metals, transition metals, halogens, noble gases, lanthanides, actinides, and hydrogen.

Transcription

1 Chemistry HP Unit 3 The Periodic Table Learning Targets (Your exam at the end of Unit 2 will assess the following:) 3. The Periodic Table 3-1. Discuss the development of the periodic table by Mendeleev Locate and state important properties of main chemical families including the alkali metals, alkaline earth metals, transition metals, halogens, noble gases, lanthanides, actinides, and hydrogen Define atomic radius explain periodic trends in this property as they relate to atomic structure Define ionization energy and explain periodic trends in this property as they relate to atomic structure. List elements that are exceptions to the general trend and use orbital notation to show why they are exceptions Define electronegativity and explain periodic trends in this property as they relate to atomic structure Define ionic radius and relate the size of an anion to a neutral atom of the same element and a cation to a neutral atom of the same element Draw electron dot diagrams for atoms, showing the correct number of valence electrons Draw Lewis structures from chemical formulas Assign bond orders for a molecule from the Lewis structure Calculate the total number of valence electrons in a polyatomic ion Draw Lewis structures for polyatomic ions Assign formal charges to atoms in polyatomic ions Draw resonance structures for polyatomic ions Classify bonds as ionic or covalent and as polar or non-polar using electronegativity values Assign shapes to molecules using VSEPR Theory and draw the VSEPR diagrams for a molecule Classify molecules as polar or non-polar using shape Compare miscible and immiscible, by definition and with example, and determine if two substances will be miscible or immiscible based on polarity 1

2 3-1. Discuss the development of the periodic table by Mendeleev. Mendeleev s Periodic Table Dmitri Mendeleev was a Russian chemistry professor who was looking for a way to the to aid his students. He listed the elements in according to their, and in when of the elements began to. Mendeleev left in his table for elements that had yet to be discovered. When the elements and were later discovered, they fit perfectly into the gaps Mendeleev had left for them. Mendeleev also believed that the placement of elements and should be switched, to correspond with their chemical properties. When the Modern Periodic Table was later organized according to and not, the elements and lined up correctly according to their chemical properties. Mendeleev Periodic Table CLOZE: Mendeleev organized elements by, NOT ( had not been discovered yet), AND organized them by similar. He left for elements which predicted the of those elements, which were discovered later. The Modern Periodic Table The Modern Periodic Table organizes elements into rows, called, according to, and NOT. Each period corresponds to the, or,. The elements are also organized into according to. The columns are called. Each column, corresponds to a different of elements 2

3 3-2. Locate and state important properties of main chemical families including the alkali metals, alkaline earth metals, transition metals, halogens, noble gases, lanthanides, actinides, and hydrogen. Chemical Groups/Families Elements in the same chemical family share similar properties. Some chemical families have been given names. Group Number and Name Hydrogen Elements Ion(s) Formed Properties I or 1 Alkali Metals II or 2 Alkaline Earth Metals 3-12 Transition Metals VII or 17 Halogens VIII or 18 Nobel Gases Lanthanides (57-70) and Actinides (89-102) 3

4 WS #1 (Learning Target 3-2: Locate and state important properties of main chemical families including the alkali metals, alkaline earth metals, transition metals, halogens, noble gases, lanthanides, actinides, and hydrogen.) Periodic Table Puzzle Fictitious symbols are used for the first 18 elements in the periodic table. Use the clues below to write the fictitious symbol in the appropriate spot on the periodic table provided. Symbols for real elements do not represent those elements. HINT: You do not have to complete each clue in order. Clue 1 U and J are alkali metals. J has more energy levels. Clue 2 T has 4 valence electrons on the 3rd energy level. Clue 3 M is a metal in period 3 with 2 valence electrons. Clue 4 X has one proton in its nucleus. Clue 5 Q has 2 energy levels, is a nonmetal, and is a solid at room temperature. Clue 6 L is a noble gas that doesn t have 8 valence electrons. Clue 7 Z and Y are members of the nitrogen family. Y is a gas at room temperature. Clue 8 D has an ending electron distribution of s 2 p 5. R has an ending electron distribution of s 2. Clue 9 Clue 10 Clue 11 G has 6 valence electrons. V and W have full outer energy levels. V has 3 energy levels. A atoms have 3 valence electrons and E atoms have 6 valence electrons. Both are in the second period. Clue 12 K has one fewer total electrons than V. Clue 13 I has 3 valence electrons on the third energy level. Answers: X L U-R-A-Q-Y-E-D-W J-M-I-T-Z-G-K-V 4

