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Unit 11:

Ionic & Metallic Bonding (Chapter 7)

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Unit 11:

Covalent Bonding (Chapter 8)

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Unit 11 ~ Problem Set #1

Read pg. 187-199. Pg. 193; 7.1 Section Assessment #9, 10, 11 9. How many electrons will each element gain or lose in forming an ion?

a. Calcium - _____________________

b. Fluorine - _____________________

c. Aluminum - _____________________

d. Oxygen - _____________________

10. Write the name and symbol of the ion formed when

a. a potassium atom loses one electron ________________________________

b. a zinc atom loses two electrons ________________________________

c. a fluorine atom gains one electron _________________________________

11. Write the electron configuration of Cd+2. ___________________________________________

Pg. 196; practice problems #12, 13 12. Use electron dot structures to determine formulas of the ionic compounds formed when

a. potassium reacts with iodine.

13. What is the formula of the ionic compound composed of calcium cations and chloride

anions? __________________

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Pg. 199; 7.2 Section Assessment #15, 16, 18, 20, 22 15. What properties characterize ionic compounds? ______________________________________

_________________________________________________________________________________

16. Define an ionic bond. _____________________________________________________________

_________________________________________________________________________________

18. Write the correct chemical formula for the compounds formed from each pair of ions.

a. K+, S-2 _____________________________

b. Ca+2, O-2 _____________________________

c. Na+, O-2 _____________________________

d. Al+3, N-3 _____________________________

20. Which pairs of elements are likely to form ionic compounds?

a. Cl, Br _____________________________

b. Li, Cl _____________________________

c. K, He _____________________________

d. I, Na _____________________________

22. Why do ionic compounds conduct electricity when they are melted or dissolved in water?

________________________________________________________________________________ ________________________________________________________________________________

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Bonding Pre-Quiz

1. Answer the following questions about the element, arsenic, Z=33. a. Write the longhand spectroscopic notation for this METALLOID. b. Draw its shorthand orbital notation diagram. c. How many valence electrons does arsenic have? _____________________________ d. Draw its Lewis dot structure. e. Predict the 3 ions this METALLOID makes. ___________________________________ f. Draw the Lewis dot structures for the three ions. 2. What types of elements make up an ionic compound? Give two examples. ____________ ______________________________________________________________________________ 3. What types of elements make up a molecular compound? Give two examples. ________ ______________________________________________________________________________ 4. a. What is the most electronegative element? __________________________________ b. Define electronegativity.___________________________________________________ ______________________________________________________________________________ 5. a. What are the names of the two types of ions? _______________________________ b. How is each one formed? _________________________________________________ ______________________________________________________________________________ c. Give two examples of each. ______________________________________________ 6. What does every element strive for? _____________________________________________

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Electron Arrangement and Bonding Video Questions

Video 1: How Atoms Bond

1. Nuclei and electrons, both within atoms and between atoms, exert ___________________________

forces on each other, while two nuclei or two electrons of adjacent atoms exert

_______________________________ forces on each other.

2. T-F: In the H—H bond, the electrons are always found in the same place. ______________________

3. Bonds in which atoms share electrons are called _______________________________ bonds.

4. The reason two helium atoms do not bond together is because the forces of

__________________________ are greater than the forces of _______________________________,

leaving the helium atoms far apart.

5. The family of atoms which normally do not bond together are called the _____________________.

6. There are two reasons why sodium’s 3s electron is held quite loosely by the nucleus. 1) This outer

electron is located far from the nucleus, thus decreasing its ______________________________ to the

nucleus, and 2) the atom’s other ten electrons _______________________________ the outer

electron, which _______________________________ its attraction to the nucleus.

7. Atoms shrink across a period because of increasing _______________________________ charge.

8. In a bond between sodium and chlorine, an electron is _______________________________ to

chlorine, resulting in the formation of ions for both sodium and chlorine.

9. The name given to the type of bond between sodium and chlorine, in which an electron is

transferred, is an ______________________________ bond.

Video 2: Molecular Substances and Covalent Crystals

1. In the Cl—Cl bond, the atoms share _______________ pair of electrons to make a single covalent

bond.

2. In the O—O bond, each atom must gain _______________ electrons to acquire the electron

configuration of neon.

3. Thus, in the diatomic element oxygen, there are _______________ shared pairs of electrons in

their molecules, creating a _______________________________ covalent bond.

4. Diatomic nitrogen atoms share _______________ pairs of electrons in their molecules, creating a

_____________________covalent bond.

5. T-F: The strength of covalent bonds varies as follows: Triple > Double > Single

_______________________________

6. In the HCl molecule, _______________________________ has a stronger attraction for electrons,

thus the electrons spend most of their time there.

7. A molecule in which partially positive and partially negative charges exists is called ____________.

8. T-F: Polar molecules have higher boiling and melting points than nonpolar molecules. _________

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9. Carbon can form up to _______________________________ single covalent bonds.

10. T-F: In general, the larger the molecule, the higher the melting and boiling points. ____________

11. Molecules such as diamond are called _______________________ ___________________________.

Video 3: Metals and Ionic Solids

1. T-F: In metals, electrons are free to move from atom to atom. _______________________________

2. Many metals have similar electron configurations. Their outer electrons are located in a(n)

________ orbital.

3. As a result of the location of the outer electrons, a bond can form in ________________ direction

and are ________________ held by the nucleus, which means that the electrons can come under

the influence of an _____________________________ atom.

4. The atoms in a metallic bond behave like a series of _____________________ positive charges in a

____________________ of free flowing electrons.

5. The “free to move” idea of electrons in metals explains a metal’s ability to conduct both

______________________ and _______________________________ .

6. Metallic bonds are ______________ directional, so the atoms can ________________________ over

each another. This makes it possible to press metals into thin ____________________ and stretch

into thin ____________________.

7. The relative size of the two ions plays a role in determining the ___________________ of the

________________________ as they pack together.

8. The fact that the ions in an ionic solid are held in fixed positions by their attractive forces make the

ionic solid ______________________________ and ______________________________. This also

enables them to be good _______________________________ of electricity and have

____________________ melting points.

9. In order to conduct electricity, an ionic solid must be dissolved or in the _______________________

state so that __________________ ____________________ ions are present.

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What Kind Of Bond Am I? Unlike Thomas Jefferson’s view of people, all atoms are not created equal. We have already

seen that there are complex patterns to the trends in ionization energy, electron affinity

and electronegativity that describe an atom’s ability to gain, lose or attract electrons. We

know that it is easier to remove an electron from K than it is from Li, but we’ve not

considered how these various atomic effects might affect bonding between atoms.

