name: honors unit 2 class: bonds, james bonds · 10/2/2019 · 1. what is the electron...

Honors Unit 2 Bonds, James Bonds

Topics/ Daily Outline:

Day A B Content: TEXT CW #: HW #: 1 10/14 10/16 Table of Elements 6.1, 6.2 1, 2 --

2 10/17 10/18 Electron configuration, Periodic trends 5.2, 6.3 3, 4 1

3 10/21 10/22 Chemical bonds and properties 7, 8 5, 6 -- 4 10/23 10/24 Covalent bonds, Molecular naming 8.1, 8.2, 9.3 7, 8 2

5 10/25 10/28 Molecular Lewis dot structures 8.2, 8.3 9 3

6 10/29 10/30 Intermolecular forces 8.4 10 4 7 10/31 11/1 Review -- -- --

8 11/4 11/5 Quarterly Assessment -- -- --

9 11/6 11/7 Ionic bonds 7.1, 7.2 11, 12 5 10 11/8 11/11 Ionic naming 9.1, 9.2 12, 13 --

11 11/12 11/13 Acid and bases naming, Metallic bonds 9.4, 7.3 14, 15 6

12 11/14 11/15 Review -- -- -- 13 11/18 11/19 Unit Test -- -- --

Homework:

• HW 1: Electron Configuration

• HW 2: Molecular Compounds

• HW 3: Intermolecular Forces Notes

• HW 4: Review for the Quarterly Assessment

• HW 5: Ionic Compounds

• HW 6: Review for Unit Test Important Due Dates:

• Who Tagged the Lab Bench Conclusion, 10/17 (A Day) and 10/18 (B Day)

• SciResearch: 6 Develop a Procedure, 10/21 (A Day) and 10/22 (B Day)

• SciResearch: 7 Get Approval to Start, 10/29 (A Day) and 10/30 (B Day)

For tutorials and additional resources: www.leffellabs.com

If you are absent, please use this sheet to determine what you missed and collect the materials from the make-up work bins up front. Get help from a friend, the links above, or the instructor.

Name: Class:

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Drills

Date: Outcome:

Date: Outcome:

Date: Outcome:

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Date: Outcome:

Date: Outcome:

Date: Outcome:

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CW 1: Mendeleev’s Table

Video Link: http://www.youtube.com/watch?v=fPnwBITSmgU

1. Was Mendeleev the only person to sort the elements into lists by similar properties?

2. What is the significance of the dashes Mendeleev placed on his periodic table?

3. How could Mendeleev predict the properties of eka-aluminum?

4. What is the relationship between gallium and eka-aluminum?

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CW 2: Properties of the Periodic Table

Go to http://chemfiesta.wordpress.com/2014/09/27/periodic-table-and-trends/ Complete the following sections by reading along on the webpage.

What’s the Periodic Table?

1. Color code and label each section of the periodic table.

Main block/group elements Inner transition metals Outer transition metals

2. Which section corresponds to the representative elements?

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Other Sections of the Periodic Table

3. Color code and label each group on the periodic table.

---- Hydrogen*

Alkali Metals*

Alkali Earth Metals*

Transition Metals*

Lanthanides (Rare Earth Metals)

Actinides

Those Other Ones

Pnictogens

Chalcogens

Halogens*

Noble Gases*

* Important ones that you should memorize

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4. Complete the table below, listing the properties of each group. Group # / Name Properties

1: Alkali Metals

2: Alkali Earth Metals

3-12: Transition Metals

18: Noble Gases

17: Halogens

Hydrogen

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Metals, Nonmetals, and Metalloids

5. Color code and label the metals, nonmetals, and metalloids on the periodic table.

Metals Metalloids Nonmetals

6. What are the properties of metals?

7. How does delocalized bonding explain the properties of metals?

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8. What are the properties of nonmetals?

9. How does localized bonding explain the properties of nonmetals?

10. Metalloids exhibit a combination of localized and delocalized bonding. How does this relate to the properties of metalloids?

Journal Write 1

Create a table to compare the type of bonds in metalloids to those in metals (delocalized) and nonmetals (localized).

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CW 3: Electron Configuration

Questions Notes

Explain the atom according to the quantum mechanical model.