5 3-3. Define atomic radius explain periodic trends in this property as they relate to atomic structure. Periodic Trends Some chemical properties of elements vary predictably down a group and across a row on the periodic table. These tendencies are called periodic trends and are based on atomic structure. Periodic trends are observed in the properties of atomic radius, ionization energy, and electronegativity. Trend Definition Atomic Radius Atomic radius is the of an atom. The atomic radius is measured as from the nucleus to the. Atomic radius is usually measure in (1 pm = m). Atoms range in size from (Helium) to (Francium). Down a Group Across a Period Example Atomic Radius down a group. Atoms become larger down a group of the periodic table. Each new element down a group adds a. Since electrons are located in, they are from the nucleus, making atoms. Atomic Radius across a row. Atoms become smaller across a row of the periodic table. Across a row, new energy levels are added, but the number of in the increases. Therefore, the positive charge on the nucleus becomes and electron shells are to the nucleus, making atoms. Which atom has a larger atomic radius? Why? (1) Na or K? (2) C or O? Atomic Radius CLOZE Atomic radius is defined as the from the nucleus to the outermost electron in an atom. It is usually measured in the unit. As you go down a on the Periodic Table, ionic radius, because, with each new period, an additional is added, making the atomic radius. Therefore, we would expect the atomic radius of Br to be than that of Cl. As you go across a period on the Periodic Table, atomic radius. This is because the number of protons as you go left to right, while the energy level. The increased number of creates a greater force between the nucleus and its electrons, bringing the outermost electrons in, and making the atomic radius. Therefore, we would expect the atomic radius of Br to be than that of Se. 5

6 3-4. Define ionization energy and explain periodic trends in this property as they relate to atomic structure. List elements that are exceptions to the general trend and use orbital notation to show why they are exceptions. Trend Definition Ionization Energy Ionization energy is the required to remove the from an atom. For example, the ionization for an atom of lithium: Li + IE Li + + e - The ionization for an atom of fluorine: F + IE F + + e - Generally, the closer an electron is to the nucleus, the it is to remove and the the ionization energy will be. Down a Group Ionization Energy down a group. Each new element down a group adds a, so the outermost electrons are located in shells from the nucleus, making the outermost electrons to remove. Across a Period Ionization Energy across a row. Across a row, are added, but the number of protons in the nucleus. Therefore, the electrons are pulled to the nucleus and are to remove, the amount of to remove the electrons, the ionization energy. Example Exceptions Example Which atom has a higher IE? Why? (1) Na or K? (2) C or O? Ionization Energy and Ion Stability Ionization energy can also depend on the of the by losing the electron. In general, ions are most stable when the highest energy orbitals are or -. The of an atom must be considered to explain some trends. Which atom has a higher IE? Why? (1) Al or Mg? We would expect Al to have a ionization energy than Mg because its outermost electrons are in the, and Mg has protons than Al. However, Mg actually has a ionization energy than Al because: 6

7 When Mg loses an electron, Mg Mg + + e -, its electron configuration changes from 1s 2 2s 2 2p 6 3s 2 to 1s 2 2s 2 2p 6 3s 1 + e -. In doing so, it goes from a 3s orbital to an 3s orbital. On the other hand, when Al loses an electron, Al Al + + e -, its electron configuration changes from 1s 2 2s 2 2p 6 3s 2 3p 1 to 1s 2 2s 2 2p 6 3s 2 + e -. In doing so, Al goes from an 3p orbital to a 3s orbital. Therefore, the ionization energy of Al is than the ionization energy of Mg. (2) S or P? Ionization Energy CLOZE Ionization Energy is defined as the required to remove the outermost electron from an atom. As you go down a on the Periodic Table, ionization energy, because, with each new period, an additional is added, bringing the outermost electron the nucleus. Therefore, we would expect the ionization energy of Br to be than that of Cl. As you go across a period on the Periodic Table, ionization energy. This is because the number of protons as you go left to right, while the energy level. The increased number of creates a greater force between the nucleus and its electrons, making it to remove an electron. Therefore, we would expect the ionization energy of Br to be than that of Se. An exception to this trend is when the ion that is formed by removing the outermost electron is more stable than the neutral atom. When this is the case, ionization energy across a period may not always from one element to the next. Boron and Beryllium are examples of this. One would expect the ionization energy of boron to be than that of beryllium. But, since removing one electron from boron results in a filled sublevel, whereas removing one electron from beryllium removes an electron from a filled sublevel, the ionization energy of boron is than that of beryllium. 7