A chemical bond is a linkage of atoms by transferring or sharing valence electrons, those outermost _______ and _______ electrons. But why do atoms bond together? We can answer that by looking at the group of atoms that do not bond — ________________________. The reason they are so stable is their outer s and p orbitals are filled, the same electron configurations that other atoms strive to achieve. This can occur by rearranging the outer electrons between two atoms. From the definition, we can see that there are two general types of bonds in compounds, those that result from electron _________________ and the second class that results from electron _______________. When electrons transfer from one atom to another, positive and negative ions are formed that are held together by an electrostatic attraction. This is called an ionic bond and is usually formed between a _________________________ and a

_________________________. EXAMPLE: NaCl ----> the Na atom has _____ valence electron; Cl atom has _____ valence electrons. In order to achieve a noble gas electron configuration, the Na atom will give up its electron to the Cl atom, creating ___________ and __________ ions. The two ions of opposite charge are attracted to one another and held together by an _____________________ ____________________.

In the second type, usually formed between _______ ________________, the electrons in the bond are shared between the two atoms, giving each atom _______ electrons. This is called a covalent bond. EXAMPLE: Cl2 ----> Each Cl atom has _______ valence electrons.

In order to achieve a noble gas electron configuration, each Cl atom will share its unpaired electron with the other Cl atom, creating a single covalent bond between the two atoms. Now each chlorine has _____ pairs of unshared electrons and _____ shared pair, for a total of _______ valence electrons. In covalent bonds, the electrons in the bond are shared by two atoms. In the example above, both chlorine atoms

have identical attractions for these shared electrons, creating a pure covalent bond ------> ________________ sharing.

Key Points

___________________

___________________

___________________

Key Points

___________________

___________________

___________________

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This is not always the case. If we examine the water molecule, we will see why. EXAMPLE: H2O ----> Each H atom has _____ valence electron, and the oxygen has _____ valence electrons. In order to achieve a noble gas electron configuration, each H atom will share its single valence electron with the oxygen’s 2 single valence electrons, creating two single covalent bonds. Now each hydrogen atom has _____ electrons, and the oxygen has _____.

In water’s two covalent H—O bonds, the electrons in the bond are not shared equally by the two atoms. Oxygen, which has a stronger attraction for electrons than hydrogen, pulls the electrons towards itself. This creates a polar covalent bond ------> _______________ sharing. In a polar covalent bond, the more electronegative element

assumes a partial negative charge (δ-), while the other atom in the bond takes on a slightly positive charge (δ

+).

In this introductory course, we can predict bond types by using the following rule: * Ionic bond --> metal + __________________

* Covalent bond --> nonmetal + ____________________

If the two nonmetals are identical --> ______________ covalent.

If the two nonmetals are different --> ______________ covalent.

This rule is only a rough approximation and works well for compounds of the Group 1 or 2 metals with the halides but doesn’t reflect reality when we consider bonds between carbon and most metals. Carbon is certainly a nonmetal, but it forms covalent bonds (sometimes highly polar, but covalent none the less) with almost all of the elements on the periodic table. There needs to be a more accurate way to determine bond types. In the 1930’s, Linus Pauling (1901 - 1994), an American chemist who won the 1954 Nobel Prize, recognized that

bond polarity resulted from the relative ability of atoms to attract electrons. Pauling devised a measure of this electron attracting power. He called it “electronegativity,” which he defined as the “power of an atom in a molecule to attract electrons to itself.” Electronegativity only has meaning in a _______________. When we are looking at a chemical bond to determine whether its actual bond type is pure covalent, polar covalent, or ionic, we must compare the electronegativities of the atoms in the bond. Simply subtract the EN values of the two elements in the bond and compare to the chart below:

0 - 0.3 ----> pure covalent

0.4 - 1.7 ----> polar covalent

1.8 - 4.0 ----> ionic

EXAMPLE: NaCl ----> ENNa - ENCl = _______ - _______ = _______, a(n) ____________________ bond. Cl2 ----> ENCl - ENCl = _______ - _______ = _______, a(n) _________________________ bond. H2O ----> ENO - ENH = _______ - _______ = _______, a(n) _________________________ bond.

Key Points

___________________

___________________

___________________

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Keeping It Together Directions: 1) Examine the elements in the compound and predict the bond type. 2) Then use the electronegativity chart to determine the difference in EN between the two elements. 3) Determine the actual bond type by using the chart below.

0 - 0.3 --> pure covalent 0.4 - 1.7 --> polar covalent 1.8 - 4.0 --> ionic 4) Lastly, determine which of the elements in the bond is more negative (i.e. has a higher

electronegativity).

Predicted Bond Type

Electronegativity Difference

Actual Bond Type

More Negative

Atom

1. CaO ________________ ________________ ________________ ________________

2. KI ________________ ________________ ________________ ________________

3. F2 ________________ ________________ ________________ ________________

4. N2O ________________ ________________ ________________ ________________

5. SnH2 ________________ ________________ ________________ ________________

6. CH4 ________________ ________________ ________________ ________________

7. SbCl5 ________________ ________________ ________________ ________________

8. PbS ________________ ________________ ________________ ________________

9. As2O3 ________________ ________________ ________________ ________________

10. Mg3N2 ________________ ________________ ________________ ________________

H = 2.1 x x x x x x He

Li = 1.0 Be = 1.5 B = 2.0 C = 2.5 N = 3.0 O = 3.5 F = 4.0 Ne

Na = 0.9 Mg = 1.2 Al = 1.5 Si = 1.8 P = 2.1 S= 2.5 Cl = 3.0 Ar

K = 0.8 Ca = 1.0 Ga = 1.6 Ge = 1.8 As = 2.0 Se = 2.4 Br = 2.8 Kr = 3.0

Rb = 0.8 Sr = 1.0 In = 1.7 Sn = 1.8 Sb = 1.9 Te =2.1 I = 2.5 Xe = 2.6

Cs = 0.7 Ba = 0.9 Tl = 1.8 Pb = 1.9 Bi = 1.9 Po = 2.0 At = 2.2 Rn = 2.4

Fr = 0.7 Ra = -.9

DIRECTIONS: Refer to the answers from 1-10 to answer the following questions. 1. What kind of elements usually form ionic compounds? _________________________________

2. What kind of elements usually form covalent compounds? _____________________________

3. What kind of bonds are usually formed when a metalloid is present in the compound? ____________

4. What kind of compound always contain pure covalent bonds? __________________________

5. Using only a periodic table, how would you predict the bond type? _____________________________

_________________________________________________________________________________________

Electronegativity Values

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Chemical Bonding

How can the physical properties of substances be explained?

Physical properties of substances can be explained in terms of intermolecular forces that exist ________________ molecules and chemical bonds found between ________________ in a molecule. These properties include conductivity, malleability, ductility, solubility, hardness, melting point, and boiling point.