What is an atomic orbital?

How does the electron configuration relate to the arrangement of electrons in atoms?

S ---- ---- p ---- ---- d ---- ---- f ----

----

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1. What is the electron configuration of the following? a. Magnesium

b. Iron

c. Krypton

d. Rubidium

2. For each of the following, write the electron configuration, then draw electrons in the

appropriate energy level of the Bohr model of the atom.

Ca: Number of valence electrons:

Ca+2: Number of valence electrons:

S: Number of valence electrons:

S-2: Number of valence electrons:

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CW 4: Trends in the Periodic Table

1. Take notes in the table below.

Property Definition Trend Across a

Row (left to right) Trend Down a

Column

Atomic Radius

Ionization Energy

Electronegativity

Metallic Character

2. For each of the following, circle the element which has the highest:

a. Atomic Radius: F, Cl, Br

b. Ionization Energy: K, Ca, Sc

c. Electronegativity: C, Si, P, N

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Shielding Effect

As you move down a family, there are more energy levels occupied by electrons. This is because the additional layers of electrons block the outer electrons from the nucleus.

Li: 1s2 2s1 Na: 1s2 2s2 2p6 3s1 K: 1s2 2s2 2p6 3s2 3p6 4s1

3. Consider the valence (2s1) electron of lithium, the valence (3s1) electron of sodium and

the valence (4s1) electron of potassium. Which would experience a greater shielding effect?

4. How does shielding affect the strength of attraction between the nucleus and the valence electrons?

5. If the attraction between nucleus and valence electron is weak, would that electron be found closer to the nucleus, or farther away? Explain.

6. How does more shielding affect the following? a. Atomic radius

b. Ionization Energy

c. Electronegativity

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Nuclear Charge

As you move across a period, the number of protons in the nuclei of the atom increases. This results in a larger positive charge of the nucleus.

Atom K Ca Sc Ti V

# of p+

7. Consider the atoms above. Which would have a larger positive nuclear charge?

8. How does nuclear charge affect the strength of attraction between the nucleus and the valence electrons?

9. If the attraction between nucleus and valence electron is strong, would that electron be found closer to the nucleus, or farther away? Explain.

10. How does a more positive nuclear charge the following? a. Atomic radius

b. Ionization Energy

c. Electronegativity

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Journal Write 2

Cut and paste the table of elements below into your journal. Then answer the questions that follow.

1. Fill in the blanks on the table of elements. 2. How does increasing shielding affect the attraction between the nucleus and electrons? 3. How does increasing nuclear charge affect the attraction between the nucleus and

electrons?


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--Page Intentionally Left Blank--

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CW 5: What Kind of Bond Will Form?

1. Define electronegativity. (HINT: CW 4)

Electronegativity Scale (Pauling Scale)

Here are the electronegativity values for many elements. In general, the closer the element is to fluorine on the periodic table, the more electronegative the element is. The closer the element is to francium on the periodic table, the less electronegative the element is.

Electronegativity Difference

Based on the electronegativity difference between two elements, we can predict what type of bond will form. In general, a larger difference behaves more like ionic compounds (electrons and transferred from one element to another), while a smaller difference behaves like polar bonds (electrons shared, but unequally). The closer the difference is to zero, the more nonpolar (electrons shared equally) the bonds become.

𝐸𝑙𝑒𝑐𝑡𝑟𝑜𝑛𝑒𝑔𝑎𝑡𝑖𝑣𝑖𝑡𝑦 𝐷𝑖𝑓𝑓𝑒𝑟𝑒𝑛𝑐𝑒= |𝐸𝑙𝑒𝑐𝑡𝑟𝑜𝑛𝑒𝑔𝑎𝑡𝑖𝑣𝑖𝑡𝑦𝑒𝑙𝑒𝑚𝑒𝑛𝑡 𝐴 − 𝐸𝑙𝑒𝑐𝑡𝑟𝑜𝑛𝑒𝑔𝑎𝑡𝑖𝑣𝑖𝑡𝑦𝑒𝑙𝑒𝑚𝑒𝑛𝑡 𝐵|

2. Use the information below and the electronegativity value table on the previous page to complete the chart below.

Electronegativity and Bond Types The electronegativity gives us a clue about the type of bond that will form. It is also important to consider the identity of the elements that are forming a chemical bond.

