Unit 4: Bonding Class Packet - Mrs. Beierman€¦ · Web viewElectron Dot Diagrams and Transfer of Electrons Video Notes: Ionic Bonding & Lewis Structures Friday 11/8 Lewis Structure

Unit 4: Bonding Class Packet

Name:_________________________ Period:__________

DATE TOPIC/ACTIVITY HOMEWORK

Monday 11/4 1. Ionic, Covalent and Metallic Bonds Inquiry Lab

Tuesday 11/5 1. Finish Lab2. Ionic versus Covalent Bonding

Wednesday 11/6 1. Work on Venn Diagram2. Octet Rule

Thursday 11/7 1. Ionic, Covalent and Metallic Properties Quiz

2. Electron Dot Diagrams and Transfer of Electrons

Video Notes: Ionic Bonding & Lewis Structures

Friday 11/8 1. Lewis Structure Practice

Monday 11/11LATE START

1. Lewis Structure Stations

Tuesday 11/12 1. Lewis Structure Quiz2. Start Video Notes

Wednesday 11/13 1. Working with Ions Worksheet2. Introduction to Polyatomic Ions

Thursday 11/14 1. Writing Binary Ionic Formulas 2. Writing Polyatomic Formulas

Friday 11/15 1. Bond with a Classmate Video Notes: Ionic Naming

Monday 11/18 1. Naming Ionic Compounds

Tuesday 11/19 1. Naming Ionic Compounds White Boards

Wednesday 11/20 1. Big Ionic Quiz Video Notes: Covalent Bonding

Thursday 11/21 1. Review Quiz2. Covalent WebQuest

Friday 11/22 1. Covalent WebQuest Finish WebQuest

Monday 11/25 1. Mixed Naming Practice

Tuesday 11/26 1. Mixed Naming Quiz

11/27-11/29 THANKSGIVING BREAK!

12/2 Monday 1. Ionic, Covalent Naming and Formula Review Video Notes: Polarity and IMF

12/3 Tuesday 1. Polarity and Intermolecular Forces

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12/4 Wednesday 1. Molecular Geometry Lab

12/5 ThursdayBEIERMAN ABSENT

1. Review2. Work on Study Guide

12/6 Friday UNIT 4 EXAM

Bonding Notes

I. Introduction to Chemical BondingA. Ionic Bonding

1. Transfer of electrons2. Attraction between large numbers of + ions and - ions

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3. Results when there is large electronegativity difference4. Generally, involves a metal and a nonmetal

B. Covalent Bonding 1. Sharing of electron pairs between two atoms2. Results when there is small electronegativity difference3. Generally, involves two or more nonmetal atoms

II. Ionic Bonding – transfer of electronsA. Attraction between cations and anionsB. Generally, between a metal and a nonmetal

1. Most metals are cations (positive ions)2. Most nonmetals are anions (negative ions)

C. Octet Rule1. Atoms lose, gain, or share electrons to have 8 in the outer level

a. (except H, He, they are fine with 2)2. Most compounds follow the octet rule

a. There are exceptions but we will discuss those next unit.

D. Electron Dot Symbol (Dot Diagrams)1. use dots to represent the valence electrons

2. K Mg Al Si

N S Br Ar

III. Ionic Bonding and Ionic CompoundsA. Ionic Compounds

1. Again, generally contain a metal and a nonmetal2. Cations and anions combined so that charges are equal3. Most form crystalline solids4. Represented by a formula unit; simplest ratio of the ions

B. Electron Dot Symbols and Ionic Bonding1. sodium and fluorine

Na + F

atom atom2. magnesium and chlorine

ClMg +

Cl

atom atoms3


IV. Ions - ReviewA. atoms are neutral because they have equal numbers of protons and electrons, so when they lose or

gain electrons the become charged. B. ions – positive or negative charges atoms that formed when atoms lose or gain electrons