  Conductivity is the ability to carry an electrical current when electrons or ions are ____________________. Malleability is the ability of solids to be ________________________ into thin sheets. Ductility is the ability of solids to be _____________________ into wires. Solubility is the ability to ____________________ in a solvent. Hardness is the ability to resist ____________________ changes. Melting point is the temperature at which a substance melts. Boiling point is the temperature at which a substance boils.

 

Ionic solids Example: NaCl(s) Note: The crystal shape is cubic.

Ionic solids have relatively strong attractions between the _____________.

Their physical properties are as follows: 1) Because there are no free moving electrons or ions, ionic solids do/do not (circle one) conduct heat

and electricity, unless the ionic solid is dissolved ( _______________________), or melted

(_______________________).

2) Because particles cannot slide over one another they are/are not malleable. This causes ionic solids to

be __________________________.

3) Because the ions are attracted to the δ+ and δ- ends of water, they have low/high solubility in water.

4) Since ionic bonds are strong, their crystals are soft/hard.

5) Because ionic solids have strong attractions between the ions, the compounds have low/high

melting and boiling points.

Molecular (covalent) Compounds Example: SiO2 (s)

Molecular compounds have relatively _____________ intermolecular forces between the molecules. As a result, they exist as solids, liquids, and gases at room temperature. In general, their physical properties are as follows:

1) Because there are no free moving electrons or ions, they do/do not conduct heat and electricity, ever.

2) In the solid state, the particles are “glued” in position and cannot slide over one another, so they

are/are not malleable. This causes molecular solids to be ____________________.

3) In general, covalent compounds have low/high solubility in water.

4) The intermolecular forces between covalent compounds are relatively weak and easy/hard to break.

Therefore, solid covalent compounds are soft/hard, and the compounds have low/high melting and

boiling points.

5) The intermolecular forces of pure covalent compounds are weaker/stronger than polar covalent

bonds.

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Metallic Solids Example: Al (s)

Metallic bonds are formed because valence _____________ are mobile. Metals have ionic/covalent/metallic bonds between the particles. ______________ electrons in the metals are not confined to each atom but are shared by all. Therefore, the _______________ ions (cations) are said to be immersed in a “moving sea” of _______________. The physical properties of metals are as follows: 1) Because there are free moving __________________________ present, they are good/bad conductors.

2) Because the positive ions can slide over one another, metals are/are not malleable and ductile.

3) They have low/high solubility in water.

4) Since metallic bonds are also strong, they are soft/hard and the atoms have low/high melting and

boiling points.

5) Examine the picture above. Metallic bonds are/are not found in compounds.

Place the 4 types of bonds (ionic, polar covalent, pure covalent, metallic) in order of increasing strength.

___________________________________________________________________________________

Place the four types of bond in order of increasing melting/boiling points. Explain your reasoning. ____________________________________________________________________________________________ ____________________________________________________________________________________________ ____________________________________________________________________________________________ Refer to the two diagrams below and explain why metals are malleable and ionic solids are not.

Metals are malleable because the free flowing __________________ act as a lubricant and allow the

__________________ to __________________ over each other.

However, in ionic solids, the __________________ and __________________ are held in

__________________ positions. When a force is applied to the crystal (examine the bottom right diagram),

__________________ charges line up. When similar ions are close to each other, they attract/repel (circle

one) one another, causing the crystal to_____________________.

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Why Atoms Bond Name _________________ Directions:

a. Go to www.wwnorton.com/web/chem1e/. Click on the picture.

b. Choose chapter 8 - Chemical Bonding and Atmospheric Molecules.

c. Scroll down to Chapter 8 Tours and choose Bonding.

d. Answer the following questions as you go through the tutorial.

1. How are chemical bonds formed? ___________________________________________________________

________________________________________________________________________________________

2. Where are valence electrons located? _________________________________________________

3. What are the two different kinds of bonds that atoms form? ___________________ and

___________________

4. How are ionic bonds formed? ______________________________________________________________

________________________________________________________________________________________

5. An atom can ___________________ an electron to become positively charged. This is called a

__________________.

6. An atom can gain an electron to become ___________________ charged. This is called an

___________________.

7. When an atom of sodium bonds with fluorine, ___________________ electron is transferred from the

___________________ atom to the ___________________ atom.

8. What happens to the size and charge of the two atoms in the bond after this electron transfer occurs?

________________________________________________________________________________________

________________________________________________________________________________________

9. When atoms form an ionic bond, each atom in the bond achieves a ___________________

___________________ ___________________ with its ___________________ electrons.

10. The attraction between the ___________________ and ___________________ forms an

___________________ ___________________.

11. An ionic bond is typically formed between a ___________________ and a ___________________.

12. In general, metallic elements ___________________ electrons to become isoelectronic to the

___________________ ___________________ ___________________. Nonmetallic elements

___________________ electrons to become isoelectronic to the ___________________

___________________ ___________________.

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13. Use the color coded periodic table to write correct formulas for three compounds that are held together

with an ionic bond. EX: NaCl, MgBr2, Al2O3 ___________________, ___________________, and

___________________

14. What kinds of elements form covalent bonds? ___________________

15. How are covalent bonds formed? __________________________________________________________

16. In the F2 molecule, how many electrons are shared in the single bond? ___________________

17. By sharing one pair of electrons, each atom in the molecule is surrounded by ___________________

electrons.

18. Do all covalent compounds share electrons equally? ___________________ Explain your answer.

________________________________________________________________________________________

________________________________________________________________________________________

________________________________________________________________________________________

19. When the electrons are shared unequally, the bond is called a ___________________ bond and

produces a ___________________ ___________________.

20. In a polar covalent bond, the element with the higher electronegativity value acquires a partial

___________________ charge.

21. Click on the atom to change the element. As the difference in electronegativity values between the

two atoms decreases, the strength of the dipole moment ___________________.

22. If an atom needs more than one electron to achieve a noble gas configuration, it will form __________

___________.

23. How many electrons does an oxygen atom need to complete its octet? ___________________

24. How many electrons are shared in a single bond? ______________ in a double bond? ______________

25. In the O2 molecule, each O atom has ___________________ shared pairs of e- and ___________________

unshared pairs of e-.

26. How many electrons does a nitrogen atom need to complete its octet? ___________________

27. How many electrons are shared in a triple bond? ___________________

28. In the N2 molecule, each N atom has ___________________ shared pairs of e- and ___________________

unshared pair of e-.

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Test your Understanding of Chemical Bonding

How can the physical properties of substances be explained?