Difference in electronegativity Most probable bond type Example

0.0 – 0.4 Covalent – Nonpolar H – H (0.0) 0.4 – 1.0 Covalent – Moderately Polar H – Cl (0.9)

1.0 – 2.0 Covalent – Very Polar H – F (1.9) ≥ 2.0 Ionic Na+Cl– (2.1)

First Element

Metal or Nonmetal?

1st Element Electroneg.

Second Element

Metal or Nonmetal?

2nd Element Electroneg.

Electroneg. Difference

Bond Type

Ca

Metal

F

4.0

Ag

1.9

Au

Metal

Metallic

H

Nonmetal

Cl

3.0

Au

2.4

Cu

Metal

Metallic

H

Nonmetal

H

2.1

Mg

1.2

Cl

Nonmetal

H

Nonmetal

O

3.5

3. Describe the electronegativity difference and the types of element (metals or nonmetals) involved in each of the following types of bonds.

Bond Type Electronegativity Difference Types of Elements Involved

Ionic

Polar Covalent

Nonpolar Covalent

Metallic

4. Use the periodic table below.

a. Find and highlight Na and Cl in yellow.

b. Find and highlight C and O in green.

c. Which pair would form an ionic bond?

d. Which pair would form a covalent bond?

e. Considering the position of the elements you highlighted, which pair of elements would you expect to have a large electronegativity difference? Explain.

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5. Complete the table below.

Element Electron Configuration # Valence Electrons

Electronegativity Value (Page 7)

Na

F

6. How many valence electrons are needed for a stable electron configuration?

7. Fluorine has 7 valence electrons. How does this explain the very high electronegativity value of fluorine?

8. How many valence electrons would sodium have if it lost its 3s1 electron, forming the cation Na+1? Is this configuration more stable, or less stable?

Na: 1s2 2s2 2p6 3s1

Journal Write 3

Considering your answer to Question 8, why does sodium have such a low electronegativity value?


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CW 6: Bonding Types Virtual Lab

Go to LEFFELlabs/ unit 2 to access the virtual lab.

1. Review the formation of each bond type. Summarize below. Include a labeled picture for each summary.

a. Ionic

b. Metallic

c. Covalent (explain the difference in electron sharing for polar and nonpolar) i. Polar

ii. Nonpolar

2. Explain the relationship between covalent bond and molecular compound.

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3. Complete the table below by summarizing and explaining each property. Property Ionic Metallic Covalent: Polar Covalent:

Nonpolar Dissolves in a polar solvent (water)

Dissolves in a nonpolar solvent (oil)

Conduct in the solid state

Conducts in the aqueous state

Melting point/ Boiling point

Malleability

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CW 7: Covalent Bonds Gizmo

Vocabulary: covalent bond, diatomic molecule, Lewis diagram, molecule, noble gases, nonmetal, octet rule, shell, valence, valence electron Prior Knowledge Questions (Do these BEFORE using the Gizmo.) 1. There are eight markers in a full set, but Flora and Frank each only have seven markers.

Flora is missing the red marker, and Frank is missing the blue marker.

What can they do so that each has a full set of markers?

2. Otto and Olivia each have six markers. Otto is missing the purple and green markers, and Olivia is missing the black and brown markers. What can they do so that each has a full set?

Gizmo Warm-up Just like the students described above, nonmetal atoms can share electrons. As you will see in this gizmo, atoms form bonds in this way. To begin, check that Fluorine is selected from the Select a substance menu. Click Play ( ) to see the electrons orbiting the nucleus of each atom. 1. The outermost electrons in each atom are called valence electrons. How many valence

electrons does each fluorine atom have? ________________________________________

2. Click Pause ( ). Drag an electron from the left atom to the right atom. Click Play.

What happens? ____________________________________________________________ 3. Click Pause, drag an electron from the right atom to the left, and then click Play.

What happens now? ________________________________________________________

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Activity A: What happens when atoms share electrons?

Get the Gizmo ready:

• Click Reset.

