# 8/19/2011. Periodic Trends and Lewis Dot Structures. Review PERIODIC Table

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1 Periodic Trends and Lewis Dot Structures Chapter 11 Review PERIODIC Table Recall, Mendeleev and Meyer organized the ordering the periodic table based on a combination of three components: 1. Atomic Number 2. Relative Mass 3. Chemical and Physical Properties Periodic Trends PERIODIC Table Additionally, from the periodic table, the design is dictated based on the elements properties. We will focus on three key trends that can be derived from the periodic table: 1. Metallic Character 2. Atomic Radii 3. Ionization Energy Using a combination of these trends, one can predict the substances new properties and reactions it undergoes. 1

2 Metallic Character BREAKDOWN Metals Metals tend to lose electrons to become cations. Nonmetals Nonmetals tend to lose electrons to become anions. REACTION: M M n+ + ne - REACTION: M + ne - M n- Metallic Character BREAKDOWN (Properties) Metals 1. Shiny in appearance. 2. Ductile (can form wires). 3. Mallable (can be rolled or hammered into sheets) 4. High densities. 5. High melting points. Nonmetals 1. Nonshiny in appearance. 2. Brittle 3. Poor conductors of heat and electricity. 4. Low densities. 5. Low boiling points. METALLIC Character (ILLUSTRATION) 2

3 METALLIC Character The two key trends for this concept are the following: 1. From left to right on a period, the metallic character decreases. 2. From top to bottom within a family, the metallic character increases. Question Which of the following has the greater metallic character? a) Si vs. S Si b) P vs. As As c) P vs. Se Not enough information!!! ATOMIC Radius This is a property where we deal with the relationship between the total number of electrons in an atom and its effect on the radius size of the atom. THERE are two key trends: 1. If you move from left to right in a period, the atomic radii decreases. 2. If you move from top to bottom in a group, the atomic radii increases. 3

4 ATOMIC Radius (ILLUSTRATION) Question Which of the following has the largest radius? a) O vs. Cl Not enough information!! b) Sn vs. As Sn c) Ba vs. Sr Ba IONIZATION Energy This a property which deals with how easily it is to remove an electron from the atom in the gaseous state. REACTION: Na (g) + Ionization Energy Na + (g) + e - 4

5 IONIZATION Energy (cont d) REACTION: Na (g) + Ionization Energy Na + (g) + e - One can use the electron configuration to justify this behavior: Na: 1s 2 2s 2 2p 6 3s 1 Na + : 1s 2 2s 2 2p 6 The first ionization energy is the amount of energy required to remove the first electron from an atom. He + first ionization energy He + + e - He + 2,372 kj/mol He + + e - The second ionization energy is the amount of energy required to remove the second electron from an atom. He + + second ionization energy He 2+ + e - He + + 5,247 kj/mol He 2+ + e - As each succeeding electron is removed from an atom, ever higher energies are required. 5

6 IONIZATION Energy (cont d) There are two trends which must be known about ionization energy: 1. If you across a period from left to right on the periodic table, the ionization energy increases. 2. If you move down from top to bottom (in a family) on the periodic table, the ionization energy decreases. IONIZATION Energy (Illustration) Question Which of the following has the greater ionization energy? a) Ca vs. Rb Ca b) Ti vs. Nb Not enough information c) Cu vs. Ag Cu 6

7 After sodium loses its 3s electron, it has attained the same electronic structure as neon. 20 After chlorine gains a 3p electron, it has attained the same electronic structure as argon. 21 7

8 A The sodium 3s electron ion (Na+) of sodium and a chloride transfers ion to (Cl the - ) 3p are orbital formed. of chlorine. The force holding Na + and Cl - together is an ionic bond. Lewis representation of sodium chloride formation. 22 In NaCl the crystal is made each up sodium of cubic ion crystals. is surrounded by six chloride ions. In the crystal each chloride ion is surrounded by six sodium ions. 8

9 The ratio of Na + to Cl - is 1:1 There is no molecule of NaCl A sodium ion is smaller than a sodium atom because: (1) the sodium atom has lost its outermost electron. (2) The 10 remaining electrons are now attracted by 11 protons and are drawn closer to the nucleus. A chloride ion is larger than a chlorine atom because: (1) the chlorine atom has gained an electron and now has 18 electrons and 17 protons. (2) The nuclear attraction on each electron has decreased, allowing the chlorine to expand. 9

10 Prediction IONIC Formulas From the main group elements, if you know where the element belongs, one can predict the expected charge when the element transforms into its respective ion. RECALL the following: Metals lose electrons to become cations. Nonmetals gain electrons to become anions. Charges MAIN Group Elements Group Charges IA +1 IIA +2 IIIA +3 IVA +4/-4 VA +5/-3 VIA +6/-2 VIIA +7/-1 VIIIA 0 Examples Predict the chemical formula for magnesium phosphide. ANSWER: Magnesium is in group IIA: Mg 2+ Phosphide is an anion from group VA: P 3- FINAL Answer: Mg 3 P 2 10

