indicates the position of the double bond.) cis trans cis-2-butene trans-2-butene In general terms, a compound is cis if its higher ranking groups are on the same side of a line running through the atoms involved in the double bond, and trans if they are on opposite sides of the line. Biological relevance of cis & trans is- to trans-retinal in Rhodopsin, The 1st biochemical step in the visual process cis- double bond light psin (big protein) trans- double bond psin (big protein)