• Select Hydrogen. Introduction: The electrons that orbit the nucleus of an atom are arranged into shells. The first shell contains up to two electrons and the second contains up to eight electrons. Most elements are stable when they have eight valence electrons—a rule of thumb known as the octet rule. (Elements with less than five electrons are stable with two valence electrons.)

4. Predict: Each hydrogen atom has one valence electron, but it needs two electrons to be stable. How can both hydrogen atoms each achieve a stable configuration?

5. Form a bond: Drag the electrons so that they move around both hydrogen atoms. Click

Play to observe them in orbit, and then click Check. You have created a covalent bond. Because the molecule has two atoms, it is a diatomic molecule.

Why is this configuration stable for hydrogen?

6. Draw a diagram: Covalent bonds are shown in Lewis diagrams. In a Lewis diagram, dots

represent unshared valence electrons and dashes represent pairs of shared electrons.

Turn on Show Lewis diagram. What is the Lewis diagram for hydrogen, H2? H H

7. Form a bond: Now select Fluorine and using the same procedure as above, create a molecule of fluorine, F2.

What is the Lewis diagram for fluorine, F2? F F

8. Think and discuss: How many electrons are represented by the line between each atom?

What do the other dots around each atom represent?

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Activity B: How do atoms share more than one pair of electrons?

Get the Gizmo ready:

• Click Reset.

• Turn off Show Lewis diagram.

• Select Oxygen.

9. Observe: Like fluorine and most other elements, oxygen atoms are most stable with a

full complement of eight valence electrons.

A. How many valence electrons does each oxygen atom have now? _______________ B. How many more electrons does each oxygen atom need to be stable? ___________

10. Form a bond: Drag electrons back and forth until the molecule of oxygen (O2) is stable.

Click Check to confirm your molecule is stable.

How many pairs of shared electrons are there in a stable molecule of oxygen? __________ How many lines are needed to represent the pairs? _________

11. Draw a diagram: Draw a Lewis diagram of the oxygen molecule in the space below at

left. To check your work, turn on Show Lewis diagram. Draw the correct diagram on the right.

Practice diagram: O O Actual: O O

12. Practice: Create covalent bonds and stable molecules for the remaining substances.

Draw Lewis diagrams for each one. (As above, draw the diagram on your own before checking your work.)

Water H O H

Carbon dioxide O C O

H

Ammonia H N H H Methane H C H H

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13. Count: Review the Lewis diagrams you drew on the previous page. Note that each element tends to form a certain number of chemical bonds. This value is the valence of the element.

For each element in the table below, use the Gizmo to find the number of valence electrons and the list the valence based on the Lewis diagram. Then find the sum of these numbers.

Element Symbol # of valence

electrons Valence Sum

Fluorine F

Hydrogen H

Oxygen O

Nitrogen N

Chlorine Cl

Carbon C

Silicon Si

14. Make a rule: If you knew the number of valence electrons in a nonmetal atom, how

would you determine the valence of the element? (Hint: Ignore hydrogen for now.)

15. Analyze: The first shell can hold a maximum of two electrons. How does this explain the valence of hydrogen?

16. Apply: Selenium has six valence electrons. What is the valence of selenium?

17. Think and discuss: The last column of the periodic table contains the noble gases,

elements that do not easily form chemical bonds. Why don’t these gases tend to form chemical bonds?

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CW 8: Naming Molecular Compounds

The following table shows the names of several molecular compounds.

NO Nitrogen monoxide NF3 Nitrogen Trifluoride NO2 Nitrogen dioxide SF6 Sulfur Hexafluoride

N2O4 Dinitrogen Tetraoxide PCl5 Phosphorus Pentachloride

1. Based on the examples above, what kinds of atoms form molecular compounds

(Metals? Nonmetals? Both?)

2. What do the following prefixes mean?

3. Based on the given examples, write a set of rules that can be used to write the formulas of molecular compounds given their name.