Group Valence Electrons Lose/Gain Charge1 1 lose 1 1+2 2 lose 2 2+

13 3 lose 3 3+14 4 shares None15 5 gains 3 3 -16 6 gains 2 2 -17 7 gains 1 1 -18 8 unreactive none

1. Cations are positive ionsa. metals lose electrons to form + ionsb. Na+ sodium ion lost one electron ( e- )c. Mg 2+ magnesium ion lost two electrons

2. Anions are negative ionsa. nonmetals gain electrons to form - ionsb. monatomic anions have –ide endingc. Cl- chloride ion gained one electrond. O2- oxide ion gained two electrons

3. Transition Metals & Post Transition Metalsa. All metals not in groups 1, 2, & 13b. most have more than one possible chargec. use Roman Numeral to show the charge of transition metals

Fe2+ iron (II) Cu+ copper (I)Fe3+ iron (III) Cu2+ copper (II)

d. Ag+, Zn2+, and Cd2+ have only one possible charge and do not get a Roman Numeral

V. Writing Formulas for Binary Ionic Compounds A. Chemical formulas are chemical symbols that tell what elements are involved in the compound and

the relative proportions1. The subscript (small number) tells you how many atoms of the element involved.

a. H2O – contains 2 hydrogens and 1 oxygenB. Binary contains only two elements (metal + nonmetal)

1. + and – charges must balance2. the positive ion (metal) is written first3. –ide ending usually indicates a binary compound

C. Charge Flip to get formula1. The positive charge on cation becomes the subscript on the anion, and the negative charge on

anion becomes the positive subscript on the cation. a. I.e. Aluminum sulfide

i. Al is +3, sulfide is sulfur, which is -24


ii. So formula is Al2S3

VI. Writing formulas for other ionic compounds A. When more than two elements are present, then the compound contains a polyatomic ion

1. polyatomic ion – group of atoms that bind covalently but act as one ion2. –ite ending indicates one less oxygen than –ate; the charge will be the same

SO42- SO3

2-

sulfate ion sulfite ion

B. Writing formulas for ionic compounds with polyatomic ions1. The charge flip is still used, you will use parenthesis if subscript is added for the polyatomic ion

(meaning you need more than one of that polyatomic ion)2. Examples

a. sodium carbonate - Na2CO3

b. calcium nitrate – Ca(NO3)2

VII. Naming Ionic Compounds

1. Name of Metal + Name of Nonmetal (-ide ending)a. If a polyatomic ion is present then you name same as before but add polyatomic ion name

for anionName of metal + polyatomic ionPolyatomic ion + name of anion (-ide ending)

2. When a transition metal is present, Roman Numerals are used to show the charge of transition metals (except Ag, Zn, Cd because we know these have a set charge)

a. There is three ways to find charge of transition metali. Do it in your head ii. Reverse charge flip

i. The subscript on nonmetal becomes positive charge on metal. ii. The subscript on metal becomes negative charge on nonmetal.

1. Check nonmetals actual charge. If it does not match, multiply each charge by that factor.

2. Use charge to write Roman numeral.iii. Use the following formula

0 = (x)(subscript of metal) + (charge of anion)(subscript of anion)Where x is the charge on the metal and should be your roman numeral

Introduction to BondingSUBSTANCE GROUP #1:

For each of the following substances, place the symbol of the first element in its spot on the periodic table using red ink. Then, place the symbol of the second element in the substance in its spot on the periodic table using black ink. Then label charges up top!

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NaCl LiBr KF ZnCl2 Fe2O3 CuI2 Al2S3

QUESTIONS FOR SUBSTANCE GROUP #1:1. Where are all the first elements located on the periodic table (red symbols)?

2. Based on your knowledge about the periodic table, what “classification” would you give these elements? Metal or Non-Metal?

3. Where are all the second elements located on the periodic table (black symbols)?

4. Based on your knowledge about the periodic table, what “classification” would you give these elements? Metal or Non-Metal

SUBSTANCE GROUP #2

For each of the following substances, place the symbol of the first element in its spot on the periodic table using a red ink. Then, place the symbol of the second element in the substance in its spot on the periodic table using black ink. Then label charges up top!