8 3-5. Define electronegativity and explain periodic trends in this property as they relate to atomic structure. Trend Definition Electronegativity Electronegativity is the measure of the ability of an atom in a chemical compound to the electrons of the other atom in a. The closer the electrons of another atom in a bond can be to the, the the electronegativity. In general, atoms have a electronegativity and atoms have a electronegativity. Electronegativity values are measured on a scale from to. Down a Group Electronegativity down a group. Down a group, atoms become, so the electrons of another atom in a bond would be from the nucleus and electronegativity is. 8

9 Across a Period Electronegativity across a row. Across a row, atoms become, so electrons of another atom in a bond can be to the nucleus and electronegativity is. Example Which atom has a higher EN? Why? (1) Na or K? (2) C or O? Electronegativity CLOZE Electronegativity is defined as the ability of an atom in a chemical bond to attract the other atom s. As you go down a on the Periodic Table, electronegativity, because, with each new period, an additional is added, increasing the size of the atom. Since the larger the atom, the the nucleus is from its outermost electrons, and the the nucleus is from the electrons of the other atom in a chemical bond, it becomes more difficult to attract the other atom s electrons toward itself. Therefore, we would expect the electronegativity of Br to be than that of Cl. As you go across a period on the Periodic Table, electronegativity. This is because the number of protons as you go left to right, while the energy level. The increased number of creates a greater force between the nucleus and its electrons, and between the nucleus and the electrons of the other atom in a chemical bond, making it to attract an electron. Therefore, we would expect the electronegativity of Br to be than that of Se. In general, noble gases have an electronegativity of, because they usually do not form chemical bonds. 9

10 3-6. Define ionic radius and relate the size of an anion to a neutral atom of the same element and a cation to a neutral atom of the same element. Trend Definition Cations Anions Example Ionic Radius Ionic radius is the of the ion of an element compared to the neutral atom of the element. Positive Ions/Cations Atoms become positively charged ions by electrons. Since electrons are lost, a positive ion is than the neutral atom. Negative Ions/Anions Atoms become negatively charged ions by electrons. Since electrons are gained, a negative ion is than the neutral atom Which has a larger radius, the neutral atom or the ion? Why? (1) Ca or Ca 2+? (2) S or S 2-? Ionic Radius CLOZE Ionic radius is defined as the from the nucleus to the outermost electron in an ion. When an atom an electron, becoming an anion, its ionic radius. Therefore, we would expect the radius of Br- to be than the atomic radius of neutral Br. When an atom an electron, becoming a cation, its ionic radius. Therefore, we would expect the radius of Na + to be than the atomic radius of neutral Na. 10

11 WS #2 (Learning Targets 3-3, 3-4, 3-5, 3-6) Periodic Trends (1) List each set of atoms from smallest to largest radius. (a) Mg, Be, Ca (c) As, P, N (b) C, B, N (d) Sc, Ti, V (2) Which of the elements listed has the highest ionization energy? (a) K, Li, Cs (c) Na, Mg, Cl (b) C, F, Be (d) Ra, Ca, Be (3) Which of the elements listed has the highest electronegativity? (a) Be, B, N (c) I, Br, At (b) Se, O, S (d) Na, Al, P (4) Complete the following table: Give Electron Configuration (a) Si or Cl? (b) F or Br? Si: Cl: F: Br: Which element has a larger atomic radius? Explain. Which element has a higher ionization energy? Explain. Which element has a higher electronegativity? Explain. (5) Which has a larger radius, the neutral atom or the ion? Explain. (a) Br or Br (b) Cs or Cs + 11