11 Additional Practice Write the chemical formulas for the following compounds. a. Lithium Sulfide b. Potassium Bromide c. Aluminum Chloride Covalent Bond What is a covalent bond? Essentially, a covalent bond is one where there is a chemical bonding occurring between two nonmetal elements. The electrons are being shared, not transferred between electrons. Let s look at some examples!!! 11

12 Substances which covalently bond exist as molecules. Carbon dioxide bonds covalently. It exists as individually bonded covalent molecules containing one carbon and two oxygen atoms. The term molecule is not used when referring to ionic substances. Sodium chloride bonds ionically. It consists of a large aggregate of positive and negative ions. No molecules of NaCl exist. Covalent bonding in the hydrogen molecule The most likely The region orbital to find of the the Two 1s orbitals from each of electrons two electrons includes is two hydrogen atoms overlap. both between hydrogen the two nuclei. nuclei. Each 1s orbital contains 1 The two nuclei are electron. shielded from each other by the electron pair. This allows the two nuclei to draw close together. 12

13 Covalent bonding in The the orbital chlorine of the molecule electrons includes Two 3p orbitals from both each chlorine of two chlorine atoms nuclei. overlap. The two nuclei are shielded from each other by the electron pair. This The most likely allows the two region to find the nuclei to draw two electrons is close together. between the two nuclei. Each unpaired Each 3p chlorine orbital now has 8 on each chlorine electrons atom in its outermost contains 1 electron. energy level. Covalent bonding with equal sharing of electrons occurs in diatomic molecules formed from one element. hydrogen chlorine iodine nitrogen A dash may replace a pair of dots. Electronegativity 13

14 What is electronegativity? It is the attractive force in which an atom has for shared electrons in a molecule or a polyatomic ion. Let s look at HCl!!! Partial positivepartial chargenegative charge on hydrogen. on chlorine. Polar Covalent Bonding in HCl δ+ : δ- H Cl : The attractive force that an atom of an element has for shared electrons in a molecule or a Chlorine has a greater attraction Shared The shared electron for the electron pair. pair shared electron pair than is hydrogen. closer to chlorine than to hydrogen. polyatomic ion is known as its electronegativity. 41 The discovery of determine the relative electronegativities of these elements were determined by Linus Pauling. 14

15 Bond Types 1. Pure Covalent (NONPOLAR Covalent) the difference between the two atoms electronegativities falls between Polar Covalent the difference between the two atoms electronegativities falls between Ionic the difference between the two atoms electronegativities is greater than 2.0. If the electronegativities are the same, the bond is nonpolar covalent and the electrons are shared equally. The molecule is nonpolar covalent. Cl Cl Electronegativity Difference = 0.0 Electronegativity 3.0 Electronegativity 3.0 Chlorine Molecule If the electronegativities are not the same, the bond is polar covalent and the electrons are shared unequally. The molecule is polar covalent. δ+ δ- H Cl Electronegativity Difference = 0.9 Electronegativity 2.1 Electronegativity 3.0 Hydrogen Chloride Molecule 15

16 If the electronegativities are very different, the bond is ionic and the electrons are transferred to the more electronegative atom. Electronegativity Difference = 2.1 No molecule exists. The bond is ionic. Na + Cl - Electronegativity 0.9 Electronegativity 3.0 Sodium Chloride Application ELECTRONEGATIVIES In cases where there is an asymmetrical molecule with a difference in electronegativity and of opposite charges, then a dipole is present. REPRESENTATION: Additionally an arrow is used to indicate the more electronegative element. An arrow can be used to indicate a dipole. Thearrowpointsto thenegativeendofthe dipole. MoleculesofHCl,HBrandH 2 Oarepolar. O H Cl H Br H H 16

17 Lewis Dot Structures Prediction Covalent Compounds When two or more nonmetals interact to form its respective molecular compound, a covalent bondis produced. In a covalent bond, recall that electrons are being shared between atoms in forming the chemical bond between them. Valence Electrons It is important to know how to determine the valence electrons for nonmetals by looking at either its electron configuration or from the group it is located. Example: O or Oxygen It has an electronic configuration of the following: 1s 2 2s 2 2p 4. How does one determine the exact number of valence electrons? 17