CCl4 P2O5 N2O4 NI3 PBr3 F2Se

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QUESTIONS FOR SUBSTANCE GROUP #2:1. Where are all the first elements located on the periodic table (red symbols)?


3. Where are all the second elements located on the periodic table (black symbols)?


TYING IT TOGETHER:1. Group 1 substances are called ionic compounds and Group 2 substances are called covalent

molecules. Write a simple rule that will allow you to classify compounds as ionic or covalent on the basis of what you have learned from the model.

2. Did the subscripts (the little numbers shown in the compound formulas) provide any insight into determining whether a substance is ionic or covalent? Why or why not?

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ELEMENT EVALUATION:Fill in the following tables identifying the type of element present in the substance. Use “M” for metal, and “NM” for non-metal.

FORMULA 1ST ELEMENT (M or NM) 2ND ELEMENT (M or NM) CLASSIFICATION(Ionic or Covalent?)

NaBr

SF6

CoBr2

BaS

NO2

C6H6

CrCl3

CO2

MnO2

PbCl2

OF2

CsF

Fill in the Venn diagram concerning ionic, covalent, and metallic using the following properties as well as any others you can think of- use PENCIL so that we can make CHANGES!

Burns easily, hard to burn, High melting points, Low melting points, High boiling points, Low boiling points, make crystal structures, electrons are shared, electrons are transferred, delocalized electrons, made of metals, made of non-metals, made of anions, made of cations, named with prefixes, named with roman numerals, large differences in electronegativity, small differences in electronegativity, make alloys, lots of energy to form bonds, has electrostatic attractions, makes molecules, makes polymers, minerals, plastic, conduct electricity when dissolved, conducts electricity as a solid, does not conduct electricity, brittle, electrolyte, non-electrolyte

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Types of Bonds

Complete the Venn diagram below with properties of ionic and Covalent Bonds:

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Metallic

CovalentIonic


Bonds formed between two nonmetals are _______________ and involve the ____________ of electrons.

Bonds formed between two metals are _______________ and involve the ____________ of electrons.

Bonds formed between metals and nonmetals are ____________ and involve the ___________ of electrons.

Describe the following as ionic, metallic, or covalent:NaCl _____ Al _____ Lithium _____CO2 _____ C6H12O6 _____ Strontium bromide _____Au _____ Ti _____ Tin (II) chloride _____MgBr2 _____ K2O _____ Nitrogen (IV) oxide _____Fe _____ CH4 _____ Hydrogen selenide _____H2O _____ H2S _____ Copper (II) phosphate _____Ca3(PO4)2 _____ PI3 _____ Lead (IV) nitrate _____

Lewis Structures Introduction

A Lewis structure is a way to represent the valence electrons of an atom using dots. These valence electrons are important because they are involved in

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Ionic Bonds

Covalent Bonds


bonding. The Lewis structure for an atom is formed by writing the element symbol and placing one dot next to the symbol for each valence electron. For example, the Lewis structure for carbon would have four dots surrounding it, while the Lewis structure of chlorine would have seven.

In a molecule, atoms are held together with covalent bonds. These bonds are formed when pairs of electrons are shared. For example, carbon will bond with four chlorine atoms to form the molecule carbon tetrachloride, CCl4. Each bond in the molecule represents a pair of shared electrons and is represented as a dash. The guiding principle behind bonding in atoms is the octet rule which states that each element is trying to get a full set of valence electrons. When carbon shares electron pairs with four chlorine atoms it has an octet. Each chlorine atom has three nonbonding pairs of electrons and one bonding pair. This arrangement is favorable because both carbon and the chlorine atoms have an octet of electrons

ElementMetal or

Nonmetal?

Lewis Dot Structure

as an ATOM

Gain or lose

electrons?