1. Which element is malleable and can conduct electricity in the solid phase?   a. iodine b. phosphorus c. sulfur d. tin

2. Which type of bond is found in sodium bromide?

a. covalent b. hydrogen c. ionic d. metallic 3. Explain, in terms of atomic structure, why liquid mercury is a good electrical conductor. _________

______________________________________________________________________________________

4. Explain, in terms of electronegativity difference, why the bond in H–Cl is more polar than the bond in H–I.

______________________________________________________________________________________

5. Explain, in terms of intermolecular forces, why hydrogen has a lower boiling point than water. ______________________________________________________________________________________

Base your answers to questions 6, 7, and 8 on the information below.      Testing of an unknown solid shows that it has the properties listed below. (1) low melting point (2) nearly insoluble in water (3) nonconductor of electricity (4) relatively soft solid

6. What type of bonding would be expected in the particles of this substance?___________________

7. Explain in terms of attractions between particles why the unknown solid has a low melting point.  ______________________________________________________________________________________

8. Explain why the particles of this substance are nonconductors of electricity.

______________________________________________________________________________________

9. Metallic bonding occurs between atoms of   a. sulfur b. copper c. fluorine d. carbon

10. The high electrical conductivity of metals is primarily due to

a. high ionization energies c. mobile electrons b. filled energy levels d. high electronegativities

11. Which of the following solids has the highest melting point?

a. H2O (s) b. Na2O (s) c. SO2 (s) d. CO2 (s) 12. Which characteristic is a property of molecular substances?

a. good heat conductivity c. low melting point b. high melting point d. good electrical conductivity

13. Which substance contains metallic bonds?

a. Hg (l) b. H2O (l) c. NaCl (s) d. C6H12O6 (s)

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14. True or False: Metallic bonds are found in compounds.

a.  true b. false

15. Which statement describes a chemical property of iron? a. Iron can be flattened into sheets. c. Iron conducts electricity and heat. b. Iron combines with oxygen to form rust. d. Iron can be drawn into a wire.

16. Conductivity in a metal results from the metal atoms having

a. high electronegativity. c. highly mobile protons in the nucleus. b. high ionization energy. d. highly mobile electrons in the valence shell.

17. A substance that does not conduct electricity as a solid but does conduct electricity when melted is

most likely classified as a. an ionic compound. b. a molecular compound. c. a metal. d. a nonmetal.

18. A chemist performs the same tests on two white crystalline solids, A and B. The results are shown in

the table below.

Solid A Solid B

Melting Point 8010 950

Solubility in H2O high low

Electrical Conductivity when melted good nonconductor The results of these tests suggest that

a. both solids contain ionic bonds. b. both solids contain covalent bonds. c. solid A contains covalent bonds and solid B contains ionic bonds. d. solid A contains ionic bonds and solid B contains covalent bonds.

Base your answers to questions 19 and 20 on the table below.

Substance Melting Point (0C) Boiling Point (

0C) conductivity

A -80 -20 none

B 20 190 none

C 320 770 as a solid

D 800 1250 in solution

19. Which substance is an ionic compound?

a. A b. B c. C d. D 20. Which substance is a metal?

a. A b. B c. C d. D 21. Which property did you use to distinguish between ionic and metallic bonding? __________________ 22. Which substance is a pure covalent compound?

a. A b. B c. C d. D 23. Which substance is a polar covalent compound?

a. A b. B c. C d. D 24. Which property(s) did you use to distinguish between the two types of covalent bonding? _________

_________________________________________________________________________________________

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Bringing it all Together!

DIRECTIONS: Use the word banks in the parentheses to answer the questions about ionic, covalent, & metallic bonds. Ionic Covalent Metallic Valence electrons are____________.

(shared, mobile, transferred)

Electrons involved? (metal, nonmetal + nonmetal, nonmetal + metal)

Particles present? cations & electrons, atoms, cations + anions)

Type of attraction? (electrostatic attraction, attraction between metallic particles, shared e- between atoms)

Melting & boiling point? (high, low)

Water solubility? (high, low)

Conductivity as a solid? (good, poor)

Conductivity when melted or aqueous?

(good, poor)

Malleable? (yes, no)

Hard, brittle, or soft? (a bond type can have more than one answers)

Is this type of bond found in compounds?

(yes, no)

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Properties of Ionic and Covalent Compounds Lab

PRE-LAB ASSIGNMENT: Chemical compounds are combinations of atoms held together by chemical

bonds. These chemical bonds are of two basic types—ionic and covalent. Use your knowledge of

ionic and covalent bonds to fill in the following blanks:

Covalent and ionic compounds exhibit different properties. Therefore, the physical properties of a

substance, such as melting point, solubility, and conductivity can be used to predict the type of bond

that binds the atoms of the compound together. Use the “Chemical Bonding” worksheet in your

packet to fill in the following blanks:

Ionic substances:

Covalent molecules:

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Define the following. If needed, use a textbook and/or your notes as a reference. 1. intermolecular forces –__________________________________________________________________

______________________________________________________________________________________

2. electrolyte – ___________________________________________________________________________

______________________________________________________________________________________

3. soluble – ______________________________________________________________________________

4. solute – _______________________________________________________________________________

5. solvent - ______________________________________________________________________________

Table 1: Typical Properties of Ionic and Covalent Compounds

Classification Ionic Covalent

Odor low to none stronger

Texture hard, brittle soft

Solubility in H2O varies varies

Conductivity in solution yes no

Melting Point/Boiling Point high low

In this experiment, you will test eight compounds to determine the above properties. Your compiled data will enable to

you classify the substances as either ionic or covalent compounds. But before you begin predict the type of bond that

exists in each of the following substances.

Name Bond Type (ionic or covalent)

table salt

chalk

sucrose

Advil (C13H18O2)

paraffin wax (approx. C25H52)

milk of magnesia ( Mg(OH)2)

isopropyl alcohol (C3H8O)

ethanol (C2H6O)

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PART A

1. Label five pieces of paper with the following names: table salt, sucrose, Advil, chalk, and wax. Obtain a FEW

CRYSTALS of each chemical and place them on the correctly labeled paper. Take the samples back to your lab

desk.

2. Notice that there is a hot plate for every 3 lab stations. You will have to share! Determine which one your lab

group will use. Do not turn it on yet! Create a “flat-bottom” boat made of aluminum foil that will fit

comfortably on the hot plate. Using the samples that you collected in step 1, place each of the five different

compounds in separate spots on the foil boat--do not allow the samples to touch each other! Label the diagram

below that shows the position of each chemical on the aluminum foil. Keep the labeled pieces of paper to use in

step 7.

3. Place the aluminum boat on the hot plate that has NOT been turned on yet. Make sure the bottom is flat so that

all the substances are exposed to the heat. Recruit two other lab groups to share the hot plate. All three “boats”

will fit!

4. Turn on the hotplate and set the temperature to 1000C. Increase the temperature to the following every five

minutes: 1500C, 2000C, and maximum. Record the order of melting in the data table below. It is not necessary

to have the exact values for the melting points of the solids. The foil will get hotter as it is heated so the order of

melting will give relative melting points. After completing all of the temperature increases, record an “n” in the

date table for each substance that did not melt. Turn off the hot plate and place your foil in the trash can after

it has cooled.