4. Use the rules that you wrote to write formulas for the following.

Name Formula

Carbon tetrachloride

Dihydrogen monosulfide

Phosphorus triiodide

Sulfur dibromide

Boron trifluoride

Nitrogen

Tetraphosphorus decaoxide

Sulfur hexafluoride

Mono ---------- Tetra ---------- Di ---------- Penta ----------

Tri ---------- Hexa ----------

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Rules for Naming Molecular Compounds

• Write the names of the elements in the order listed in the formula

• Use prefixes appropriately to indicate the number of each kind of atom

• Never start with “mono”

• Remove the ending of the second element and add the suffix “- ide” instead

5. Name the molecular compounds below using the rules above. Formula Name Formula Name

N2H4 -------_______--___________---- SBr2 -------_______--___________----

XeF4 P4O3

NH3 BCl3

CO2 H2O

Journal Write 4

Rewrite and correct any formulas you got wrong in Question 4.


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CW 9: Lewis Dot Structures (Molecular)

Bonding Basics

1. Write the number of valence electrons above each column.

2. How many electrons are shared in a single bond?

3. How many electrons are shared in a double bond?

4. How many electrons are shared in a triple bond?

5. What type of bond is always formed with hydrogen?

6. What is the main difference between drawing Lewis structures for ionic compounds and covalent molecules?

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Lewis Structures for Covalent Molecules

Molecule CO2 HCN CH2O

Draw the Lewis dot structure of the

individual atoms.

Find number of bonds each can

form (each unpaired electron

= 1 bond).

Determine the central atom (atom with the least # of valence electrons,

never H).

Draw the Lewis structure (place

electrons in bonds, then place remaining

electrons around their correct

atoms).

Shapes of Molecules (VSEPR)

Molecule CO2 HCN CH2O Central

atom (see part 2)

# Bonding atoms on

central atom

# Lone pairs on central atom

Shape

Build molecule and draw

picture

7. Compare the 3D model you built to the shapes shown on the VSEPR chart. Do they

match up?

8. What is missing from the models that you can see in the Lewis dot structure?

CW 10: Intermolecular and Intramolecular Forces

Types of Attractive Forces Notes

Type Intermolecular Intramolecular

Dispersion Forces Dipole-dipole Forces H-Bonding Covalent Bond Ionic Bond

Definition

Diagram

Other Info

Relative Strength

IMFs and State of Matter

The three states of matter are solids, liquids, and gases. The molecules of each state of matter attract each other, and the force of attraction increases as the distance between molecules decreases.

1. When liquid water evaporates into gaseous water, are any bonds between H atoms and O atoms within a molecule broken?

2. On average, are the intermolecular forces stronger in liquid water or gaseous water?

3. Consider the two molecules below to complete the table.

Molecule Picture Intramolecular Forces

(circle one) Intermolecular Forces (circle all that apply)

Water

H2O

Polar Covalent Nonpolar Covalent

Ionic

Ion-dipole H-bonding

Dipole-dipole Dispersion

Hydrogen sulfide

H2S

Polar Covalent Nonpolar Covalent

Ionic



4. Which of the substances above would have weak attractions and be more likely to be a

gas at room temperature? Explain in terms of the intramolecular and intermolecular forces present.

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5. Methane and methanol have nearly the same chemical formula. Identify the forces present in each molecule.

Molecule Picture Intramolecular Forces

(circle one)

Intermolecular Forces (circle all that

apply)

Methane

CH4

Polar covalent Nonpolar covalent

Ionic



Methanol

CH3OH


Ionic



6. In terms of intermolecular forces, why is methane a gas and methanol is a liquid?

IMFs and Surface Tension of Liquids

Surface tension is a measure of how strongly attracted molecules within a liquid are to each other. This means that the stronger the attraction between molecules, the higher the surface tension.

When examining a bead of water on a car, the attractions between the water molecules are much stronger than the attraction of the water molecules to the car surface (or to the air). Therefore, water forms “beads” rather than spreading out.

7. Place one drop of isopropyl alcohol and one drop of water onto the desk. Based on your

observations, which substance would have stronger intermolecular forces?

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IMFs and Phase Changes

Evaporation occurs when the attractive forces between molecules are weakened enough that molecules can escape the liquid phase and enter the gas phase.

8. Use your finger “push up” the two drops from Question 7 about two inches, so you have two “streaks”. Observe which substance evaporates more quickly.

9. In Question 7, which substance had the stronger intermolecular forces? Does this agree

with your observations about which evaporated more quickly?