How many

e-?

Lewis Dot Structure

of Stable ION

Becomes like

which noble gas?

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Fluorine

Lithium

Aluminum

Sulfur

Radium

Phosphorous

Using your chart, draw Lewis structures for the following compounds:

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LiF Li3P

AlF3 Aluminum Sulfide

Radium Fluoride Radium Sulfide

Li2S

Aluminum Phosphide

Radium Phosphide

Summarize:-Why do metals bond with nonmetals?

-What determines how many atoms of the metal and nonmetal will combine to form the compound?


Lewis Dot Structures Practice

How many valence electrons do cations show in the Lewis dot diagrams? ____ anions? ____

Draw the following ionic compound.Lewis Diagram Lewis Diagram

NaF CsCl

K2O MgO

Rb3N SrI2

CaBr2 BaS

SrO Fe2O3

Mg3P2 Ag2S

AlI3 CuO

CuS NiCl3

CrN TiO2

MnO2 PtCl2

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Working with Ions

1. What is an ion?

2. What is a cation?

3. What is an anion?

For each element below, tell how many electrons it would prefer to gain/lose in order to

fulfill the Octet Rule The first one is done for you as an example.

Element Gain or Lose Electrons? Ion Formed Type of Ion

Mg lose 2 electrons Mg2+ cation

Na

F

O

K

Al

Li

Cl

Br

Any Noble Gas

5. What is the “octet rule”?

6. Which elements are “exceptions” to the octet rule?

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Lewis Structures Part 2Molecules that don’t follow octet rule

There are some cases in which all of the elements in a Lewis structure will not follow the octet rule. Recall that the octet rule states atoms will share valence electrons in a way that gives them a complete set (8). In a few examples the elements have less than octet. Boron trihydride is an example of a compound where an element has less than an octet. Other molecules may contain elements with more than an octet of electrons. In the phosphorus pentafluoride molecule the bonded phosphorus atom has 10 electrons around it. Elements in period 3, 4, 5, and 6 can take extra electrons because they have an empty d orbital available.

Resonance Resonance occurs when more than one valid Lewis structure can be written for a compound. The

nitrate ion, NO3– has a double bond between nitrogen and one of the oxygen atoms. The double bond could form with any of the three oxygen atoms to give a working model of the molecule. Experimentation shows that there are not three distinct models for this ion. Instead there is one structure that is a hybrid of all three – a sort of average. Resonance structures only differ by the placement of their electrons. These Lewis structures are drawn by creating models for each and putting double-headed arrows between each structure.

Draw Lewis Structures for the following. 1. BF3 6. ICl4–

2. AsF6– 7. BeH2

3. ClF3 8. TeF4

4. SF6 9. SnCl3–

5. RnCl2 10. PO43-

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Bond with a Classmate

You will be given either a positive or negative ion. Your job is to find other people to “bond” with. When you have finished filling in the information for 4 compounds, go to Mrs. Beierman to get a new ion. Then, bond 5 more times. Keep going until you’ve filled in all the blanks below.

Cation Anion Lewis Dot Diagram (Show the transfer of electrons) Final Formula

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Cation Anion Lewis Dot Diagram Final Formula

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Writing and Naming Binary Ionic Formulas

1. Na2S

2. Li2O

3. BeO

4. AlCl3

5. AlN

6. MgS

7. Ag2S

8. CaF2

9. MgCl2

10. Al2O3

11. Li3P

12. BeF2

13. K3P

14. Mg3P2

15. CaO

16. NaCl

17. Sr3P2

18. BaO

19. Calcium bromide

20. Potassium phosphide

21. Beryllium fluoride

22. Aluminum oxide

23. Magnesium oxide

24. Strontium phosphide

25. Potassium sulfide

26. Aluminum nitride

27. Beryllium oxide

28. Sodium fluoride

29. Lithium selenide

30. Barium oxide

31. Aluminum sulfide

32. Sodium chloride

33. Aluminum iodide

34. Sodium iodide

35. Rubidium nitride

36. Magnesium carbide

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Writing and Naming Binary Ionic Compounds (transition metals with single element anions)