Compound Melting order

table Salt

chalk

sucrose

Advil

paraffin wax

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PART B

5. Repeat step one. In addition to the five solid substances, obtain a small scoop of milk of magnesia.

6. Write a brief description of each of the five substances in the data table. Make special note of the color and

texture of each substance. To assess its texture, obtain a small amount between two fingers. Rinse and dry your

hands between each chemical.

7. Place one of the solid substances into a plastic cup filled with about 25 mL of distilled water. Stir with a glass

rod and test for odor by carefully wafting—do not sniff directly. Observe and record your results in the date

table. In the solubility column, use “s” (soluble) if the substance dissolves and “i” (insoluble) if it does not

dissolve.

8. Use the conductivity tester to determine whether or not the solution conducts electricity. Place only the tip of

the electrodes into the solution. If a bright light is observed then the solution conducts electricity. If a dim light

is observed, the solution does not conduct electricity. Record your results in the data table.

9. Empty the contents down the drain, rinse out the beaker, dry the beaker and electrodes, and repeat steps 7 and

8 with all of the remaining solids, one at a time. Record your observations in the data table.

10. Obtain 10 mL of isopropyl alcohol. Pour it into the plastic cup and describe its appearance. Then test its odor

and conductivity. You do not have to test the texture. Record your observations in the data table.

11. Empty the contents down the drain, rinse out the beater, dry the beaker and electrodes. Repeat step 10 with

ethanol.

12. Clean all equipment, wipe off the lab desk, and wash your hands.

13. Use your data to answer the conclusion questions. Use complete sentences when necessary.

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Names _______________________________________________________________________________ Period __________________

DATA TABLE

Compound

Description

Odor

Texture

Solubility in H2O

Conductivity (in solution)

table Salt

chalk

sucrose

Advil

Paraffin wax

milk of magnesia

isopropyl alcohol

ethanol

CONCLUSION: Answer the following questions. 1. List three properties of ionic and covalent compounds (HINT: Use Table 1 on page two as a guide.)

Ionic: _______________________________________________________________________________________________________ Covalent:____________________________________________________________________________________________________

2. List the five solid substances in order of increasing melting point.

____________________________________________________________________________________________________________

26

3. Based on your melting point data, compare and contrast the strength of covalent and ionic bonds. Be sure to include how you came to your conclusion. ________________________________________________________________________________________________________________ _______________________________________________________________________________________________________________________________ _______________________________________________________________________________________________________________________________

4. Based on your data and the properties of ionic and covalent compounds, make a final decision of what type of bond each is found in the following substances:

Name

Bond Type (ionic or covalent)

table Salt

chalk

sucrose

Advil

Paraffin wax

milk of magnesia

isopropyl alcohol

ethanol

5. If you are faced with two unknown solids in the lab, make a rule that you can follow to determine if a compound is molecular or ionic.

_____________________________________________________________________________________________________________________ _____________________________________________________________________________________________________________________

27

TYPES OF ELECTRON PAIRS In the Lewis structure of a molecule or polyatomic ion, valence electrons occur in pairs. There are two kinds of pairs.

1. A shared pair, or bonding pair, exists ___________________________of the compound. It can be represented by a ______________________________ between the bonded atoms.

2. An unshared pair, or lone pair, belongs_____________________________________________. It is represented by a __________________________________________ on that atom.

Count the lone pairs and shared pairs in the hydroxide ion: [ O — H ] — lone pairs ________ shared pairs ________

Rules for Drawing Lewis Structures It is important to learn to draw Lewis structures. They can help predict the molecular structure and geometry of the molecule.

1. Draw a skeleton structure for the molecule or ion by joining the atoms with single bonds.

a. Hydrogen is always an end or terminal atom. It is connected to only one other atom.

b. The least electronegative atom in the molecule or ion is usually the central atom.

c. Oxygen is NEVER the central atom unless bonded with fluorine or hydrogen.

2. Count the number of valence electrons.

a. For molecules, simply add up the group numbers of the elements.

b. For ions, add to the sum the charge on a negative ion and subtract the charge on a

positive ion.

3. Determine the number of electron pairs that are available to satisfy the octet rule of each

element by dividing the total number of valence electrons in half.

4. Determine the number of electron pairs still available for distribution by subtracting the number

of single bonds (drawn in step 1) from the total number of pairs available (calculated in step 3).

5. Place lone pairs around each terminal atom (except H) to satisfy the octet rule. If pairs are left at

this point, assign them to the central atom. (If the central atom is from the third or higher

period, it can have more than 8 electrons surrounding it. This is called an expanded octet and

will be discussed in greater detail later.)

6. If the central atom is not yet surrounded by 8 electrons, multiple bonds must be formed by

moving one or more terminal lone pairs to a bonding location.

a. If the central atom is two electrons short, move one lone pair to convert a single bond to

double bond.

b. If the central atom is four electrons short, move two lone pairs to convert a single bond

to a triple bond (or two single bonds to two double bonds).

c. Halogens NEVER multiple bond, and if another structure is possible, oxygen will not triple

bond.

∙ ∙ ∙ ∙ ∙ ∙

28

EXAMPLE: Draw the Lewis structure for water. Follow the six steps for writing Lewis structures.

1. Draw a skeleton for the molecule or ion by joining the atoms with single bonds.

2. Count the number of valence electrons.

3. Determine the number of electron pairs.

4. Determine the number of electron pairs still available for distribution.

5. Place lone pairs around terminal atoms. If pairs are left, assign them to the central atom.

6. If the central atom is not yet surrounded by 8 electrons, from multiple bonds.

EXAMPLE: Draw a Lewis structure for carbon dioxide by following the six steps.

29

LEWIS STRUCTURES

DIRECTIONS: Draw Lewis Structures for the following. Show all of your work.

1. AsBr3 2. CSe2

3. ICN 4. NH2-

5. O3 6. BH4-

30

Exceptions to the Octet Rule _______________________________________________________________________________

__________________________________________________________________________________

__________________________________________________________________________________

__________________________________________________________________________________

1. Write Lewis structures for the following ions or molecules, which have central atoms that do not obey the octet rule.

a. BrF5 d. BH2-1

b. ClF3 e. SbCl5

c. SI5

+1 f. IBr2-1

2. Why do phosphorus and sulfur expand their octets in many compounds but nitrogen and

oxygen never do?________________________________________________________________

______________________________________________________________________________

31

IRECTIONS: Draw the Lewis structures for the following molecules or ions. Determine which ones are exceptions to the octet rule and then color or shade in their block. The uncolored blocks will reveal the symbol for an element that is sometimes a culprit for an expanded valence.