10. Both water and isopropyl alcohol can form hydrogen bonds. Count the number of hydrogen bonds that each substance can form.

Water

Isopropyl alcohol

Number of hydrogen bonds: Number of hydrogen bonds:

11. Which molecule is “better” at forming hydrogen bonds?

12. How does this explain the difference in evaporation time between water and isopropyl

alcohol?

13. Revisit the comparison between methane and methanol from Question 5. One of the substances has a boiling point of 65°C, while the other has a boiling point of -161.5°C. Which one is which? Explain.

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IMFs and Odor

Highly volatile substances easily evaporate because they have weak attractive forces between them. These compounds will usually have a strong odor.

For example, vinegar contains acetic acid, a volatile compound. This means that the vapor

above the vinegar will contain many evaporated acetic acid molecules, and if you sniff this vapor, your nose will collect a lot of acetic acid molecules. There they will dissolve into your nasal fluids, which will become acidic, and you will experience the sour, pungent smell of vinegar.

14. Waft the following substances to determine if they are volatile or not. Then determine

which intermolecular forces are present in each molecule.

Substance Picture Volatile or Not?

Intramolecular Forces (circle one)

Intermolecular forces (circle one)

Sodium chloride

Polar covalent

Nonpolar covalent Ionic



Water

Polar covalent




Salt water +

Polar covalent




Ethyl acetate

Polar covalent




Ethyl acetate

and water +


Ionic



Journal Write 5

Explain your observations of the volatility of the substances in terms of the intermolecular forces present in each molecule.

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CW 11: Lewis Dot Structures (Ionic)

Formation of Cations

1. Complete the table below using a table of elements. All will form cations.

Element Property Before Making an Octet After Making an Octet

Na

Electron configuration

1s2 2s2 2p6 3s1 1s2 2s2 2p6

# Protons 11 11

# Electrons 11 10

Charge 0 +1

Lewis Dot Structure

Mg


# Protons

# Electrons

Charge

Lewis Dot Structure

Al


# Protons

# Electrons

Charge

Lewis Dot Structure

2. Explain why these atoms tend to form cations.

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Formation of Anions

3. Complete the table below using a table of elements. All will form anions.

Element Property Before Making an Octet After Making an Octet

N


1s2 2s2 2p3 1s2 2s2 2p6

# Protons 7 7

# Electrons 7 10

Charge 0 -3

Lewis Dot Structure

O


# Protons

# Electrons

Charge

Lewis Dot Structure

F


# Protons

# Electrons

Charge

Lewis Dot Structure

4. Explain why these atoms tend to form anions. 5. Would neon form an anion? Explain why or why not.

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Ionic Compounds

In ionic compounds, one atom (metal) transfers one or more valence electrons to another atom (nonmetal). A Lewis structure indicates the direction of this transfer as well as the charge of the resulting ions.

6. Draw Lewis dot structures for the following ionic compounds. a. NaBr

b. MgO

c. Na3N

d. AlF3

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CW 12: Build an Ionic Formula

Polyatomic Ions

Some ions consist of a group of covalently bonded atoms that tend to stay together as if they were a single ion. Such ions are called polyatomic ions. An example is the nitrate ion, NO3

-1. This polyatomic ion contains one nitrogen atom covalently bonded to three oxygen atoms, and has an overall charge of -1. Note that many of the names of polyatomic ions end with –ate or –ite. When you see these endings, it should signal to you to use this table.

1. Complete the table below by filling in your own polyatomic ions from the table.

Polyatomic Ion Name Polyatomic Ion Formula Number and Kind of Atoms Overall charge

Nitrate NO3-1 N: 1 O: 3 -1

Phosphate

Sulfate

Acetate

Carbonate

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2. Use the following blocks to represent your ions. Indicate the charge of each. Cations Formula and Charge Anions Formula and Charge

Light pink – calcium Black – sulfate

Lime – sodium Gray – chloride

Light Blue – copper Orange – acetate

Dark Blue – aluminum Yellow – oxide

Green – lead Red – nitride

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3. Build a Lego representation of each of the compounds below. Place cation blocks on top and anion blocks on the bottom. Each compound must have equal lengths of Lego blocks top and bottom.