1. CuS

2. SnF2

3. PtO2

4. PbO2

5. SbP

6. CuCl2

7. SnCl4

8. Au3P

9. AuAs

10. Cu2O

11. SnS2

12. CrP2

13. FeCl3

14. SnBr2

15. SnI

16. PbS2

17. CuCl

18. BiBr3

19. AuF3

20. Hg3P2

21. Iron (III) chloride

22. Copper (II) oxide

23. Antimony (III) sulfide

24. Mercury (II) oxide

25. Tin (IV) iodide

26. Iron (III) phosphide

27. Cobalt (II) sulfide

28. Platinum (IV) sulfide

29. Nickel (II) oxide

30. Copper (II) bromide

31. Tin (IV) fluoride

32. Manganese (II) chloride

33. Nickel (II) phosphide

34. Cobalt (III) oxide

35. Vanadium (V) sulfide

36. Gold (III) arsenide

37. Mercury (II) phosphide

38. Titanium (III) carbide

39. Tin (II) fluoride

40. Lead (IV) nitride

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Writing and Naming Polyatomic Compounds (All metals bound to a polyatomic ion)

1. Li2SO4

2. Na2SO4

3. Al(ClO3)3

4. AgNO3

5. MgSO4

6. K3PO4

7. K2CO3

8. LiClO3

9. Sn(SO3)4

10. PbSO4

11. Sn(OH)4

12. AuOH

13. Sb2(SO4)3

14. Cu(NO3)

15. Sr(ClO3)2

16. Pb(ClO3)4

17. Mg(ClO3)2

18. Mg(NO3)2

19. Tin (II) phosphate

20. Uranium (IV) nitrate

21. Magnesium nitrate

22. Magnesium hydroxide

23. Barium sulfate

24. Calcium sulfate

25. Lithium hydroxide

26. Sodium hydroxide

27. Magnesium chlorate

28. Chromium (III) phosphate

29. Copper (II) sulfate

30. Nickel (II) sulfate

31. Mercury (I) chlorate

32. Vanadium (V) carbonate

33. Ammonium sulfate

34. Nickel (II) carbonate

35. Cobalt (II) hydroxide

36. Lead (II) chlorate

Covalent WebQuest

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Open the following PBS Covalent Bonding tutorial link and answer the following questions as you read through each page: https://pbslm-contrib.s3.amazonaws.com/WGBH/arct15/SimBucket/Simulations/chemthink-covalentbonding/content/index.html

● Page 2-3:How is the movement of electrons different when the atoms are close?

What happens if you try to move the atoms very close to each other?

● Pages 6-9Describe what is happening in a covalent bond.

How is this different from an ionic bond?

● Page 11-13What kind of atoms have the strongest attraction for another atom’s electrons? What is this property called?

*Skip pages 14-20*● Pages 21-24

Fill in the following table:

Type of Bond Number of e- Shared Relative Strength

● Pages 25-35Looking at the examples on pages 32-33, how can you tell that these chemical formulas are covalent and not ionic just by looking at them? In other words:An ionic bond is between a _______________ and a _________________.A covalent bond is between a _________________ and a ________________.

After completing the practice in the tutorial, fill in the missing names or formulas for the following. The prefixes can also be found on the back of your periodic table.

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Chemical Name Chemical Formula

Antimony tribromide

Hexaboron monosilicide

Chlorine dioxide

Dinitrogen trioxide

SeF6

CH4

NF3

P4S5

Open the following website:https://phet.colorado.edu/sims/html/molecule-shapes/latest/molecule-shapes_en.htmlClick on the “Model” page. Check the “Show Bond Angles” box. Take about 5 minutes and play around with this molecule builder. Get to know some of things you can do with it, then answer the following questions.