Exceptional Element! BCl3

IF4

+1

I3-

BrF3

IF3

SF4

SCl2 CO2 PO4

-3 PF6

-

XeF2

SO4

-2 BeCl2 F2 XeO2F2

ICl3

CS2 PCl3 CCl4 ClF5

KrF4

CO BF3 ClF2

- BeF2

PF5

N2

XeF4 IF4

- SF6

The element is _________________________________________________.

Name ___________________________________________________________________ Period ___________

32

33

Covalent Bonding

1. The bonds between the following pairs of elements are covalent. Circle the more polar bond in each pair.

a. H—Cl or S—Cl b. H—C or H—H c. H—F or H—O 2. Predict what type of bond (ionic, polar covalent, or pure covalent) will form between atoms of the

following elements. If you predict a polar covalent bond, show the polarity of the bond by assigning δ+

and δ- to the appropriate atoms.

a. Ca and Cl ____________________________ b. N and O ____________________________

c. C and S ____________________________ d. H and O ___________________________

e. Mg and F ____________________________ f. S and O ____________________________ 3. Explain why helium is monatomic but hydrogen exists as a diatomic molecule. 4. Define and give an example of the following:

a. a single covalent bond

b. an unshared pair of electrons 5. How many electrons are shared by two atoms in a double covalent bond? ______________

In a triple covalent bond? _______________

6. Draw the Lewis structure for each of the following molecules.

a. bromine monochloride b. ammonium ion

c. HOOH d. SOCl2

34

7. Draw a Lewis structure for each of the following compounds.

a. hydrobromic acid b. phosphorus trichloride

c. sulfate ion d. dichlorine monoxide 8. Draw the Lewis structures for these molecules.

a. nitrogen triiodide b. hydrochloric acid

c. HCCH d. hydrogen cyanide e. nitrogen gas f. oxygen difluoride

9.

a. Draw the Lewis structures for formaldehyde (H2CO) and carbon monoxide.

b. In which molecule is the C - O bond shorter? Why? ____________________________________

__________________________________________________________________________________

c. In which molecule is the C - O bond stronger? Why? ___________________________________

__________________________________________________________________________________

35

Draw the Lewis Structure for sulfur dioxide.

There are two valid structures that can be written for sulfur dioxide. Both indicate that the molecule has one single bond and one double bond. Yet experiments show that there is only one kind of bond in the molecule. An explanation for this phenomenon is that the actual bond is a resonance hybrid, or an average of the valid structures. In this case, the true bond is an intermediate between a single and double bond. The concept of

Resonance

refers to bonding in molecules that cannot be correctly represented by a single Lewis structure. For the electron-dot drawing to comply with experimental observations, equivalent structures called resonance structures are drawn. 1. When drawing resonance structures, there are several things to notice. 2. Resonance structures differ only in location of the bonding and lone electron pairs, not in the location of

the atoms. 3. The resonating bonds in a given molecule or ion are a hybrid of the predicted bonds. These bonds do

not oscillate back and forth. They are a weighted average of the valid resonance structures. If the resonating structures are symmetrical, as in the sulfur dioxide molecule, then the bonds are identical.

4. When halogens occur as terminal atoms, they never multiple bond. 5. If another structure is possible, oxygen will avoid forming a triple bond. Draw all possible Lewis structures for the carbonate ion.

36

1. Since the 3 resonance structures for the ion are symmetrical, then all 3 bonds are _________________.

2. Since the 3 bonds are identical, they are a _____________________ average of the ______ single bonds

and _______ double bonds present in the ion, making the bond length closer to that of a

_____________________ bond.

3. The bond is (weaker/stronger) than a single bond but (weaker/stronger) than a double.

Using Lewis structures draw the resonance structures for both the nitrite and nitrate ions. Which ion contains the longer bonds? Explain your answer. Compare the nitrogen-oxygen bond lengths in the nitronium ion, NO2

+1, and in the nitrite ion. Which ion has the

37

Name _________________________________________________________________ Period ____________

Lewis Structure Word Scramble

DIRECTIONS: Choose the correct answer for each question, and record the letter of that answer at the end of the puzzle. You

must draw the Lewis structures for questions 2-18. Unscramble the letters to determine the two word term related to writing

Lewis structures. The answers to questions 1-9 will be used to form the first word of the two word answer, while the answers

to questions 10-18 will form the second word.

1. Which of the following is the correct Lewis structure for hydrochloric acid? A. H – Cl B. H – Cl C. H – Cl

2. Which of the following has resonance structures?

A. SO3 B. NH4+ C. CO

3. Which of the following contains a multiple bond? P. H2 Q. F2 R. N2

4. Which of the following contains one lone pair of electrons? M. HBr N. NH3 O. SCl2

5. Which of the following contains only bonding pairs of electrons? D. CCl4 E. SiH4 F. H2S

6. Which of the following contains only single bonds? S. CH3I T. HCCH U. NO2

-

7. Which diatomic molecule contains a double bond? D. Cl2 E. O2 F. H2

8. Which of the following molecules contains no multiple bonds? M. O3 N. CSe2 O. CH4

9. Which of the following contains two lone pairs and two bonding pairs of electrons? M. H3O

+ N. H2O O. OH

-

●●

●●

●●

●●

●●

●●

●●

●●

●●

●●

●●

38

10. The C-N bond in the cyanide ion is a ________________ bond. S. double T. single U. triple bond

11. Which of the following contains the most lone pairs of electrons? S. PO4

-3 T. BH4

- U. HF

12. Which Lewis structure contains a double bond? S. NH3 T. ClO4

- U. CO3

-2

13. Which of the following has three resonance structures? S. PO3

-3 T. NO3

- U. SO4

-2

14. Which of the following contains a triple bond? D. PCl3 E. ICN F. Br2

15. Which central atom has no lone pairs of electrons? P. AsF3 Q. NF3 R. PCl5

16. Which Lewis structure contains six lone pairs of electrons? A. HI B. SCl4 C. SO2

17. Which Lewis structure contains the greatest number of lone pairs? S. XeF2 T. SiCl4 U. IF3

18. Which of the following has resonance structures? P. SF4 Q. CO2 R. SeO3

Letters in the first word: ______ ______ ______ ______ ______ ______ ______ ______ ______

Letters in the second word: ______ ______ ______ ______ ______ ______ ______ ______ ______

Unscrambled Lewis structure term: ___________________________ ___________________________

39

Name ____________________________________________________ Period ___________________

IONIC & METALLIC BONDING - Vocabulary Review

Match the correct vocabulary term to each numbered statement. Write the letter of the correct term on the line. Each answer can only be used once.