Rules for Writing Ionic Formulas:

• Cation is first, anion second • The overall charge is always zero

• Subscripts indicate the number of atoms in the compound, and ones are never written.

• Parentheses are used around polyatomic ions to show that the subscript applies to the entire ion

Combining Ions

Cation Formula

Anion Formula

Number of

Cations

Number of

Anions

Total +

charge

Total –

charge Ionic Formula

Aluminum

chloride

Calcium

sulfate

Lead (IV)

oxide

Copper (II)

chloride

Aluminum

oxide

Copper (II)

nitride

Sodium

sulfate

Calcium

nitride

Journal Write 6

Examine the combining ions and the formulas you wrote in your data table. Develop a general procedure explaining how to write a formula for an ionic compound, given its charges. What is the relationship between the size of the Lego piece and the charge of the ion?


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CW 13: Naming Ionic Compounds

1. Explain why sodium obtains a stable electron configuration by losing its valence electron.

Na: 1s2 2s2 2p6 3s1

2. For each of the following, write if the atom will gain or lose electrons, the number of electrons it will gain or lose, and the resulting charge of the ion.

Atom Gain or Lose?

How many?

Charge? Atom Gain or Lose?

How many?

Charge?

Li N

Be O

B F

C Cl

3. The charges that an ion

forms are known as oxidation states. Write the oxidation state for groups 1, 2, 13, 15, 16, and 17 on the table below.

4. Why is it difficult to

assign an oxidation state to group 14?

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Name to Formula

• Write the symbol and charge for the cation. o Cations with one ionic charge: If an element belongs to group 1, 2, or 3, the charge

of the cation is known. Look up the symbol for the element and determine its charge from its position on the table of elements.

o Cations with more than one ionic charge: Some transition metals may have more than one possible charge. In this case, the charge of the cation is indicated using roman numerals.

o Polyatomic ions: check the polyatomic ion list for these cations.

• Write the symbol and charge for the anion. o Element: Change the “-ide” ending back into “-ine”. Look up the symbol for the

element and determine its charge from its position on the table of elements. o Polyatomic ions: check the polyatomic ion list for these anions.

• Place parenthesis around polyatomic ions.

• Ensure that the overall charge is neutral (zero) by adding subscripts.

5. Write formulas using the names below. a. Magnesium Acetate

b. Calcium Phosphate

c. Aluminum Sulfide

d. Iron (III) Oxide

e. Lead (II) Sulfate

f. Vanadium (V) Phosphate

g. Copper (II) Chloride

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Formula to Name

• Name the cation. o Cations with one ionic charge: If an element belongs to group 1, 2, or 3, it has only

one ionic charge. Write the name of the element exactly as it appears on the table of elements.

o Cations with more than one ionic charge: Place parenthesis around polyatomic ions (if present). Determine the charge of the cation from the anion. Indicate the charge in the name using roman numerals.

o Polyatomic ions: check the polyatomic ion list for these cations.

• Name the anion: o Element: Find the symbol of the element on the table of elements. Change the

ending to “- ide”. o Polyatomic ions: check the polyatomic ion list for these anions.

6. Write the names of the ionic compounds below.

a. NaCl

b. Ca(OH)2

c. MgBr2

d. FeCl2

e. FeCl3

f. Zn(OH)2

g. Co3N2

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CW 14: Naming Acids and Bases

An acid is a compound that contains one or more hydrogen atoms and produces hydrogen ions when dissolved in water. In general, an acid molecule consists of an anion combined with as many hydrogen ions as needed to make the molecule neutral.

Naming Common Acids

Anion Ending Example Stem Acid Name Example

-ide chloride, Cl-1 chlor hydro-(stem)-ic acid hydrochloric acid -ite sulfite, SO3

-1 sulf (stem)-ous acid sulfurous acid

-ate nitrate, NO3-1 nitr (stem)-ic acid nitric acid

When writing formulas for acids from a name, use the same guideline as above, then check charges to ensure that the molecules are neutral.