1. Build a molecule with two atoms and a single bond. Draw this below:

This shape is called LINEAR. Why do you think this name was chosen?

2. Build a molecule with three atoms and two double bonds. Draw this below including bond angles:

This shape is also called LINEAR. Why do you think this name was chosen?

3. Build a molecule with 4 atoms, 3 single bonds. Draw this below including bond angles:

This shape is called TRIGONAL PLANAR. Why do you think this name was chosen?

4. Build a molecule with 3 atoms, 2 single bonds, and 1 lone pair of electrons. Draw this molecule below including bond angles:

This shape is called BENT. Why do you think this name was chosen?

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How is this shape similar and different from the trigonal planar shape?

5. Build a molecule with 5 atoms and 4 single bonds. Draw this below including bond angles:

This shape is called TETRAHEDRAL. Why do you think this name was chosen?

6. Build a molecule with 4 atoms, 3 single bonds, and 1 lone pair of electrons. Draw this below including bond angles:

This shape is called PYRAMIDAL. Why do you think this name was chosen?

How is this shape similar and different from the tetrahedral shape?

7. **What happens to the bond angle as more atoms or lone pairs of electrons are added in? Why do you think the atoms do this? (Hint: What else is around the outside atoms that we aren’t being shown?)**

Use the information from number 1-7 above to fill in the chart below:

Number of Atoms

Number of Bonding Areas (single or double bond) Around Central Atom

Number of Lone Pairs Around Central Atom

Sketch of Molecule Shape Name

2 NA NA

23


3 2 0

3

1 or 2 lone pairs

4

4

5

You will be responsible for knowing the shapes above (MEMORIZE THEM) . . . there are also molecules that have 5 or 6 atoms which are shown below.

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On the bottom of the screen, switch the type to “Real Molecules”. On the top right, there is a drop down box for you to change the molecules. Look through them and decide on the name of the shape for each molecule.

Formula Shape Formula Shape

H2O CH4

CO2 SF4

SO2 XeF4

XeF2 BrF5

BF3 PCl5

ClF3 SF6

NH3

Covalent Bond Practice

1. Fill in the missing information on the chart.

Element # of Protons # of Electrons # of Valence Electrons

Carbon

Hydrogen

Chlorine

Helium

Phosphorus

Oxygen

Sulfur

Nitrogen

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2. For each of the following covalent bonds: Write the number of valence electrons. Draw the Lewis Structure

a) SiCl3 b) SCl2

c) NOCl d) NCl3

e) COCl2 f) SiF4

g.) O2 h.) CO

i.) I2 j.) AsH3

3. How does the octet rule relate to the electron configuration of a noble gas?

4. Explain the difference between an ionic and covalent bonds.

5. Why do some molecules show resonance structures?

6. Explain the sharing of electrons with respect to polar and nonpolar covalent bonds.

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7. Rank the intermolecular forces from strongest to weakest.

8. Explain VSEPR theory.

9. Indicate the number of valence electrons and draw Lewis structures for the following. You must show unshared electron pairs.

a. I2 b. OF2 c. CS2

d. OH- e. PO43- f. H3O+

10.Indicate whether the bonds in these compounds are ionic, polar covalent, or nonpolar covalent.

a. KClb. MgF2 c. HCl d. H2O

11. Indicate the number of valence electrons and draw the resonance structures for, carbonate ion CO3

2-.