a. malleable j. octet rule b. anion k. chemical formula c. ionic bond l. metal d. pure covalent bond m. nonmetal e. polar covalent bond n. ionic compound f. valence electrons o. metallic bond g. Lewis dot structure p. ductile h. electrostatic attraction q. cation i. resonance r. expanded octet

___________ 1. a bond formed between a cation and an anion

___________ 2. the attraction of free-flowing valence electrons for positively charged metal

cations

___________ 3. negatively charged ions

___________ 4. positively charged ions

___________ 5. a diagram that shows valence electrons as dots

___________ 6. positive and negative ions are held together by this type of attraction

___________ 7. type of bond formed by the equal sharing of electrons

___________ 8. electrons in the highest occupied energy level of an atom

___________ 9. type of bond formed by the unequal sharing of electrons

___________10. are good conductors of heat and electricity only when melted or in an

aqueous solution

___________11. are good conductors of heat and electricity

___________12. are poor conductors of heat and electricity

___________13. when forming compounds, atoms tend to react so as to acquire the stable

electron configuration of a noble gas ; 8 valence electrons

___________14. the lowest whole number ratio of ions in an ionic compound

___________15. the ability to be hammered into thin sheets

___________16. the ability to be drawn into wire

___________17. bonding in molecules cannot be correctly represented by a single Lewis dot

structure

___________18. period 3 elements and beyond can have more than 8 valence electrons

because they have empty “d” sublevels for excited electrons to jump into

40

41

BONDING

REVIEW

PART A: Fill in the word(s) that will make each statement true. 1. The electrons in the highest occupied energy level of an atom are called the _____________ electrons. 2. The _______________ rule states that atoms desire to have eight electrons in their outer energy level 3. An oxygen atom attains a stable electron configuration by __________ 2 electrons. 4. Ionic compounds conduct electricity when they are in the ____________ state. 5. When atoms share electrons to gain the electron configuration of a noble gas, the bonds formed are

called ____________________.

6. ___________ tend to lose electrons when they react to form compounds. 7. Ionic bonds are held together by a/an _________________ __________________. 8. Metals are ductile because they have free moving _________ that allow the particles to slide over one

another. 9. Halogens, when terminal atoms, never form ______________ _______________. 10. During the formation of the compound NaCl, _______ electron is transferred from a sodium atom to a

chlorine atom. 11. Among the representative elements, the group number of each element is equal to the number of

______________ electrons in an atom of that element.

12. Most covalent compounds are composed of two or more _______________. 13. Molecular compounds tend to have ___________ melting and boiling points, while ionic compounds

tend to have _____________ melting and boiling points.

14. When it is possible to write two or more valid Lewis structures for a molecule or ion, each formula is

referred to as a _____________ _________________. 15. When identical atoms are joined by a covalent bond, the bonding electrons are shared

_________________, and the bond is ___________. When atoms in a bond are not the same, the

bonding electrons are shared ___________, and the bond is _______.

16. A ____________ covalent bond exists when three pairs of electrons are shared by two bonded atoms.

42

PART B: Answer the following in the space provided. 17. Predict whether the following compounds contain polar covalent bonds, pure covalent bonds, or

ionic bonds.

a. KF _____________________ c. MgCl2 _____________________

b. SO2 _____________________ d. Cl2 _____________________ 18. Using only the periodic table, predict which bond in each of the following groups will be most polar.

a. C-F, Si-F, Ge-F _______________ c. S-F, S-Cl, S-Br _______________

b. P-Cl, S-Cl _______________ d. C-H, Si-H, Sn-H _______________ 19. Name the two elements that can have a deficient octet (less than 8 electrons in its outer energy level). Which elements can have an expanded octet? Why? 20. Compare the strength of single, double, and triple bonds. 21. Compare the length of single, double, and triple bonds. PART C. Vocabulary Review ~ match the correct vocabulary term with the correct statement. 22. _____ compounds composed of cations and anions A. malleability 23. _____ the ability to dissolve in a solvent B. valence electrons 24. _____ the attraction of free-flowing valence electrons for C. Lewis dot structure positively charged metal ions D. octet rule 25. _____ the electrostatic attraction that binds oppositely E. halide ion charged ions together F. ionic compounds 26. _____ negatively charged ions G. conductivity 27. _____ a diagram that shows valence electrons as dots H. alloy 28. _____ a negative ion formed when a halogen atom gains an e- I. metallic bonds 29. _____ when forming compounds, atoms tend to react in order to J. solubility acquire the stable electron configuration of a noble gas K. anions 30. _____ a mixture of two or more elements, at least one of L. cations which is a metal M. ionic bond 31. _____ positively charged ions 32. _____ ability to carry an electrical current when e- or ions are free to move 33. _____ the ability of solids to be hammered into thin sheets 34. _____ the electrons in the outermost s and p sublevels

43

Bonding Standardized Test Prep

DIRECTIONS: Select the choice that best answers each question or completes each statement. (1) Which of these is not an ionic compound?

A. KF B. Na2SO4 C. SiO2 D. Na2O

(2) Which statements are correct when barium and oxygen react to form an ionic compound?

I. Barium atoms lose 2 electrons and form a cation. II. Oxygen atoms form oxide anions (O2−). III. In the compound the ions are present in a one-to-one ratio.

A. I and II only B. II and III only C. I and III only D. I, II, and III

(3) How many valence electrons does arsenic have?

A. 5 B. 4 C. 3 D. 2

(4) For which compound name is the incorrect formula given?

A. magnesium iodide, MgI2 B. potassium selenide, K2Se C. calcium oxide, Ca2O2 D. aluminum sulfide, Al2S3

(5) Which electron configuration represents a nitride ion?

A. 1s22s23s24s2 B. 1s22s22p3 C.1s22s22p6 D. 1s2

(6) When a bromine atom gains an electron

A. a bromide ion is formed. B. the ion formed has a -1 charge. C. the ion formed is an anion. D. all the above are correct.

The lettered choices below refer to Questions 7–10. A lettered choice may be used once, more than once, or not at all.

A. gains two electrons B. loses two electrons C. gains three electrons D. loses one electron E. gains one electron

Which choice describes what happens as each of the following elements forms its ion? (7) iodine _________ (9) cesium _________ (8) magnesium _________ (10) phosphorus _________ (11) Which substance contains both covalent and ionic bonds?

A. NH4NO3 B. CH3OCH3 C. LiF D. CaCl2

(12) Which of these bonds is most polar?

A. H—Cl B. H—Br C. H—F D. H—I

44

(13) How many valence electrons are in a molecule of phosphoric acid, H3PO4?

A. 7 B. 16 C. 24 D. 32

(15) True or False. The nitrate ion has three resonance structures because the nitrate ion has three single bonds. Include Lewis Dot Structures to explain your answer.

__________________________________________________________________________________ __________________________________________________________________________________ (16) Which of the following gases found in earth’s atmosphere would you predict is a pure

covalent compound?

