Unit 11: Ionic & Metallic Bonding

Home , Metallic bonding

(Chapter 7)

Unit 11: Covalent Bonding

(Chapter 8)

2 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? ______

3 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?

______

4 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? ______

6 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. ______

7 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.

8 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 ______.

Key Points  ______

 ______

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. Key Points

 ______

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.

9 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 _____. Key Points

 ______

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.

10 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 Electronegativity Actual Bond More Bond Type Difference Type Negative Atom

1. CaO ______2. KI ______

3. F2 ______

4. N2O ______

5. SnH2 ______

6. CH4 ______

7. SbCl5 ______8. PbS ______

9. As2O3 ______

10. Mg3N2 ______Electronegativity Values 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? ______

12 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.

13 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______.

14 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 ______.

15 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-.

16 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)

17 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? ______

______

18 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)

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:

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)

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

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.

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.

______

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.

______

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.

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.

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

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.

-1 a. BrF5 d. BH2

b. ClF3 e. SbCl5

+1 -1 c. SI5 f. IBr2

2. Why do phosphorus and sulfur expand their octets in many compounds but nitrogen and oxygen never do?______

______

30 Name ______Period ______

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! +1 - BCl3 IF4 I3 BrF3 IF3

-3 - SF4 SCl2 CO2 PO4 PF6

-2 XeF2 SO4 BeCl2 F2 XeO2F2

ICl3 CS2 PCl3 CCl4 ClF5

- KrF4 CO BF3 ClF2 BeF2

- PF5 N2 XeF4 IF4 SF6

The element is ______.

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

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? ______

______

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.

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.

+1 Compare the nitrogen-oxygen bond lengths in the nitronium ion, NO2 , and in the nitrite ion. Which ion has the

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

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? -3 - S. PO4 T. BH4 U. HF

12. Which Lewis structure contains a double bond? - -2 S. NH3 T. ClO4 U. CO3

13. Which of the following has three resonance structures? -3 - -2 S. PO3 T. NO3 U. SO4

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: ______

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

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.

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

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

(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

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.)

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

VSEPR Theory Valence Shell Electron Pair Repulsion theory 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. 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?

Shapes of Covalent Molecules SHAPE UP!

50 NAME ______PERIOD ______

VSEPR Pre Lab Assignment # BONDING LOCATIONS # LONE TOTAL # OF FORMULA LEWIS STRUCTURE ON PAIRS ON ELECTRON 3d SKETCH WITH VSEPR SHAPE CENTRAL CENTRAL LOCATIONS BOND ANGLES & NAME ATOM ATOM

SO2

NH3

IF5

52 BONDING SHAPES

Total e- Bonding Lone locations on = locations on + pairs on Shape Bond Angles Example central atom central atom central atom

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

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.**

Molecular Shapes of Compounds

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 lone shape bond angle polar? locations pairs (Y/N)

-1 PF6

OSF4

BeCl2

-1 ClO3

-1 NH2

55 substance Lewis structure bonding lone shape bond angle polar? locations pairs (Y/N)

-2 S2O3

-3 BO3

BrF3

XeF4

IF5

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:

900 a.) b.) c.) 0 d.) 1200 109.5 1200 1200

______

0 90 0 e.) 90 0 f.) g.) h.) 90 0 0 < 109.5 < 120

______

0 0 0 90 < 90 < 90 0 i.) j.) < 120 k.) < 900

______

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. ______

58 e. selenium tetrabromide 1. ______

2. ______

3. ______

f. sulfite ion 1. ______

2. ______

3. ______

g. SiSCl2 1. ______

2. ______

3. ______

h. POCl3 1. ______

2. ______

3. ______

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