1. For each of the following, identify the anion, the anion ending and the stem.

Acid Anion Anion Name (Highlight Ending) Stem

HCl ____ ______

HClO

HClO2

HClO3

HClO4

2. Name each of the following acids.

a. H2S

b. HCl

c. H2SO4

d. HClO4

e. H2SO3

f. HClO2

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3. Complete the table below by filling in either the name or the formula as needed. Name Formula

HNO3

HNO2

HBr

Hydrofluoric acid

Carbonic Acid

Nitrous Acid

How do you determine the name and formula for a base?

A base is generally an ionic compound that produces hydroxide when added to water. They are named in the same manner as other ionic compounds, but the anion is always hydroxide (OH-1).

4. Name or write the formula for the following bases. a. NaOH

b. Potassium hydroxide

c. Fe(OH)2

d. Iron (III) hydroxide

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CW 15: Metallic Bonding

Go to LEFFELlabs/ unit 2 to complete the PowerPoint while answering the questions below. 1. How do metals form cations? Explain using an electron configuration.

2. Where do the electrons in the electron sea come from?

3. Draw a picture that shows the electron sea model. Label: metal cation, delocalized electron, electron sea.

4. Write a caption for your picture. Words to use: repulsion, delocalized electron, cation.

5. Explain each of the following in terms of the electron sea model. a. Metals are shiny.

b. Metals are good conductors.

c. Metals are malleable and ductile.

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Metallic Bonding Model

6. What do the large pearl beads represent? What do the same silver beads represent?

7. Describe the motion of the electrons relative to the cations.

8. Use your model to define alloy. What did the replacement (wooden) beads represent?

9. How might making your metal into an alloy affect the properties of the metal?

Journal Write 7

Evaluate your model. What does it illustrate well? What changes could you make to create a more accurate model?


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Reference Materials

Polyatomic Ions

H3O+ hydronium CrO42– chromate

Hg22+ dimercury (I) Cr2O7

2– dichromate

NH4+ ammonium MnO4

– permanganate

C2H3O2–

CH3COO– acetate

NO2– nitrite

NO3– nitrate

C2O42– oxalate O2

2– peroxide CO3

2– carbonate OH– hydroxide

HCO3– hydrogen (bi)carbonate CN– cyanide

PO43– Phosphate SCN– thiocyanate

ClO– hypochlorite SO32– sulfite

ClO2– chlorite SO4

2– sulfate

ClO3– chlorate HSO4

– hydrogen sulfate

ClO4– perchlorate S2O3

2– thiosulfate

Electronegativity Values

H 2.1

Li 1.0

Be 1.5

B 2.0

C 2.5

N 3.0

O 3.5

F 4.0

Na 0.9

Mg 1.2

Al 1.5

Si 1.8

P 2.1

S 2.5

Cl 3.0

K 0.8

Ca 1.0

Ga 1.6

Ge 1.8

As 2.0

Se 2.4

Br 2.8

Rb 0.8

Sr 1.0

In 1.7

Sn 1.8

Sb 1.9

Te 2.1

I 2.5

Cs 0.7

Ba 0.9

Ti 1.8

Pb 1.9

Bi 1.9

Electronegativity and Bond Types Difference in electronegativity Most probable bond type Example

0.0 – 0.4 Covalent – Nonpolar H – H (0.0)

0.4 – 1.0 Covalent – Moderately Polar H – Cl (0.9) 1.0 – 2.0 Covalent – Very Polar H – F (1.9)

≥ 2.0 Ionic Na+Cl– (2.1)

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VSEPR Chart

Considering the central atom…

# of Bonded Atoms # of Lone Electron Pairs Molecular Shape 2 0 Linear

3 0 Trigonal planar 2 1 Bent

4 0 Tetrahedral

3 1 Trigonal pyramidal 2 2 Bent

5 0 Trigonal bipyramidal 4 1 Seesaw

3 2 T-shape

2 3 Linear 6 0 Octahedral

5 1 Square pyramidal 4 2 Square planar

VSEPR Images

Linear

Trigonal planar

Tetrahedral

Trigonal

pyramidal Trigonal bipyramidal

Bent

Seesaw

T-shape

Octahedral

Square

pyramidal

Square planar

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Table of Elements

name: honors unit 2 class: bonds, james bonds · 10/2/2019 · 1. what is the electron...

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