I. Oxygen II. Hydrogen III. Nitrogen IV. Helium A. I only B. I and II only C. IV only D. I, II, and III

45

WHAT dO I need to know??

Unit 10: Bonding

PART I: Scantron

Multiple choice o Lewis structures ~ you will have to

draw several in order to answer the questions. The structures are worth points!

o Definitions: polarity ionic bond polar vs. pure covalent bond metallic bond valence electrons shared pair of electrons lone pair of electrons resonance electronegativity single, double, triple bonds

o strength of single, double, and triple bonds o expanded octets o deficient octets o resonance structures o electronegativity trend o physical properties of ionic, covalent, and metallic

bonds (i.e. melting points, boiling points, hardness, conductivity, solubility, etc.)

46

True & False o conductivity of metals o states of matter (aqueous vs. liquid) o resonance structures o physical properties of ionic, covalent, and metallic

bonds

Matching o choices include: ionic, metallic, pure covalent,

polar covalent match the correct type of bond with the

appropriate statement Example: bond that is formed by an

electrostatic attraction between ions = ionic

47

VSEPR Theory Valence Shell Electron Pair Repulsion theory

There are three types of repulsion between the electron pairs:

Bonding Pair – Bonding Pair: least repulsion Lone Pair – Bonding Pair: more repulsion Lone Pair – Lone Pair: greatest repulsion

This strength of the repulsion helps to determine how a molecule will adjust its shape to minimize the electron repulsion.

In order to determine the shape of the molecule, spread the bonds as far apart as possible. Since the central atom determines the shape of the molecule, count the number of electron locations around the central atom to determine the maximum degree of separation.

Example: CO2: a molecule with two bonding locations around the central atom. 1. Draw the Lewis structure: 2. If you follow the rule and spread the bonds as far apart as possible, which one is the correct shape? Example BH3: a molecule that has three bonding locations around the central atom. 1. Draw the Lewis structure. 2. Once again, follow the rule and spread the bonds as far apart as possible. Which one is the correct shape? Example CH4: a molecule with four bonding locations around the central atom. 1. Draw the Lewis structure. 2. This is the first example where the structure is not planar but three-dimensional, so the bond angle is not 900, as I am sure you drew. What should the correct shape be?

Because electrons naturally repel one another, molecules will adjust their shape so

that the valence electron pairs around the central atom are as far apart as possible.

48

49

Shapes of Covalent Molecules

SHAPE UP!

50

51

VSEPR Pre Lab Assignment

FORMULA

LEWIS STRUCTURE

# BONDING

LOCATIONS

ON

CENTRAL

ATOM

# LONE

PAIRS ON

CENTRAL

ATOM

TOTAL # OF

ELECTRON

LOCATIONS

3d SKETCH WITH

BOND ANGLES

VSEPR SHAPE

& NAME

SO2

NH3

IF5

NAME __________________________________________________________________ PERIOD _______________

52

53

Total e-

locations on central atom

Bonding

locations on central atom

Lone

pairs on central atom

Shape

Bond Angles

Example

1.

2 0

2.

3 0

3.

2 1

4.

4 0

5.

3 1

6.

2 2

7.

5 0

8.

4 1

9.

3 2

10.

2 3

11.

6 0

12.

5 1

13.

4 2

= +

BONDING SHAPES

54

Determining the Polarity of Molecules

For a molecule to be POLAR:

1. It must have at least one polar bond. For our purposes, assume any bond between two different nonmetals to be polar.

2. It must be nonsymmetrical so the polar bonds do not cancel each other out.

The following are guidelines to follow when determining if a molecule is a polar compound:

These apply ONLY to substances where the central atom

is bonded to identical terminal atoms. 1. If there are NO LONE PAIRS on the central atom, it is NONPOLAR.

- These shapes are all symmetrical and the polar bonds would cancel each other out. 2. If there are LONE PAIRS on the central atom, it is POLAR.

- EXCEPTIONS: square planar and linear with lone pairs; since their shapes are symmetrical, they are nonpolar.

**When the molecule has different terminal atoms, it loses its symmetry, making the molecule polar.**

55

Draw a Lewis structure for each of the following molecules or ions. Using this structure along with the VSEPR Theory, fill in the chart for each compound or ion.

substance Lewis structure bonding

locations

lone

pairs

shape bond angle polar?

(Y/N)

PF6-1

OSF4

BeCl2

ClO3-1

NH2-1

Molecular Shapes of Compounds

56

substance Lewis structure bonding

locations

lone

pairs

shape bond angle polar?

(Y/N)

S2O3-2

BO3-3

BrF3

XeF4

IF5

O3

57

1200

MOLECULAR GEOMETRY REVIEW 1. State the basic assumption of the VSEPR theory. 2. True or False. If a molecule possesses polar bonds, then it means that the molecule is polar. Explain your answer. 3. Do the following molecules have a dipole moment (are they polar)?

a. CH2Cl2 _________ d.) CCl4 _________

b. CHCl3 _________ e.) CO2 _________

c. N2O _________

4. Label the following models with the correct molecular geometry:

a.) b.) c.) d.) __________________ __________________ __________________ __________________ e.) f.) g.) h.) __________________ __________________ __________________ __________________ i.) j.) k.) _________________ __________________ __________________

109.50

900

1200

1200

900

900

< 109.50 < 120

0

900

< 900

< 1200

< 900 90

0

< 900

58

5. Draw Lewis structures for the following molecules or ions in a - h. Then describe, 1) the shape, 2) bond angle, & the 3) polarity of each. Show all resonance structures where applicable.

a. XeOF2 1. _____________________________

2. _____________________________

3. _____________________________

b. sulfate ion 1. _____________________________

2. _____________________________ 3. _____________________________

c. beryllium diiodide 1. _____________________________

2. _____________________________ 3. _____________________________

d. OCN- 1. _____________________________

2. _____________________________ 3. _____________________________

59

e. selenium tetrabromide 1. _____________________________

2. _____________________________ 3. _____________________________

f. sulfite ion 1. _____________________________

2. _____________________________ 3. _____________________________

g. SiSCl2 1. _____________________________

2. _____________________________ 3. _____________________________

h. POCl3 1. _____________________________ 2. _____________________________ 3. _____________________________

60

61

WHAT dO I need to know??

Unit 11: VSEPR Theory

PART I: Scantron

o 25 questions multiple choice o Lewis structures ~ you will have to

draw several in order to answer the questions

(They will be worth points!)

o VSEPR Shapes and bond angles o Definitions:

VSEPR lone pair, bonding pair single bond, double bone resonance polar vs. nonpolar repulsions & their strengths

o lone pair vs. lone pair o loan pair vs. bonding pair o bonding pair vs. bonding pair

o expanded octets o deficient octets

62