12 12

13 3-14. Classify bonds as ionic or covalent and as polar or non-polar using electronegativity values. Atoms react with each other and form bonds in order to with electrons. This is why noble gases generally do not bond with other elements; they already have full outer shells. The types of bonds they form depend upon the difference in electronegativity between the two atoms forming the bond. Recall that electronegativity is a measure of the ability of an atom in a chemical compound to in a chemical bond. A bond in this sense then, is a tug of war between electrons. The greater the electronegativity difference between the two atoms forming a bond, the more the electrons will around the atom. As an example, let s consider the bond sodium and chlorine form when making sodium chloride, also known as table salt. Both sodium and chlorine are in period of the periodic table. Sodium is in the first group, the, whereas chlorine is a, in the second to last group. According to the periodic trends, we would expect chlorine to have a electronegativity than sodium, because chlorine has a nuclear charge (it has more in its nucleus), while both sodium s and chlorine s outermost electrons are distances from the nucleus, being in the same. We can confirm this by comparing the electronegativity values of sodium and chlorine. From the chart, you can see that the electronegativity of chlorine is, and the electronegativity of sodium is. That s an electronegativity difference of - =. This electronegativity difference is so, that chlorine literally sodium s valence electron, forming the, leaving sodium as the ion. The bond that sodium and chlorine form is therefore called an, because the electronegativity difference between the elements is so great that the valence electrons are literally from the electronegative atom to the electronegative atom, turning each atom into an. An ionic bond can therefore be defined as a bond between, caused by an electronegativity difference of or greater. What holds these ions together in an ionic bond is the between the cation ( ion) and the anion ( ion). The rule of thumb is that an ionic bond forms when a bonds with a. This makes sense when you consider that metals are on the side of the periodic table and nonmetals are on the side. Nonmetals have a electronegativity then metals, because they generally have a nuclear charge. Atoms whose electronegativity difference lies between and are called bonds. The term covalent is made up of the prefix co- which is a, and -valent which refers to the. Covalent bonds are therefore a of electrons. Neither of the two atoms in a covalent bond have sufficiently more electronegativity than the other to literally take the other s valence electrons. Rather, in a covalent bond, the valence electrons are. An example of a covalent bond is the bond between oxygen and hydrogen in the molecule water. Hydrogen has an electronegativity of and oxygen has an electronegativity of. The electronegativity difference between these two atoms is - =. Since is than 1.7, oxygen and hydrogen form a bond. Because oxygen has a electronegativity than hydrogen, the electrons will more around the oxygen, giving the oxygen an excess charge, and leaving the hydrogen with an excess charge. But, because the electronegativity difference is high enough, the electrons will completely hydrogen to join oxygen. 13

14 3-14. Classify bonds as ionic or covalent and as polar or non-polar using electronegativity values. A covalent bond can therefore be defined as a bond between two atoms in which the electronegativity difference of is less than or equal to. The rule of thumb is that a covalent bond forms when bond with. This makes sense when you consider that all the nonmetals are on the side of the periodic table, so their electronegativity values are to each other. Covalent compounds can be further classified as or. Polar comes from the word, describing the existence of positive and negative regions, or within a molecule. In the last example comparing the electronegativities of hydrogen and oxygen, we saw that the electronegativity difference between hydrogen and oxygen was, which is less than the difference required for ionic compounds. However, since there a difference in electronegativity, the electrons in the compound will be attracted to the atom more than the atom. This will leave the oxygen atom with a charge, also called a, and the hydrogen atom with a charge because of its of electron density, also called a. We draw these poles using the following symbols: The bond between oxygen and hydrogen in water is referred to as a, because, even though the bond is covalent, it does contains positive and negative. If the electronegativity difference between two atoms in a covalent compound is less than or equal to, we refer to the bond as. If the electronegativity difference is between and, we refer to the bond as. The following chart summarizes these electronegativity differences: Electronegativity Difference Covalent Ionic Nonpolar Covalent Polar Covalent 0 to to and up 14

15 3-14. Classify bonds as ionic or covalent and as polar or non-polar using electronegativity values. To Read Before the Next Lesson: The type of bond shared between elements in a compound determines some of the properties of that compound. Ionic compounds, formed when ions bond together, exist as crystal structures with repeating cation-anion arrangements, such as the example drawn below of NaCl. The large green spheres are the sodium ions, and the smaller blue spheres are the chloride ions. They assemble into this crystal lattice structure to maximize electrostatic attraction between the cations and anions, and to minimize the electrostatic repulsion between like charges. Here is a 3D image of the same structure: Since a sodium ion is bonded within an entire crystal of repeating positive and negative ions, we don t refer to NaCl as a molecule, but rather a formula unit. NaCl is the formula unit of the repeating Na-Cl-Na-Cl-Na-Cl crystalline structure. Because of this strong electrostatic attraction between the positive and negative ions, the melting points of ionic compounds are high. It takes a lot of energy to loosen the hold these ions have on each other. In addition, ionic compounds are brittle. And, they act as electrolytes, conducting electricity in solution. In summary, the properties of ionic compounds are as follows: 1. They exist as repeating units within a crystalline structure. 2. High melting points. 3. Brittle 4. Strong electrolytes (conduct electricity when in solution). Covalent compounds are the result of nonmetals bonding with each other and sharing their electrons in such a way that every atom in the bond has a full set of valence electrons. Covalent compounds are referred to as molecules, not formula units. In general, covalent compounds have: 1. Lower melting and boiling points. 2. Are non-electrolytes (They do NOT conduct electricity). 3. Exist as mostly liquids or gases at room temperature. 15