Write the formula for the covalent compound 12. nitrogen tribromide

13. hexaboron monosilicide27


14. chlorine dioxide

15. hydrogen moniodide

16. iodine pentafluoride

17. dinitrogen trioxide

18. ammonia (nitrogen trihydride)

19. phosphorus triiodide

20. dihydrogen monoxide

21. diphosphorous pentoxide

Write the names for the following compounds:

22. P4S5

23. O2

24. SF6

25. Si2Br6

26. CCl4

27. CH4

28. B2Si

29. NF3

30. H2O

31. N2O5

Lewis Structure Practice for Covalent CompoundsDraw the Lewis structure for the following ions:

32. ClO4-

33. NH4+

34.SO32-

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35.N3-

36.OH-

37.BrO3-

38.Draw resonance Lewis structures for NO3-

39.Draw resonance Lewis structures for CO3-2

Mixed Practice Naming

Write the names of the following compounds:

1. PbCl22. SBr2

3. P4H10

4. NaC2H3O2

5. AgNO3

6. Ba(OH)2

7. NH4NO3

8. Na2SO4

9. CuSO4

10. Cu2SO4

11. Fe2(SO4)3

12. C2H6

13. Cr(OH)3

14. Ba(NO3)

15. KMnO4

16. K2CrO4

17. P2O5

18. N2O3

19. CoF3

20. SnI4

29


21. CaO

22. PbCl4

23. Ba(C2H3O2)2

24. PBr5

25. (NH4)3PO4

26. CaCO3

27. NH4C2H3O2

28. AlPO3

Write the formulas of the following compounds:1. iron (II) sulfide

2. copper (I) oxide

3. dinitrogen pentoxide

4. potassium permanganate

5. cesium chloride

6. lead (II) nitrate

7. xenon hexafluoride

8. lithium phosphide

9. lithium phosphate

10. ammonium oxide

11. dinitrogen monoxide

12. tetraphosphorus decoxide

13. magnesium chloride

14. magnesium hydroxide

15. mercury (II) bromide

16. cobalt (II) bicarbonate

17. tin (IV) fluoride

18. ammonium sulfate

19. diarsenic trisulfide

20. barium phosphate

21. aluminum permanganate

22. strontium oxide

23. strontium chloride

24. calcium carbonate

25. calcium sulfide

26. calcium sulfate

27. calcium sulfite

28. potassium nitrate

Bond Polarity Introduction

1 2 13 14 15 16 17 18

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In the empty spaces of the chart above, fill in the element symbols and electronegativity values. Notice the transition metals are absent.

1. Electronegativity values generally _______________ down a

group and ____________ across a period.

2. Metals tend to have ____________ electronegativity values

and nonmetals are _____________ values.

3. When lithium bonds with fluorine they form an

________________ bond.

a. What is the electronegativity difference of lithium and fluorine that might help characterize the

properties they have? _____

b. List some of the properties:

4. When fluorine bonds with another fluorine atom they form a ________________ bond.

a. What is the electronegativity difference of the two fluorine atoms that might help characterize the

properties they have? _____


5. When hydrogen bonds with fluorine atom they form a ________________ bond.

a. What is the electronegativity difference of the hydrogen and fluorine atoms that might help

characterize the properties they have? _____


6. HF and F2 are both ________ compounds, but they actually have some slightly different properties. F2 is

not attracted to an electromagnetic field, where HF is. HF has a high boiling point, but F2 is very low.

Based on the information you obtained so far, what characteristic might cause these differences?

7. Fill in the chart below:

Electronegativity Difference Type of Bond

0.0-0.4

0.5-1.4

1.5-4.0

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Using your table above find the electronegativity difference for each substance. If more than one bond is formed, find all differences. Then, check which bonds are present.

Substance Electronegativity difference(s) Ionic Covalent Polar NonpolarI2

PCl3SiO2

Br2

CO2

NaClCH4

N2O5

NH3

KClNaNO3

KClO3

Ca(ClO3)2

Molecular Polarity

Fill in the chart below.

32


1. In terms of lone pair electrons, how can you determine if a molecule is polar?

2. What molecular shapes are always polar?

3. What molecular shapes are always nonpolar?

4. How can a molecule be nonpolar if it contains polar bonds?

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Documents

Unit 4: Bonding Class Packet - Mrs. Beierman€¦ · Web viewElectron Dot Diagrams and Transfer of Electrons Video Notes: Ionic Bonding & Lewis Structures Friday 11/8 Lewis Structure