18 Example: O or Oxygen OXYGEN Electronic Configuration: 1s 2 2s 2 2p 4. Look at the highest principle shell where the electrons are located furthest from the nucleus. If successful, O has 6 valence electrons. What if we didn t know that, can it be determined by its main group location? Valence Electrons (BREAKDOWN) Atom Group Valence Electrons Cl 7A 7 H 1A 1 C 4A 4 O 6A 6 N 5A 5 S 6A 6 P 5A 5 I 7A 7 Writing Lewis Structures To write correct a Lewis structure for a covalent compound, follow these steps: 1. Write the correct skeletal structure for the molecule. 2. Calculate the total number of electrons for the Lewis structure by summing the valence electrons of each atom in the molecule. 3. Distribute the electrons among the atoms, giving octets (or duets for hydrogen) to as many atoms as possible. 4. If an atoms lack an octet, form double or triple bonds as necessary to give them octets. (ADDITIONAL Tip: You must find the one who has the most metallic character which will be the central atom.) 18

19 Example Write a Lewis structure of CO Write the correct skeletal for the molecule. O C O 2. Calculate the total number of electrons for the Lewis structure by summing the valence electrons of each atom in the molecule. TOTAL Electrons= # valence e - for C + 2 (# valence e - for O) = 4 + 2(6) = 16 Example Write a Lewis structure of CO Distribute the electrons among the atoms, giving octets (or duets for hydrogen) to as many atoms as possible. O C O 4. If any atoms lack an octet, form double or triple bonds as necessary to give them octets. O C O O C O Polyatomic Ions Similarly, one can write out the Lewis structures for ones which has a formal charge. EXAMPLE: CN - 1. Write the correct skeletal structure for the structure. C N 2. Calculate the total number of electrons for the Lewis structure. TOTAL Electrons= # valence e - for C + # valence e - for N -charge = (-1) = 10 19

20 Polyatomic Ions Similarly, one can write out the Lewis structures for ones which has a formal charge. EXAMPLE: CN - 3. Distribute the electrons among the atoms, giving octets (or duets for hydrogen) to as many atoms as possible. C N 4. If any atoms lack an octet, form double or triple bonds as necessary to give them octets. C N C N What is resonance? Resonance Essentially, resonance is a term used to explain that there are multiple structure for a compound or ion in which they have the same number of total valence electrons. Let s look at first SO 2. Analysis (SO 2 ) First, count the # of valence electrons. Total e - =6 + 2(6)=18 valence e - Sulfur s Valence Electrons Oxygen s Valence Electrons 20

21 Analysis (SO 2 ) (cont d) Second, start drawing first the shared electrons. O S O Third, draw in the non-bonded electrons. O S O Fourth, a lone pair from oxygen can be used to form a double bond with sulfur. Example (SO 2 ) O S O O S O Check Structure (SO 2 ) O S O One can also check if this structure is valid by using this table: Atom Max e Used e Diff. E S O O S O O Total electrons=3 shared pairs of bonds 21

22 Geometries GEOMETRIES Once the electronic structure of a molecular compound has been drawn out, then one can determine its geometry. The geometries are in respects by Lewis theory in combination with VSEPR theory. VSEPR= Valence Shell Electron Pair Repulsion Let s illustrate this theory in action!!! Carbon Dioxide (CO 2 ) If you look at this structure, there are two double bonds between C and O atoms. There are 16 total valence electrons. There are two atoms paired to the center atom. GEOMETRY: Linear (Molecular and Electronic) 22

23 Formaldehyde (CH 2 O) If you look at this structure, there is a double bond between C and O + 2 single bonds between C and H atoms. There are 12 total valence electrons. There are three atoms paired to the center atom. GEOMETRY: Trigonal Planar (Molecular and Electronic) Methane (CH 4 ) If you look at this structure, there are four single bonds between C and H atoms. There are 8 total valence electrons. There are four atoms paired to the center atom. GEOMETRY: Tetrahedral (Molecular and Electronic) Sulfur Dioxide (SO 2 ) If you look at this structure, there is one single bond and a double bond between S and O atoms + a lone pair of electrons. There are 18 total valence electrons. There are two atoms paired to the center atom. GEOMETRY: Trigonal Planar (Electronic) Bent (Molecular) 23

24 Ammonia (NH 3 ) If you look at this structure, there are two double bonds between N and H atoms plus 1 lone pair of electrons. There are 8 total valence electrons. There are three atoms paired to the center atom. GEOMETRY: Tetrahedral (Electronic) and Trigonal Pyramidal (Molecular) Water (H 2 O) If you look at this structure, there are two single bonds between H and O atoms + 2 lone pairs of electrons. There are 8 total valence electrons. There are two atoms paired to the center atom. GEOMETRY: Tetrahedral (Electronic) & Bent (Molecular) Electron Groups ELECTRON and MOLECULAR Geometries Bonding Groups Lone Pairs Electron Geometry Angle between Electron Groups Molecular Geometry Linear 180 o Linear Example (Electron) (Ball and Stick) Trigonal Planar Trigonal Planar 120 o Trigonal Planar 120 o Bent Tetrahedral o Tetrahedral Tetrahedral o Trigonal Pyramidal Tetrahedral o Bent 24

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