16 WS #3 (Learning Target Classify bonds as ionic or covalent and as polar or non-polar using electronegativity values.) Directions: Classify the following bonds as ionic, polar or nonpolar covalent. Draw the bond. If the bond is polar covalent, indicate the distribution of charge using the arrow and delta symbols. Compound Electronegativity Difference Type of Bond Draw Bond (a) Br 2 (b) MgO (c) LiF (d) SeO (e) ICl (f) BrCl (g) CO (h) NaCl (i) HCl (j) O 2 16

17 Answers. Compound Electronegativity Difference Type of Bond Draw Bond (a) Br 2 0 (b) MgO (c) LiF (d) SeO (e) ICl (f) BrCl (g) CO (h) NaCl (i) HCl (j) O 2 17

18 3-7. Draw electron dot diagrams for atoms, showing the correct number of valence electrons. We represent both ionic and covalent bonds using a notation known as Lewis Dot Structures, which are a way of representing the valence electrons around an atom. How to write Lewis Dot Structures (electron dot diagrams) 1. Write the chemical symbol of the element. 2. Put one dot for each valence electron. 3. Don t pair electrons until you have to. Element Number of Valence Electrons Lewis Dot Structure Ion Number of Valence Electrons of Ion Lewis Dot Structure of Ion N Cl Al Mg He Na Lewis Dot Structures of Ionic Compounds Draw the Lewis Dot Structures of the Ions making up the ionic compound next to each other. Ionic Compound NaCl Lewis Dot Structure of Cation Lewis Dot Structure of Anion Lewis Dot Structure of Ionic Compound MgO BeS MgCl2 18

19 WS #4 (Learning Target 3.7: Draw electron dot diagrams for atoms, showing the correct number of valence electrons. ) Directions: Draw the Lewis Dot Structures for the following atoms, ions and ionic compounds. Atom, Ion or Ionic Compound 1) Ca 2+ Lewis Dot Structure 2) I - 3) Ge 4) As 5) Rb + 6) RbI 7) AlCl3 8) CaO 9) LiBr 10) Na2S 19

20 3-8. Draw Lewis structures from chemical formulas Assign bond orders for a molecule from the Lewis structure To Draw Lewis Dot Structures for Covalent Compounds, follow these 5 steps: Example: Draw the Lewis Dot Structure for CF4 What is the bond order for each C-F bond? Example: Draw the Lewis Dot Structure for H2O What is the bond order for each O-H bond? 20

21 3-8. Draw Lewis structures from chemical formulas Assign bond orders for a molecule from the Lewis structure Example: Draw the Lewis Dot Structure for CO2 What is the bond order for each C-O bond? Example: Draw the Lewis Dot Structure for N2 What is the bond order for the N-N bond? Exceptions: Incomplete Octet Examples: Expanded Valence Examples: 21

22 WS #5 (Learning Target 3-8, 3-9) Directions: Draw Lewis structures for the following compounds. (a) I 2 (b) CF 4 (c) NCl 3 (d) SCl 2 (e) SiTe 2 (f) PF 3 (g) GeBr 4 (h) TeCl 2 (i) CSe 2 (j) H 2 (k) S 2 (l) P 2 (m) C 2 (n) C 2H 6 (o) C 2H 4 (p) C 2H 2 (q) HSCN (r) BBr 3 (s) PI 5 (t) SAt 6 22

23 3-10. Calculate the total number of valence electrons in a polyatomic ion Draw Lewis structures for polyatomic ions Assign formal charges to atoms in polyatomic ions Draw resonance structures for polyatomic ions. A polyatomic ion is a group of -- acting together as. Examples: When drawing the Lewis structure for a polyatomic ion, for each charge and for each charge. Example: Draw the Lewis Dot Structure for CN -. What is the bond order for the C-N bond? Formal Charge The charge on each atom in the compound can be determined by calculating the. The formal charges for all atoms must add up to the charge for the polyatomic ion. The formal charge on each atom in a polyatomic ion can be calculated as follows: Formal Charge = number of valence electrons number of unshared electrons ½ number of shared (bonding) electrons Formal Charge of CN - # Valence Electrons # Non-bonding Electrons ½ # Bonding Electrons Formal Charge Carbon Nitrogen 23

24 Example: Draw the Lewis Dot Structure for NO3 -. What is the bond order for each N-O bond? Formal Charge of NO3 - # Valence Electrons # Non-bonding Electrons ½ # Bonding Electrons Formal Charge Nitrogen Oxygen (1,2) Oxygen (3) Resonance Structures Resonance Structures = When 2 or more are necessary to describe the bonding in a molecule or ion. Example: Resonance Structures of NO3 -. What is the bond order for each N-O bond? When do you draw resonance structures? 1. When you have more than one possible Lewis Dot Structure. 2. Common in Polyatomic Ions containing oxygen. Bond Order of Resonance Structures = Total Number of Bonds Total Number of Bonding Locations. 24

25 WS #6 (Learning Targets 3-10, 3-11, 3-12, 3-13) (1) Determine the total number of valence electrons in each of the following polyatomic ions. Draw Lewis structures and calculate the formal charge on each atom. (a) OCl (b) OH (c) H 3O + (d) CO 3 2 (e) NO 2 (f) PO 3 3 (g) CN (h) SO 4 2 (2) Draw all possible resonance structures for the polyatomic ion. Calculate the formal charge on each atom. Determine the bond order. (a) SiO 3 2 (b) PO

26 Answers. (1) Determine the total number of valence electrons in each of the following polyatomic ions. Draw Lewis structures and calculate the formal charge on each atom. (a) OCl (b) OH (c) H 3O + (d) CO 3 2 (e) NO 2 (f) PO 3 3 (g) CN (h) SO 4 2 (2) Draw all possible resonance structures for the polyatomic ion. Calculate the formal charge on each atom. Determine the bond order. (a) SiO 3 2 (b) PO

27 3-15. Assign shapes to molecules using VSEPR Theory and draw the VSEPR diagrams for a molecule Classify molecules as polar or non-polar using shape Compare miscible and immiscible, by definition and with example, and determine if two substances will be miscible or immiscible based on polarity VSEPR Theory = Stands for. Because valence electrons each other, a molecule assumes a shape that keeps each valence electron pair as from each other as possible. Hence, VSEPR Theory predicts the of molecules. In order to determine the shape of a molecule, you need to know: 1. How many atoms are bonded to the central atom. 2. The number of lone pairs attached to the central atom. Lone pair = A pair of electrons. Steric Number = The number of bonded to the central atom, plus the number of. To Determine the Shape of a Molecule 1. Draw Lewis Dot Structure 2. Determine steric number and number of lone pairs of central atom. 3. Use VSEPR Table to identify shape. Polarity of Molecules The VSEPR shape can be used to determine the polarity of a molecule. A molecule with a three dimensional shape will be non-polar. Testing for Polarity Substances can be tested for their polarity by mixing them with liquids of known polarity. Generally, (two polar substances will be and two non-polar substances will be, but a polar and a non-polar substance will be ) Miscible: substances each other (ex. salt and water) Immiscible: substances do dissolve each other (ex. oil and water) 27

28 Steric Number General Shape # Lone Pairs Specific Shape Picture Polar or Nonpolar? 1 Linear 2 Linear Nonpolar 3 Trigonal Planar 0 Trigonal Planar Nonpolar 3 Trigonal Planar 1 Bent Polar 4 Tetrahedral 0 Tetrahedral Nonpolar 4 Tetrahedral 1 Trigonal Pyramidal Polar 4 Tetrahedral 2 Bent Polar 5 Trigonal Bipyramidal 0 Trigonal Bipyramidal Nonpolar 5 Trigonal Bipyramidal 1 Seesaw Polar 5 Trigonal Bipyramidal 2 T-shaped Polar 5 Trigonal Bipyramidal 3 Linear Nonpolar 6 Octahedral 0 Octahedral Nonpolar 6 Octahedral 1 Square Pyramidal Polar 6 Octahedral 2 Square Planar Nonpolar 28

29 Example: Determine the VSEPR shape and polarity of SF2. Example Molecule CO2 Lewis Structure Steric Number # lone pairs on central atom Shape VSEPR Diagram Polar or Nonpolar? BH3 CH4 NH3 H2O PCl5 SF6 29

30 WS #7 (Learning Targets 3-15, 3-16, 3-17) Complete the following table. Example Molecule Lewis Structure Steric Number # lone pairs on central atom VSEPR Diagram and Shape Polar or Nonpolar? SiI4 SeF2 PBr5 BAt3 SBr6 NF3 CS2 SF6 30

31 Answers. Molecule Lewis Structure Steric # lone pairs on VSEPR Diagram and Shape Polar or Nonpolar? Number central atom SiI4 4 0 Nonpolar SeF2 4 2 Polar PBr5 5 0 Nonpolar BAt3 3 0 Nonpolar SBr6 6 0 Nonpolar NF3 4 1 Polar CS2 2 0 Nonpolar 31

32 WS #9 (Unit 3 Review) 1. (Learning Target 3-1) Discuss the development of the periodic table by Mendeleev. What are the differences between Mendeleev s Periodic Table and the Modern Periodic Table? 2. (Learning Target 3-2) Locate and state important properties of main chemical families including the alkali metals, alkaline earth metals, transition metals, halogens, noble gases, lanthanides, actinides, and hydrogen. Identify the period and group that each of the following elements belongs to, the number of valence electrons its neutral atom has, and a property that all elements in that group share. Element Period Group Valence Electrons Property Ca F O Ar Fr Ag xxx U xxx 3. (Learning Target 3-3) Define atomic radius and use atomic structure to explain the periodic trend in atomic radius as you go down a group on the Periodic Table and across a period. 4. (Learning Target 3-4) Define ionization energy and use atomic structure to explain the periodic trend in ionization energy as you go down a group on the Periodic Table and across a period. List the two groups of elements that are exceptions to the general trend and use orbital notation to explain why. 5. (Learning Target 3-5) Define electronegativity and use atomic structure to explain the periodic trend in electronegativity as you go down a group on the Periodic Table and across a period. Which group of elements has zero electronegativity? 6. (Learning Target 3-6) Define ionic radius and compare the size of an anion to a neutral atom of the same element and a cation to a neutral atom of the same element. Which type of elements from anions? Which type of elements form cations? 7. (Learning Target 3-6) Which has the larger radius, the neutral atom or the ion? Explain. a. O or O 2- b. K or K + 32

33 WS #9 - p. 2 (con t) (Unit 3 Review) 8. (Learning Targets 3-3, 3-4, 3-5, 3-6) Complete the following table on Periodic Trends. Electron Configuration (a) Na or S? (b) F or I? Which element has a larger atomic radius? Explain. Na: S: F: I: Which element has a higher ionization energy? Explain. Which element has a higher electronegativity? Explain. 9. (Learning Target 3-7) Draw electron dot diagrams for atoms, showing the correct number of valence electrons. Draw Lewis Dot Structures for the following atoms or ions. Element or Ion Lewis Dot Structure Element or Ion Lewis Dot Structure Na + H Br N 3- Ga C He 33

34 WS #9 p. 3 (con t) (Unit 3 Review) 10. (Learning Targets 3-8, 3-9, 3-14) Complete the following table. Chemical Formula CsBr Electronegativity Difference Type of Bond: Ionic, Polar Covalent or Non-Polar Covalent Distribution of Charge (if Polar Covalent) Lewis Dot Structure NF3 CaO SO3 I2 BrF3 O3 11. (Learning Targets 3-15, 3-15) VSEPR Shapes 34

35 WS #9 p.4 (con t) (Unit 3 Review) Complete the following table. Molecule Lewis Dot Structure VSEPR Shape VSEPR Drawing Polar or Non-Polar? AsCl3 SiH4 OBr2 GeO2 PAt5 SI6 BI3 12. (Learning Targets 3-11, 3-12, 3-13) Polyatomic Ions. Determine the total number of electrons in each of the following polyatomic ions. Draw Lewis structures and indicate the formal charge on each atom. Draw all possible resonance structures for the polyatomic ion. (a) NO2 - (b) BO3 - (c) C2O (Learning Target 3-17) Define miscible and immiscible. 14. Identify two ways to tell whether atoms will form ionic or covalent bonds. Identify four properties of ionic compounds and two properties of covalent compounds. 35

36 Answers. (1) Mendeleev s Periodic Table arranged elements by atomic mass, whereas the modern Periodic Table organizes elements by atomic number. Mendeleev left gaps for elements whose properties were predicted based upon their location on his Periodic Table. Mendeleev did not predict the existence of the noble gases which were discovered later. (2) Element Period Group Valence Electrons Property Ca 4 Alkaline Earth Metals 2 Reactive metals; not found freely in nature F 2 Halogens 7 Salt-formers ; exist diatomically, exist as either solid, liquid or gas at room temperature O 2 VIA 6 Non-metals Ar 3 Noble Gases 8 Do not react to form compounds. Fr 7 Alkali Metals 1 Highly reactive metals; not found freely in nature Ag 5 Transition Metals 1 Metals: malleable, ductile and conductive U 7 Inner Transition Metals/ Actinides Varies Metals: malleable, ductile and conductive; Radioactive (3) Atomic radius is the distance from the nucleus of an atom to the outermost electron. Atomic radius increases as you go down a group, because, with each additional period, another energy level of electrons is being added. Atomic radius decreases as you go across a group, because the number of protons in the nucleus increases, while no new energy levels are being added, drawing the outermost electrons closer to the nucleus. (4) Ionization energy is the energy required to remove the outermost electron. Ionization energy decreases as you go down a group because, with each additional period, another energy level of electrons is being added, increasing the distance between the outermost electrons and the nucleus, decreasing the attractive force. As you go across a period, the increase in nuclear charge without the addition of more energy levels, increases the attractive force between the nucleus and its outermost electrons, making the electrons harder to remove. The two groups that are exceptions to this trend are the Boron group and the Oxygen group. With the Boron group, the removal of its outermost p 1 electron leaves behind a stable ion with a filled s-orbital, making that electron easier to remove than the outermost s 2 electron in the alkaline earth metal group. With the Oxygen group, the removal of its outermost p 4 electron leaves behind a semi-stable half-filled p-orbital, making that electron easier to remove than the outermost p 3 electron in the nitrogen group. (5) Electronegativity is the ability of a nucleus in a chemical bond to attract the electrons of the other atom toward itself. Electronegativity decreases as you go down a group because, with each additional period, another energy level of electrons is being added, increasing the distance between the outermost electrons and the nucleus, decreasing the attractive force. As you go across a period, the increase in nuclear charge without the addition of more energy levels increases the attractive force between the nucleus and its outermost electrons, making the nucleus better able to attract electrons toward itself. The exception to this are the noble gases which have an electronegativity of zero because they do not form bonds. (6) Metals form cations by losing electrons. Since electrons are lost, the cation is smaller in radius than its neutral atom. Nonmetals form anions by gaining electrons. Since electrons are gained, the anion is larger in radius than its neutral atom. (7) (a) O 2- is larger than O because it has 2 more electrons. (b) K is larger than K + because K has 1 more electron. (8) Electron Configuration (a) Na or S? Na: 1s 2 2s 2 2p 6 3s 1 S: 1s 2 2s 2 2p 6 3s 2 3p 4 (b) F or I? F: 1s 2 2s 2 2p 5 I: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 5 Which element has a larger atomic radius? Explain. Na has a larger atomic radius because it has fewer protons than S and its outermost electrons are in the same energy level as sulfur s. The fewer protons means that the Na I has a larger atomic radius than F because its outermost electrons are in the 5 th energy level, whereas F s outermost electrons are only in the 2 nd energy level. 36

37 Which element has a higher ionization energy? Explain. Which element has a higher electronegativity? Explain. (9) has a smaller nuclear charge and exerts a smaller attractive force on the outermost electrons than S. S has a higher ionization energy because it has a greater nuclear charge than Na, while its outermost electrons are in the same energy level as sulfur s. This greater nuclear charge exerts a greater attractive force, requiring more energy to remove an electron. S has a higher electronegativity because it has a greater nuclear charge than Na, while its outermost electrons are in the same energy level as sulfur s. This greater nuclear charge exerts a greater attractive force on its own and other atoms electrons in a bond. F has a higher ionization energy than I. Since fluorine s outermost electrons are closer to the nucleus than iodine s, they experience a greater attractive force which requires more energy to remove them. F has a higher electronegativity than I. Since fluorine s outermost electrons are closer to the nucleus than iodine s, the nucleus is able to exert a greater attractive force on its own and other atom s electrons in a bond. Element or Ion Na + Lewis Dot Structure Na + Element or Ion H Lewis Dot Structure Br N 3- Ga C He (10) Chemical Formula Electronegativity Difference Type of Bond: Ionic, Polar Covalent or Non-Polar Covalent CsBr 2.1 Ionic N/A Distribution of Charge (if Polar Covalent) Lewis Dot Structure NF3 1.0 Polar Covalent 37

38 CaO 2.5 Ionic N/A SO3 1.0 Polar Covalent I2 0 Non-Polar Covalent N/A BrF3 1.2 Polar Covalent O3 0 Non-Polar Covalent N/A 11) Molecule Lewis Dot Structure VSEPR Shape VSEPR Drawing Polar or Non- Polar? AsCl3 Trigonal Pyramidadl Polar SiH4 Tetrahedral Non-Polar OBr2 Bent Polar GeO2 Linear Non-Polar 38

39 PAt5 Trigonal Bipyramidal Non-Polar SI6 Octahedral Non-Polar BI3 Trigonal Planar Non-Polar (12) (a) (b) (c) (13) Miscible = completely soluble in each other; Immiscible = insoluble in each other (14) Ionic compounds are between metals and nonmetals and ΔEN > 1.7. Covalent compounds are between non-metals only and ΔEN <= 1.7. Ionic compounds form crystalline structures, are brittle, have high melting and boiling points and conduct electricity in when molten or in solution. Covalent compounds have lower melting and boiling points and do not conduct electricity 39